:-NRLF E77 fl?D INTERNATIONAL ATOMIC WEIGHTS, 1915 Symbol. Atomic weight. Symbol. Atomic weight. Aluminum Antimony Al 27.1 120.2 39.88 74.96 137.37 9,1 208.0 11.0 79.92 112.40 132.81 40.07 12.00 140.25 35.46 52.0 58.97 93.5 63.57 162.5 167.7 152.0 19.0 157.3 69.9 72.5 197.2 3.99 163.5 1.008 114.8 126.92 193.1 55.84 82.92 139.0 207.10 6.94 174.0 24.32 54.93 200.6 Molybdenum ..... Mo Neodymium ... Nd Neon Me Nickel Ni Nitoa (radium emanation) Nt Nitrogen N 96.0 144.3 20.2 58.68 222.4 14.01 190.9 16.00 106.7 31.04 195.2 39.10 140.6 226.4 102.9 85.45 101.7 150.4 44.1 79.2 28.3 107.88 23.00 87.63 32.07 181.5 127.5 159.2 204.0 232.4 168.5 119.0 48.1 184.0 238.5 51.0 130.2 172.0 89.0 65.37 90.6 Sb A Arsenic . . As Ba Be Bi . . Os B O Br Palladium . . . Pd . . Cd Phosphorus . . . . . . P . . Pt Caesium Cs Calcium Ca . . . K Carbon c Praseodymium . . Radium Rhodium .... . . . Pr . . Ra . . . Rh . . Rb Ce Chlorine Cl Cr Cobalt Co . . . Ru Columbium Cb . . Sa Cu Scandium .... . . . Sc Dysprosium Dv . . . Se . . . Si Er Eu Silver . . Ac F . . . Na Gadolinium Gd . . . Sr Ga Sulphur . . . S Germanium Gold Ge Tantalum .... Ta Au Tellurium .... . . . Te He . . . Tb Holmium . Ho Thallium . ... . . . Tl . . Th . . Tm H Indium ... In I Tin Sn Ir Ti Fe W Krypton . . Lanthanum Lead . . . Lithium . . Kr La Uranium . . . U . . . V Pb Li . . Xe Ytterbium (Neoytterbium) Yttrium . ... Zinc . . . Yb . . . Yt . . Zn Lutecium . Lu Magnesium Manganese Mg Mn He Zirconium .... . . . Zr From the collection of the n PreTinger v iJibrary San Francisco, California 2006 SYNTHETIC INORGANIC CHEMISTRY A LABORATORY COURSE FOR FIRST YEAR COLLEGE STUDENTS BY ARTHUR A. BLANCHARD, PH.D. A ssociate Professor of Inorganic Chemistry at the Massachusetts Institute of Technology SECOND EDITION, WITH SUPPLEMENT FOURTH THOUSAND NEW YORK JOHN WILEY & SONS, INC. LONDON : CHAPMAN & HALL, LIMITED 1916 COPYRIGHT, 1908, 1910, 1916 BY ARTHUR A. BLANCHARD Stanbope ]press F. H. GILSON COMPANY BOSTON. U.S.A. PREFACE THIS series of notes was designed to serve as a guide for laboratory work and study in Inorganic Chemistry during the second term of the first year at the Massachusetts Insti- tute of Technology. It had been felt for some time that Qualitative Analysis, which was previously made the basis for laboratory practice during that period, did not fully meet the requirements and that a course based upon the actual preparation of typical chemical substances might prove more satisfactory. In consequence, notes in essen- tially the form now published were prepared during the year 190607, they being the direct outcome of several years' pre- vious trial of a limited amount of preparation work with the classes. The present book is a thorough revision of those notes in the light of experience in their actual application. During the first term's study of chemistry there can be little doubt that a course of simple experiments, such as has long been in use, in the methods of formation and in the study of the properties of the non-metallic elements oxy- gen, hydrogen, the halogens, sulphur, nitrogen, and carbon and their compounds, is the most satisfactory. But when it comes to the study of the metallic elements, three options as to laboratory work present themselves : First, a continua- tion of experiments similar in nature to those of the first term ; second, Qualitative Analysis ; third, Preparation Work. The disadvantages of the first plan are that the experiments are so quickly performed and so alike in character that they fail to arouse much enthusiasm in the student or to leave very vivid impressions on his mind. .Qualitative analysis is in many ways a most excellent basis for teaching the chemistry of the metallic elements ; but its chief disadvan- tages are : First, that it is one-sided, it dealing as it does iii 304847 IV PREFACE almost exclusively with the chemistry of solutions and the formation of highly insoluble bodies ; second, that it requires the sequence followed in the lectures to be that of the quali- tative procedure instead of a more natural one based on the periodic classification ; and third, that it is well-nigh impos- sible to keep from the student's mind the false idea that the end and aim of qualitative analysis is principally " to get the unknowns correct." Some of the advantages which seem to be possessed by a course of preparation work such as outlined in the following pages are: 1. The sequence of the exercises may follow that of the lectures. 2. Very varied types of chemical change are illustrated, both those in the furnace and those in solution. In solu- tion advantage is taken not only of high degrees of insol- ubility, but also of differences in solubility among the more soluble bodies as well as of differences in the effect of temperature on solubility. 3. The danger of the work becoming a mechanical following of directions is reduced by the introduction of study questions and experiments with each exercise. 4. In its effect in awakening the student's interest this line of work has proved particularly successful, the making of preparations is, in fact, in its very nature one of the most fascinating forms of chemical work. Since each preparation requires a good deal of time and thought, and the product when obtained is something definite and tangible, the knowl- edge thus gradually absorbed is more definite and less easily forgotten than when the laboratory work consists of a large number of test tube reactions. After the completion of such a course as this, if the student commences analytical work with some conception of the sources and methods of obtaining the substances PREFACE V which he is to use as reagents, etc., there can be no doubt that the latter work will then have a much deeper meaning. The plan kept in mind in preparing this course is, briefly, as follows : The greater part of the preparations selected are of industrial importance, and for the starting point of each either natural products or crude manufactured materials are used so far as is possible. The course does not aim to be an exhaustive one in chemical preparations, but a limited number of exercises are selected to illustrate the most impor- tant types of compounds of the common elements and the most important methods. Two Or three times as many exercises are furnished as any one student will be able to complete in the time usually allotted ; thus different students may be assigned different preparations. The notes for each exercise are divided into three parts : I. A discussion of the object of the exercise, with an out- line of the principle of the method and the reasons for the steps involved. II. Working directions which, if carefully followed, should result in obtaining a satisfactory product. It is believed far better to make the directions very explicit, for the reason that the inexperienced student may easily become discouraged by failures due to difficulties which he is unable to foresee. Difficulties enough are sure to arise to develop originality and resourcefulness. III. Questions for study which involve additional laboratory experiments, the consulting of text-books, and original reasoning. At the end of each group of exercises is furnished a set of general study questions, and this arrangement of the exercises in groups is such as to bring out the relationships shown in the periodic classification of the elements. In the discussions and questions given with the various exercises it is assumed that the student has an elementary knowledge of the electrolytic dissociation theory and of the principle of mass action. In the opinion of the author Vl PREFACE a great opportunity is lost for bringing out relationships among chemical phenomena if these principles are not taught during the first term's study of college chemistry and their applications pointed out in connection with later work in inorganic preparations and in analytical chemistry. The effort has been to make the questions such as cannot be answered mechanically. Some of the questions may, in consequence, seem rather difficult and incapable of direct answers ; the object of the questions is, however, not solely to bring forth correct statements of facts and theories, but is also to teach the student to use his head in seeking for the significance of facts and in reasoning from one fact to another. Acknowledgment is due to many sources for the outline of the greater part of the methods given. The details of all of them have, however, been very carefully worked over and adapted for the purpose in view. In conclusion the author wishes to express his sense of obligation to Professor Henry P. Talbot, head of the Depart- ment of Chemistry, at whose request the preparation of these notes was undertaken ; also to other members of the instructing staff at the Massachusetts Institute of Tech- nology for helpful criticism and suggestions, and particularly to Mr. J. W. Phelan, to whose efficient management of the laboratory instruction is due any success with which this course has met at this Institute. This little book is presented for publication with the desire to offer for the consideration of those in charge of the instruction at other institutions a plan of first-year work which has quite new and perhaps advantageous features. It is hoped that it may do its part in securing recognition of the importance of synthetic or preparative work in a well-balanced course of chemical training. ARTHUR A. BLANCHARD. March, 1908. PREFACE TO SECOND EDITION THE plan of work embodied in the first edition has remained unchanged in the second edition, but many minor improvements have been made and considerable new mate- rial has been added. Several of the procedures have been altered, so that good results may be more confidently expected from inex- perienced students. A few new preparations have been added to the list in order to include important types of processes which were not well represented in the first edi- tion. The general questions following each group of exer- cises have been entirely rewritten, and now present a more consistent plan to bring out the main characteristics of the various groups and the relations among the groups. The number of these questions has been considerably reduced, and it is now felt even more strongly than before that all of the general questions should be mastered by every student who takes the course. For the convenience of teachers as well as of enthusiastic students, a number of additional gen- eral questions have been placed in the appendix. An intro- ductory section has been included, explaining principles and details of laboratory manipulation. A new chapter (Chap- ter VII) has been added, which embraces the non-metallic elements, and is intended to recall and to broaden the knowledge which the student was supposed to possess be- fore commencing the study of the metallic elements. Some useful tables have been added in the appendix, and a few cuts have been inserted in the text. June, 1910. CONTENTS INTRODUCTORY TO THE STUDENT i NOTES ON LABORATORY MANIPULATION 5 1. Precipitation-; Crystallization 5 2. Pouring 6 3. Transferring Precipitates or Crystals 6 4. Filtering; Collecting Precipitates 7 (a) A Coarse-Grained Crystal Meal 7 (b) Filtering with Suction ; Witt Filter 7 Suction 8 (c) Filtering without Suction 9 (d) Filtering Corrosive Liquids 10 (e) Cloudy Filtrates u (/) To Keep Liquids Hot during Filtration 1 1 (g) Cloth Filters u 5. Washing Precipitates 12 (a) Washing on the Filter , 12 (b) Washing by Decantation 13 6. Evaporation 14 7. Dissolving Solid Substances 15 8. Crystallization 15 9. Drying 18 10. Pulverizing 19 11. Dry Reactions; Furnaces 20 CHAPTER I. ALKALI AND ALKALINE EARTH METALS ... 23 1. Potassium Nitrate from Sodium Nitrate and Potassium Chloride 25 2. Caustic Potash from Wood Ashes 28 3. Sodium Carbonate by the Ammonia Process 31 4. Chemically Pure Sodium Chloride 34 5. Ammonium Bromide 37 6. Strontium Hydroxide from Strontium Sulphate .... 39 7. Strontium Chloride from Strontium Sulphate 42 8. Barium Oxide and Barium Hydroxide from Barium Car- bonate 45 GENERAL QUESTIONS I ................. 47 X CONTENTS CHAPTER II. ELEMENTS OF THE THIRD GROUP OF THE PERIODIC SYSTEM 51 9. Boric Acid 53 10. Alum from Cryolite . . 54 11. Aluminum Sulphide 59 GENERAL QUESTIONS II 61 CHAPTER III. HEAVY METALS OF THE FIRST Two GROUPS OF THE PERIODIC SYSTEM 63 12. Crystallized Copper Sulphate from Copper Turnings . . 65 13. Cuprous Chloride 66 14. Ammonium and Copper Sulphate 69 15. Ammonio-Copper Sulphate 71 16. Zinc Oxide 74 17. Mercurous Nitrate . . . . 76 18. Mercuric Nitrate 77 19. Mercuric Sulphocyanate 78 GENERAL QUESTIONS III 80 CHAPTER IV. ELEMENTS OF THE FOURTH GROUP OF THE PERIODIC SYSTEM 83 20. Stannous Chloride 85 21. Stannic Sulphide (Mosaic Gold) 88 22. Stannic Chloride (Anhydrous) 90 23. Lead Nitrate 92 24. Lead Dioxide 93 25. Red Lead 95 GENERAL QUESTIONS IV 96 CHAPTER V. ELEMENTS OF THE FIFTH GROUP OF THE PERIODIC SYSTEM 99 26. Ortho-Phosphoric Acid 101 27. Arsenic Acid 103 28. Antimony Trichloride from Stibnite 106 29. Sodium Sulphantimonate 108 30. Antimony Pentasulphide no 31. Metallic Antimony in 32. Basic Bismuth Nitrate (Bismuth Subnitrate) 112 GENERAL QUESTIONS V 114 CONTENTS XI CHAPTER VI. HEAVY METALS OF THE SIXTH, SEVENTH, AND EIGHTH GROUPS OF THE PERIODIC SYSTEM ... 117 33. Potassium Bichromate from Chromite 119 34. Potassium Chromate from Potassium Bichromate ... 122 35. Chromic Anhydride 122 36. Chromic Alum 124 37. Chromium Metal by the Goldschmidt Process 127 38. Manganese Chloride from Waste Manganese Liquors . 128 39. Potassium Permanganate from Manganese Dioxide . . 131 40. Manganese Metal by the Goldschmidt Process .... 134 41. Ferrous Ammonium Sulphate 135 42. Ferric Ammonium Alum 136 GENERAL QUESTIONS VI 138 CHAPTER VII. NON-METALLIC ELEMENTS OF THE SIXTH AND SEVENTH GROUPS OF THE PERIODIC SYSTEM, 143 43. Potassium Iodide 145 44. Hydriodic Acid 146 45. Potassium Chlorate 149 46. Potassium Bromate 153 47. Potassium lodate 155 48. lodic Acid; Iodine Pentoxide 157 49. Potassium Perchlorate 159 50. Sodium Thiosulphate 160 GENERAL QUESTIONS VII 163 APPENDIX Additional General Questions 165 Atomic Weights 176 Periodic Arrangement of the Elements 177 Table of Solubilities 178 SYNTHETIC INORGANIC CHEMISTRY Introductory to the Student THE following exercises are designed to illustrate the principles and methods involved in the preparation of a number of the most important chemicals. Where possible the method employed resembles that actually used on an industrial scale ; where this is, however, impossible on the limited scale of the laboratory, mention is made of the fact, with reasons therefor. On account of the limitations connected with work on a laboratory scale, it is of course impossible to get as high percentage yields as could be obtained on a large commercial scale. The amounts ob- tained of each preparation are to be weighed and recorded, but the chief stress is to be laid upon the excellence of the product rather than upon its quantity. A larger number of preparations is given than it is expected that the student can accomplish in fifteen weeks with but four hours per week in the laboratory. Each student, therefore, will be assigned certain of the exercises which he is expected to thoroughly master, and which he is expected to perform entirely independently. But almost equal in importance is it for him to know the work which the students about him in the laboratory are performing. To this end it is important that the directions, and more especially the first sections which discuss the principles involved, be studied for each exercise, and then that the work of neighboring students actually at work upon the preparations be observed and discussed with them in I 2 ' - t\NTHfeTie. INORGANIC CHEMISTRY the odd moments which will invariably occur when waiting for evaporations or filtrations to take place. Directions for laboratory work: The notes for each preparation are divided into three parts: I. Discussion of the general principles involved. II. Directions for actual manipulation. III. Study questions. Part I is to be read and understood before commencing work in the laboratory. Part II, being the working directions, is to be kept av hand while carrying out the manipulations. These direc- tions do not need to be recorded in the laboratory note< book ; but it is essential, nevertheless, to keep a laboratory notebook in which to enter all important observations and data ; such as, for example, appearance of solutions (color, turbidity) ; appearance of precipitates or crystals (color, size of grains, crystalline form) ; results of all weighings or measurements ; number of recrystallizations ; results of tests for purity of materials, etc. Part III constitutes directions for study based upon the particular preparation. This will involve : (i) labora- tory experiments and direct entries in the laboratory note- book ; (2) consultation of reference books, of which all that are necessary will be found upon the shelf in the laboratory ; (3) original reasoning. The answers to the questions should be written in the laboratory notebook following the entries for the exercise, and this book should be submitted at the same time as the preparation for the approval of an instructor. Besides the specific study questions for each prepara- tion there are, accompanying each group of exercises, general questions relating to the whole group; and these are to be worked out by every student. The answers to these questions are to be written on a certain prescribed INTRODUCTORY 3 kind of paper and handed in at the office, neatly bound, within the times which will be posted. In preparation work it is frequently necessary to wait for considerable periods of time for evaporations, crystalliza- tions, etc., to take place. This time may be utilized for work upon the study questions and experiments, but even then it is advisable to have usually more than a single preparation under way. Thus no time need be wasted by the energetic student who plans his work we 1 !. NOTES ON LABORATORY MANIPULATION THESE notes are intended simply to help the student in foreseeing and in overcoming some of the difficulties that arise in experimental work. They by no means make it unnecessary for him to exercise ingenuity and originality in planning and carrying out the details of laboratory work. At the outset these notes should be read through carefully ; then when in the later work references to specific notes are made their general bearing will be better appreciated. i. PRECIPITATION; CRYSTALLIZATION In the majority of chemical processes which are carried out in the wet way, separations are accomplished by taking advantage of differences in solubility. In case a certain product is extremely insoluble and is formed almost instan- taneously when solutions containing the requisite compo- nents are mixed, the process is called precipitation and the insoluble substance is called the precipitate. If the product to be formed is less insoluble, so that it separates more slowly, or only after evaporating away a part of the solvent, the process is called crystallization. In some cases the precipitate, or the crystals, constitute the desired product ; in other cases a product which it is necessary to remove from the solution before the desired product can be obtained pure. In either case it is neces- sary to make as complete a separation as possible of the solid from the liquid. This involves the manipulations described under Notes 2, 3, and 4. 5 O NOTES ON LABORATORY MANIPULATION 2. POURING In pouring a liquid from a vessel, either into a filter or into another vessel, care must be taken not to slop the liquid nor to allow it to run down the outside of the vessel poured from. To this end touch a stirring rod to the lip of the FIG. I FIG. 2 dish or beaker (Fig. i) and allow the liquid to run down the rod. 3. TRANSFERRING PRECIPITATES OR CRYSTALS If large crystals have separated from a liquid they may be picked out, or the liquid may be poured off. If a precipitate or a crystalline meal has formed it must be drained in a filter funnel. First pour off the liquid (see Note 2) through the filter if necessary, so as to save any FILTERING 7 floating particles of the solid then pour the main part of the damp solid into the filter. A considerable part of the solid will adhere to the dish ; most of this may be scraped out by means of a spatula, but the last of it is most easily rinsed into the filter. For rinsing, a jet of water from the wash bottle (Fig. 2) may be used if the solid is very insoluble. If the solid is soluble in water, some of the saturated solution may be poured back into the dish from out of the filter bottle, and by means of this the last of the solid may be removed to the filter. 4. FILTERING; COLLECTING PRECIPITATES (a) A coarse-grained crystal meal can best be collected in a filter funnel in which a perforated porcelain plate is placed, and the mother liquor clinging to the crystals can best be removed with the aid of suction (see next para- graph). (fi) Pilfering with Suc- tion ; Witt Filter. With a fine-grained crystal meal, or a precipitate which is not of such a slimy character as to clog the pores of the filter paper, a Witt filter is most advantageously used. For most general use a 5-inch filter funnel should be fitted tightly by means of a rub- ber stopper into the neck of a 500 cc. filter bottle (Fig. 3). Place a i^-inch FIG. 3 perforated filter plate in the 8 NOTES ON LABORATORY MANIPULATION funnel and on this a disk of filter paper cut so that its edges will turn up about 3 mm. on the side of the funnel all the way around. Hold the disk of dry paper in the right position, wet it with a jet from the wash bottle, draw it firmly down against the filter plate by applying the suction, and press the edges firmly against the side of the funnel, so that no free channel shall remain. In pouring the liquid, direct it with a stirring rod (Note 2) onto the middle of the filter ; do not allow it to run down the side of the funnel, as this might turn up the edge of the paper and allow some of the precipitate to pass through. After bringing all of the solid upon the filter it may be freed from a large part of the adhering liquid by means of the suction, and it may then be purified by washing with a suitable liquid (see Note 5). The Witt filter is very generally useful for the purpose of separating a solid product from a liquid. In cases that the liquid runs slowly, the rate of filtration can be increased by using a larger filter plate or still better a Biichner fun- nel and thereby increasing the filtering area. The student should, however, avoid using the suction indiscriminately, for in many cases it is, as explained in paragraph c, a positive disadvantage. Suction. The most convenient source of suction is the Richards water pump, which can be attached directly to the water tap. If the water is supplied at a pressure of somewhat over one atmosphere (34 feet of water), a vacuum of very nearly an atmosphere can be obtained. If the pres- sure is insufficient, an equally good vacuum can be obtained by means of the suction of the escaping water. To this end the escape pipe must be prolonged by a tube sufficiently constricted to prevent the sections of the descending water column from breaking and thus allowing air to enter from the bottom. FILTERING 9 To keep the suction pump working continuously, how- ever, is extravagant of water as well as being a nuisance in the laboratory on account of the unnecessary noise. Consequently this rule is made and must be observed: The suction pump must never be kept in operation more than two minutes at one time. If suction must be applied for more than that length of time, the vacuum which is produced inside of the two minutes may be maintained in the suction bottle by closing the latter air tight. For this purpose the bottle is to be fitted as follows : Connect a short piece of rubber tube with the side arm of the filter bottle. Provide this tube with a screw cock and connect its further end with a short piece of glass de- livery tube tapered a little at each end and rounded in the flame. (See Fig. 3 on page 7.) The short glass tube can be attached to and removed from the pump at will, and a vacuum once produced in the bottle can be preserved by closing the screw cock. Thus, for example, if all the joints of the bottle are tight, a slimy precipitate may be left filtering under suction over night or even longer. (V) Filtering without Suction. A slimy or gelatinous precipitate can be collected much better without suction. Suction drags the solid matter so completely into the pores of the filter that in most cases the liquid nearly ceases to run. A filter funnel and filter should be chosen large enough to hold the entire precipitate. The filter paper should be folded twice and then opened out in the form of a cone and fitted into the funnel (Instructions). The corners of the filter should be cut off round, and the upper edge of the filter should come about one-half inch below the rim of the funnel. It is usually best to fit the paper carefully into the funnel, to wet it and press it up IO NOTES ON LABORATORY MANIPULATION against the glass all around, so that there will be no air channels. In the case of slow-running liquids, if a large filter is used, it may be filled at intervals and left to take care of itself the rest of the time while other work is being done. In case a considerable weight of liquid is to come on the point of the filter, this may be reenforced by means of a piece of linen cloth, which should be placed under the middle of the filter paper before it is folded, and should then be folded in with it so as to strengthen the point. After the precipitate is collected in the filter and drained, it should if necessary be washed (see Note 5 on page 12). Both filtration and washing take place much more rapidly if the liquid is hot. Time can also usually be saved if the precipitate is allowed to settle as completely as possible before commencing to filter. The clear liquid can then be decanted off, or if necessary poured rapidly through the filter before the latter becomes clogged with the main part of the precipitate. (d) filtering Corrosive Liquids. Solutions of very strong oxidizing agents, concentrated solutions of the strong acids and bases, and concentrated solutions of a few salts of the heavy metals notably zinc chloride and stannous chloride attack filter paper strongly. Ordinary paper is thus entirely unserviceable for filtration, but a felt made of asbestos fibers is in many cases very useful. Shredded asbestos, which has been purified by boiling with hydro- chloric acid and subsequent washing, is suspended in water ; the suspension is poured onto a perforated plate placed in a filter funnel ; and suction is applied whereby the water is removed and the fibers are drawn together to form a compact felt over the filter plate. Enough asbestos should be used to make a felt i to 3 mm. thick, and care must be taken to see that it is of uniform thickness and that no free FILTERING I I channels are left through which solid matter may be drawn. Before it is ready for use a considerable amount of water should be drawn through the filter, and the loose fibers should be rinsed out of the filter bottle. Before pouring the liquid onto the filter the suction should be started gently, and the liquid should be directed by means of a stirring rod (Note 2) onto the middle of the filter. If these precautions are not observed the felt may become turned up in places, so that the precipitate will pass through. (e) Cloudy Filtrates. When a filtrate at first comes through cloudy, it is usually sufficient to pour the first por- tion through the filter a second time. The pores of the filter soon become partially closed with the precipitate, so that then even the finest particles are retained. With some very fine-grained precipitates, repeatedly pour- ing the filtrate through the same filter will finally give a clear filtrate. Special kinds of filter paper are made to retain very fine precipitates, but they allow the liquid to pass much more slowly than ordinary filters, and their use is by no means essential in any of the following preparations. (f) To Keep Liquids Hot during Filtration. When liquids must be kept hot during a slow filtration, as, for example, when cooling would cause a separation of crystals that would clog the filter, it sometimes becomes necessary to surround the funnel with a jacket which is heated with steam or boiling water. In the following preparations the use of such a device will not be necessary, although there are several instances where it is necessary to work quickly to avoid clogging the filter. () Cloth Filters. In preparations made on a small scale, paper filters placed in ordinary filter funnels are inva- riably used if the liquid is not too corrosive. On a larger scale or in commercial practice, cloth is much used for fil- 12 NOTES ON LABORATORY MANIPULATION ters, and it can be made in the shape of bags or it can be stretched over wooden frames. The cloth or other filtering medium (asbestos, paper pulp, sand, etc.) has to be chosen in each case with reference to the nature of the precipitate and the corrosiveness of the liquid. Many of the preparations in this book, if carried out on a larger scale than given in the directions, would require the use of such cloth filters. It is often advantageous to tack one piece of cloth permanently across a wooden sup- port and on top of this to lay a second cloth. The precipi- tate can then be easily removed together with the unfastened cloth. For devices for rapid filtration and filtration in general on a large scale, a work on Industrial Chemistry should be consulted. 5. WASHING PRECIPITATES (a) Washing on the Filter. To remove completely the impurities contained in the mother liquor clinging to pre- cipitates or crystals, the solid is washed. Pure water is used for washing, provided the solid is not too soluble or is not decomposed (hydrolyzed) by it. Special directions will be given when it is necessary to use other than pure water. First, the solid should be allowed to drain as completely as" possible, then the wash liquid should be applied, prefer- ably from the jet of a wash bottle, so as to wet the whole mass and to rinse down the sides of the filter. If suction is used, suck the solid as dry as possible, then stop the suction while applying the washing liquid ; after the solid is thoroughly wet, suck out the liquid and repeat the washing. A little thought will make it clear that the washing is much more effective if the liquid is removed as completely WASHING PRECIPITATES as possible each time before applying fresh wash liquid, and that a number of washings with a small amount of liquid each time is more effective than fewer washings with much greater quantities of wash liquid. It is, of course, evident that with each washing the liquid should penetrate to all parts of the solid material. (b) Washing by Decantation. In case a precipitate is very insoluble it can be most thoroughly and quickly washed by decantation. This consists in allowing it to settle in a deep vessel and then in pouring (decanting) or siphon- ing off the clear liquid. Following this the precipitate is stirred up with fresh water and allowed to settle, and the liquid is again decanted off. By a sufficient number of repetitions of this process, the precipitate may be washed entirely free from any soluble impurity, after which it may be thrown on a filter, drained, and then dried. Most precipitates, even after they have settled as com- pletely as possible in the liquid from which they were thrown down, are very bulky, and their apparent volume is very large as compared with the ac- tual volume of the solid matter itself. For example, a precipitate of basic zinc carbonate (No. 16, page 74), after it has settled as completely as possible in a deep jar (Fig. 4), may still occupy a volume of 400 cc. When this bulky precipitate is dried, however, it shriv- els up into a few small lumps whose total volume is not more than 4 or 5 cc. If a precipitate, which is at first uniformly suspended in a liquid, is allowed to settle in a tall jar until it occupies but J of the original volume FIG. 4 14 NOTES ON LABORATORY MANIPULATION of the mixture (Fig. 4), any soluble substances will still remain uniformly distributed throughout the whole volume. If now the upper f, consisting of the clear solution, is drawn away, it follows that practically |- of the solution, containing of the soluble impurities, remains with the pre- cipitate. By stirring up the solid again with pure water, the soluble impurities become uniformly distributed through the larger volume, and on letting the precipitate settle and drawing off f of the liquid, as before, there will remain with the wet precipitate only X i = ^5 of the original soluble matter. After the third decantation the remaining suspen- sion will contain ^ X sV lir ^ tne original impurities, and so on. 6. EVAPORATION When it is necessary to remove a part of the solvent from a solution, as when a dissolved substance is to be crystallized out, the solution is evaporated. In some cases, where the dissolved substance is volatile or is decomposed by heat, the evaporation must take place at room temper- ature, but ordinarily in the following preparations the liquid may be boiled. The boiling down of a solution should always be carried out in a porcelain dish of such a size that at the outset it is well filled with the liquid. (Never evaporate in a beaker.) * The flame should be applied directly under the middle of the dish where the liquid is deepest ; the part of the dish against which the flame plays directly should be protected with wire gauze. Under no circumstances should the flame be allowed to play up over the sides of the dish : first, because, by heating where the dish is part cooled by liquid and part uncooled, there is great danger of breaking; second, because by heating the sides the film of liquid which creeps up is evaporated and the solid deposited becomes baked hard and in some CRYSTALLIZATION 1 5 cases is decomposed. To prevent the formation of a solid crust around the edges, which even at best will take place to some extent, the dish should occasionally be tilted back and forth a little, so that the crust may be dissolved, or loosened, and washed back into the middle of the dish. While evaporating a liquid over a flame it should be carefully watched, for if it should be forgotten and evapo- rate to dryness the dish would probably break and the product be spoiled. If a precipitate or crystals separate from the liquid and collect in a layer at the bottom, the dish will probably break, because where the solid prevents a free circulation of the liquid the dish becomes superheated, and then when in any one place the liquid does penetrate, the sudden cooling causes the porcelain to crack. Usually when a solid begins to separate from a boiling liquid the evaporation should- be stopped and the liquid left to crystal- lize. After that the mother liquor may be evaporated further in a smaller dish. 7. DISSOLVING SOLID SUBSTANCES The process of dissolving solid substances is hastened, first by powdering the substance as finely as possible, and second by raising the temperature. The solid and solvent should be heated together in a porcelain dish (not in a beaker), and care should be taken to keep the mixture well stirred, for if the solid should settle in a layer on the bot- tom, that part of the dish would become superheated and would be apt to break (see last paragraph in Note 6). 8. CRYSTALLIZATION (a) A great number of pure substances are capable of assuming the crystalline condition when in the solid form. 1 6 NOTES ON LABORATORY MANIPULATION Crystals are bounded by plane surfaces, which make definite and characteristic angles with each other and with the so- called axes of the crystals. The external form of a crystal reflects in some manner the shape or structure of the individual molecules of the substance ; for the crystal must be regarded as being built up by the deposition of layer on layer of molecules, all of which are placed in the same definite spatial relation to the neighboring molecules. When a substance takes on the solid form very rapidly (as when melted glass or wax cools) its molecules do not have an opportunity to arrange themselves in a regular order, and consequently the solid body is amorphous. The axes of the individual molecules point in every direction without regularity, and consequently the solid body possesses no crystalline axes or planes. It is evident from the above that the essential condition favoring the formation of perfect crystals is that the solid shall be built up very slowly. This is the only general rule which can be given in regard to the formation of perfect crystals. The excellence of a chemical preparation is in many cases judged largely from its appearance. The more uni- form and perfect the crystals, the better appearance the preparation presents. In the following preparations sometimes a pure melted substance is allowed to crystallize by simply cooling; in such a case the cooling should take place slowly. More often crystals are formed by the separation of a dissolved substance from a saturated solution. Perfect crystals can best be obtained in this case by keeping the solution at a constant temperature and allowing it to evaporate very slowly. This is easily accomplished in industrial works where large vats of solution can be kept at a uniform tern- CRYSTALLIZATION I/ perature with steam coils and allowed to evaporate day and night. On the laboratory scale it is almost impossible, first on account of variations in temperature, and next on account of dust which must fall into an uncovered dish. The majority of substances are more soluble at higher temperatures than at lower. If a solution just saturated at a high temperature is allowed to cool very slowly, it is possible for the solid to separate so slowly as to build up perfect crystals. This is an expedient that can be adopted to advantage in several of the following preparations. In many cases, however, when a saturated solution cools it becomes supersaturated, sometimes to a high degree. Then when crystallization is once induced it occurs with such rapidity that a mass of minute crystals, instead of a few large, perfect ones, is produced. To avoid this supersatu- ration a few seed crystals (i. e., very small crystals of the kind desired) may be placed in the solution. These form nuclei on which large crystals can be built up, and when they are present it is impossible for the solution to remain supersaturated. In carrying out the following preparations the principles just stated should be kept carefully in mind ; but in many instances specific suggestions will be given as to the easi- est method for obtaining good crystals of any particular substance. Large crystals, it is true, present a pleasing appearance, but oftentimes they contain a considerable quantity of the mother liquor inclosed between their crystal layers. Hence if purity of product is the sole requisite, it is often more desirable to obtain a meal of very fine crystals. Such a meal is obtained by crystallizing rapidly and stirring while crystallizing. Some substances are so difficult to obtain in large crystals that it is more satisfactory to try only to obtain a uniform crystal meal. 1 8 NOTES ON LABORATORY MANIPULATION () Purification by Recrystallization. When a given substance crystallizes from a solution, it most generally separates in a pure condition irrespective of any other dis- solved substances the solution may contain. Thus a sub- stance can be obtained in an approximate state of purity by a single crystallization. Portions of the mother liquor (containing dissolved impurities) are, however, usually en- trapped between the layers of the single crystals, not to mention the liquid which adheres to the crystal surfaces. By dissolving the crystals, the small amount of impurity likewise passes into the solution, but only a small fraction of this impurity is later entrapped by the crystals when they separate from this new mother liquor. By several recrys- tallizations, then, a substance can ordinarily be obtained m a very high state of purity. 9. DRYING (a) A preparation that is not affected by the atmosphere can be dried by being spread in a thin layer on a plate of glass, on filter paper, or, best of all, on an unglazed porous porcelain plate. 1 The substance may with advan- tage be turned over occasionally with a spatula ; and if it is not decomposed by heat it can be dried more rapidly in a warm place, such as over the steam table. A product con- taining water of crystallization should never be dried at an elevated temperature. During the drying the preparation must, of course, be carefully protected from dust. (V) Substances which decompose on standing exposed to the air may be quickly dried if they are first rinsed with alcohol, or with alcohol and then ether. Rinsing with alcohol re- moves nearly all of the adhering water, and a further rinsing 1 Dishes which are imperfect, and on that account have not been glazed, can be obtained very cheaply from the factories. PULVERIZING 19 with ether removes the alcohol. Alcohol evaporates more rapidly than water, but ether evaporates so rapidly that a preparation wet with it may be dried by a very few minutes' exposure to the air. Alcohol and ether are both expensive and should be used sparingly. They can be used most effectively as fol- lows : After all the water possible has been drained from the preparation, transfer the latter to an evaporating dish and pour over it enough alcohol to thoroughly moisten it; stir it with a spatula until the alcohol has penetrated to every space between the crystal grains, then pour off, or drain off, the alcohol and treat the preparation in like man- ner with another portion of fresh alcohol. After that wash it once or twice with ether in exactly the same manner. If it is necessary to wash the preparation on the filter, drain off the water as thoroughly as possible, stop the suction, add just enough alcohol to moisten the whole mass, and after letting it stand a few moments drain off the liquid completely. Apply a second portion of alcohol and portions of ether in the same manner. 10. PULVERIZING In chemical reactions in which solid substances are involved the action is limited to the surface of the solid, and for this reason it is evident that it must be much slower than reactions which take place between dissolved substances ; it is also evident that the more finely powdered a solid substance, the greater is its surface, and therefore the more rapidly it will react. Most solid raw materials for the following preparations are supplied in the powdered form, but they are rarely powdered finely enough ; they should in general be further pulverized until they no longer feel gritty beneath the pestle or between the fingers. 2O NOTES ON LABORATORY MANIPULATION For grinding any quantity of a substance a large porce- lain mortar (say 8 inches in diameter) with a heavy pestle is preferable to the small mortars usually supplied in the desks. One or more such mortars is placed in the laboratory for general use. If a hard substance can be obtained only in large pieces, it should first be broken with a hammer, then crushed into small particles in an iron or steel mortar, after which it is to be ground in the porcelain mortar. In the final grinding it is often advisable to sift the fairly fine from the coarser particles, then to finish grinding the former by itself and to crush and grind the coarser particles apart. DRY REACTIONS: FURNACES Dry solid substances do not react appreciably with each other at ordinary temperature. Reactions are made possible in two ways : first, the wet way, in which the substances are dissolved and thus brought into most intimate contact. In many cases solution also produces ionization, which, as is known, greatly increases chemical activity. Reactions in the dry way are rendered possible by heat. Heat alone increases the rapidity of a chemical reaction, it being a general law that the speed is increased from two to three times for every increase of 10 C. in temperature. In cases in which one or more of the reacting substances are melted by the heat, the same sort of intimate contact is brought about as in the case of solutions. Fusion is likewise a means of producing electrolytic dissociation, and on this account also it increases chemical activity. In some of the furnace reactions in which none of the substances are melted, as, for example, in the reduction of strontium sulphate to strontium sulphide by means of char- coal (see Preparation No. 6), the process probably takes FURNACES 21 place in virtue of a certain amount of gas which is con- tinuously regenerated. A little of the hot charcoal is oxidized to carbon monoxide, which then reduces some of the strontium sulphate, it being itself changed to carbon dioxide thereby ; the latter gas comes in contact with incandescent charcoal, and carbon monoxide is again produced. Reactions in the dry way are usually carried out in crucibles of iron, clay, or graphite, according as to which is least attacked by the reagents. For rather moderate tem- peratures the crucible may be heated over a flame, but in most cases the requisite temperature can best be obtained in a furnace. The form of furnace most to be recommended for this work is repre- sented in Figure 5. It consists of a cylinder of FIG. 5 fire clay, 7 inches high and 6^ inches in external diameter, which is surrounded by a sheet iron casing. It is heated, as shown, by a gas- wind flame, introduced through an opening in the lower part of one side. If a suitable air blast is not available, a gasoline blowpipe (such as is commonly used by plumbers) is almost equally serviceable. When such a furnace as that described is heated as hot as possible with a well-regulated mixture of gas and air, a temperature of about 1,350 can be obtained. For carrying out ordinary chemical preparation work an accurate enough measure of the temperature is given by the color of the glowing interior of the furnace, and the approximate cen- 22 NOTES ON LABORATORY MANIPULATION tigrade values corresponding to different colors are as follows : Incipient red heat . . . . '55 Dull red heat 650 Red heat 800 Bright red heat 1,000 Yellow heat ...... 1,200 White heat i>35o CHAPTER I ALKALI AND ALKALINE EARTH METALS These metals constitute the left hand or A families in the first two groups of the periodic classification of the elements, as shown in the table which appears in the ap- pendix, and which is also placed for convenience inside the front cover of the book. The metals of these two families are studied together because they are the extremely active base-forming elements. On account of their great activity they are never found uncombined in nature, and it is only by the aid of the most powerful reducing agencies (for example, by electroly- sis of their molten salts) that the metals themselves are extracted from their compounds. The alkali metals are monovalent. Their hydroxides, MOH, are extremely soluble and are highly dissociated as bases ; on account of the corrosive properties of the latter they are known as the caustic alkalies hence the designa- tion, alkali metals. The compounds of the alkali metals are, with a very few exceptions, soluble in water, and they are all strong electrolytes. The radical ammonium, NH 4 , is classed with the alkali metals on account of its ability to form the same kinds of compounds. The alkaline earth metals are divalent ; their hydroxides, M(OH) 2 , are less soluble than those of the alkali metals, but are nevertheless very strongly basic. The compounds of these metals are not so generally soluble as those of the alkali metals, and in particular the carbonates and sulphates are mostly insoluble. 2 3 I. POTASSIUM NITRATE FROM SODIUM NITRATE AND POTASSIUM CHLORIDE The most important source of nitrates is Chili saltpeter, sodium nitrate. This is not suited for use in explosives ^n account of its property of attracting moisture and rendering the explosive preparation damp. Potassium nitrate is not open to this objection, and hence large quantities of it are prepared, using sodium nitrate as a source of the nitrate radical. When two ionizable salts are dissolved in water the resulting solution will contain, besides the undissociated molecules and the ions of these two salts, also the undis- sociated molecules of the two new salts which form by the interaction of the ions present. Which of these four salts will crystallize first out of solution depends upon their relative solubility. Thus if sodium nitrate and potassium chloride are dissolved together in water the resulting solu- tion will contain Na + , K + , NO 8 ~, and Cl~-ions, together with undissociated molecules of NaNO 8 , NaCl, KNO 3 , KC1. The solubility of some salts varies very much with the temperature, while that of other salts varies very little. This is seen from the following table and diagram, and practical use is made of these facts by crystallizing succes- sively the two different salts at different temperatures. GRAMS OF SALT SOLUBLE IN 100 GRAMS OF WATER At io b At 100 KNO.3 21 246 NaCl 36 40 KC1 31 56 NaNO 3 8 1 1 80 2 5 7U 26 ALKALI AND ALKALINE EARTH METALS KN0 3 50 Temperature/ 100 Procedure, Dissolve 100 grams of sodium nitrate and 88 grams of potassium chloride in 200 cc. of water and evaporate in a porcelain dish to half that volume. Without letting the liquor cool, separate it from the crystals which have formed during the evaporation. This is best accom- plished with the -aid of suction (see Note 4 (b) on page 7, Witt filter). The liquid is poured through the filter and then the crystals are thrown upon the plate and pressed with a spatula, while applying gentle suction in order to remove as much as possible of the liquid clinging to them. Pour the filtrate into a beaker and set it aside to cool ; then examine the crystals left on the filter and convince yourself POTASSIUM NITRATE 2 7 that they consist in the main of sodium chloride. (Examine with a microscope. The crystals should be cubical. Com- pare the taste with that of pure sodium chloride and that of pure potassium nitrate.) By means of running tap water cool the filtrate to about 10, and then separate the crys- tals of potassium nitrate from the liquid in the same man- ner as above (see Note 3 on page 7, last sentence). The filtrate from these crystals is saturated with both sodium chloride and potassium nitrate, and the larger part of the latter should be saved. Evaporate the solution in a smaller dish until a considerable quantity of sodium chloride crys- tallizes from the boiling liquid. Filter hot, as above, and crystallize potassium nitrate from the filtrate by cooling. Unite this crop of potassium nitrate crystals with the first. Test a very small portion of them for sodium chloride by dissolving about o.i gram in 2 cc. of water and adding a drop of silver nitrate solution. They are not pure and must be purified. Weigh roughly the crystals while still moist, and dissolve them in from half to three-quarters of their weight of hot water. Cool and separate the crystals from the mother liquor. The latter should now contain nearly all of the sodium chloride which was mixed with the first crop of crystals. Test as above to see if this crop is free from sodium chloride. If not, repeat the recrystallization as many times as is . necessary to get a perfectly pure product. A little of this should when dissolved give no turbidity with silver nitrate solution, and when held in the flame on a platinum wire should color it the violet color of potassium, with none of the yellow sodium color. Spread the preparation on an unglazed porcelain plate and allow it to dry by standing exposed to the air; then put up the salt in a test tube or a small bottle, and label it neatly. Questions i. Define metathesis, 28 ALKALI AND ALKALINE EARTH METALS 2. When a metathetical reaction is carried out in the wet way, why is the solubjlity of the substances involved of importance? Explain why, according to this point of view, the reactions AgNO 3 -\- KC1 = AgCl -f KNO 3 and BaCl 2 -f" Na 2 SO 4 = BaSO 4 -|- 2NaCl are much more com- plete than the reaction NaNO 3 + KC1 = KNO 3 -f NaCl. 3. Explain why fewer operations should be required to prepare potassium nitrate from potassium sulphate and barium nitrate than by the foregoing procedure. 2. CAUSTIC POTASH FROM WOOD ASHES Of the mineral constituents of plants, potassium salts form an important part, and, so far as these are salts of organic acids, they are converted into potassium carbonate when the plant is burned. On an average, wood ashes contain about 10 per cent, of potassium carbonate, and before the advent of the Leblanc Soda Process this was almost the sole supply of alkali. Even after this process came into general use, by which sodium carbonate could be obtained from common salt, wood ashes remained for some time the important source of potassium carbonate. In recent years, however, the greater part of the production of potassium carbonate has been derived by the Leblanc Process from potassium salts found in deposits in the earth, principally at Stassfurt, Germany. Potassium carbonate being the principal soluble con- stituent of wood ashes, it is extracted with water; but the extract so obtained contains, as well, the other soluble mineral constituents, and also a considerable amount of tarry coloring matter which was not destroyed in the com- bustion of the wood. This tarry matter is destroyed by calcination of the residue obtained on evaporating the aqueous extract, and the calcined mass is what is known POTASH FROM WOOD ASHES 29 as crude potash. A better grade of commercial potash can be obtained by dissolving this mass in water, filtering, and evaporating the solution. In order to obtain potassium hydroxide or caustic potash from potassium carbonate the solution of the latter is treated with milk of lime (calcium hydroxide). With this it interacts, yielding insoluble calcium carbonate and soluble potassium hydroxide, Ca(OH) 2 + K 2 C0 3 = CaC0 3 + 2 KOH. Procedure. Tie a piece of cloth over the mouth of a thistle tube and insert it beneath the surface of a layer of sand, one-half inch deep, in the bottom of a tall 2-liter bottle. Mix i kilogram of wood ashes with 800 cc. of hot water in a pail, and transfer the moist mass to the bottle. Pour 200 cc. more of hot water over the surface of the ashes. Connect the thistle tube with a siphon and draw off as much liquid as possible, perhaps 200 cc., using suction and drawing the liquid into a suction bottle if it does not otherwise run rapidly enough. Commence evap- orating this liquid in an 8-inch evaporating dish ; then pour 300 cc. of hot water on top of the ashes and stir around the surface layer. Again draw off about 300 cc. of liquid and add it to that in the evaporating dish, and repeat the operation until the liquid drawn off is nearly colorless. Not more than 2 to 2\ liters need be drawn off in all. When the liquid is all evaporated remove the dry residue to an iron dish and heat it strongly with a Bunsen burner to destroy the tarry matter. More than a moderate red heat should not be exceeded, as it is not desirable to fuse the salt, but the heating should be continued until the ash is white or at most contains only black specks of completely charred carbon. When cooled, weigh the material ; assum- ing it to consist wholly of potassium carbonate, calculate 3<3 ALKALI AND ALKALINE EARTH METALS the amount of quicklime (calcium oxide) necessary to react with it according to the equation, CaO + H 2 + K 2 C0 3 = 2KOH + CaCO 3 . Slake 20 per cent, more than that amount of lime by covering it in a porcelain dish with water and quickly pouring off the excess of water. If the lime is of suitable quality it will soon grow hot and crumble to a powder, Ca(OH) 2 . Take sufficient water to make ten times the weight of the crude potash. Stir up the slaked lime with half of it, thus making milk of lime. Dissolve the potash in the other half, bring it to boiling, and add the milk of lime with stirring. Let the mixture boil for 15 minutes and then filter, using a suction bottle (see Note 4 (), Witt filter). Measure the volume of the solution of caustic potash obtained and preserve it in a stoppered bottle. Test the strength of the solution. Measure 15 cc. with a pipette into a beaker and add a drop of litmus solution. Run into this from a burette a solution of normal hydro- chloric acid (36.5 grams per liter), drop by drop, until the color just changes from blue to red. If the right point is overstepped begin again with a fresh sample of the solu- tion. From the amount of acid taken to neutralize the sample, calculate the amount of KOH obtained from the wood ashes. Preserve the solution in a bottle labeled with the number of cubic centimeters of the solution, with its strength in mols per liter of potassium hydroxide, and with the actual amount in grams of the potassium hydroxide. Questions i. The calcium hydroxide used in causticizing the potash is a slightly soluble solid suspended in water, its solubility being 1.7 grams 'per liter. Explain how, in spite of its limited solubility, the required amount can enter into reaction, SODA BY THE AMMONIA PROCESS 3 I 2. Explain why the caustic potash solution obtained contains practically no calcium ions, even in case an excess of calcium hydroxide may have been used for causticizing. 3. Analyses of the crude potash obtained from various grades of wood ashes have given results which fall within the limits given in the table: K 2 CO 3 . , . 38-78 per cent. Na 2 CO 3 . . . 0-12 per cent. K 2 SO 4 . . . 13.5-40.5 per cent. KC1 . . . 0.9-10.0 per cent. Insoluble matter . 0.1-9.2 per cent. What substances other than KOH would you expect then to be present in the caustic potash solution which you have prepared ? Look up the solubility of calcium sulphate and calcium hydroxide, and decide whether, in the presence of a large amount of potassium hydroxide, milk of lime would react with a small amount of potassium sulphate according to the equation : K 2 S0 4 + Ca(OH) 2 = CaS0 4 + 2 KOH. 4. If a solid substance crystallizes from the caustic potash solution after it has stood, decide what it is by consulting the table in Question 3 and a solubility table. 3. SODIUM CARBONATE BY THE AMMONIA PROCESS The principle employed in the manufacture of sodium carbonate from sodium chloride by the Solvay Process is exceedingly simple. It depends primarily upon the fact that sodium acid carbonate is but sparingly soluble in water ; this compound is produced by the interaction of so- 32 ALKALI AND ALKALINE EARTH METALS dium chloride, in a saturated salt solution, with ammonium acid carbonate, (i) Nad + NH 4 HCO 3 ^ NaHCO 3 + NH 4 C1. Since ammonia and salts of ammonium are very much more expensive than sodium carbonate, it is evident that the process can be of no commercial value unless ammonia can be recovered and used again. This is accomplished in fact by treating the mother liquor, after separation from the sodium bicarbonate, with calcium hydroxide, (2) Ca(OH) 2 -f 2 NH 4 C1 = CaCl 2 + 2 NH 3 + 2 H 2 O. In practice the process is usually carried out as follows : A nearly saturated salt solution is purified of iron, mag- nesia, lime, etc., which would otherwise get into the final product; there is then passed into it ammonia gas until it has absorbed 60 to 70 grams of NH 3 per liter, whereupon carbon dioxide is passed in until it has reacted with the ammonia to form ammonium bicarbonate, NH 4 OH -f- C0 2 = NH 4 HC0 3 , which in turn reacts with the salt according to (i). Exten- sive precautions necessarily have to be observed that prac- tically no ammonia escape during the process, so that the entire amount may be used over and over again. It is also essential that as little carbon dioxide shall be wasted as possible. Thus the carbon dioxide is utilized which is produced in converting limestone into quicklime, CaCO 3 -> CaO + CO 2 , and in converting sodium bicarbonate into sodium carbonate, 2 NaHC0 3 - Na 2 C0 3 + H 2 O + CO 2 . Procedure. To 50 cc. of concentrated ammonium hy- droxide (sp. gr. 0.90) add 150 cc. of water. Place in a flask SODA BY THE AMMONIA PROCESS 33 and add 60 grams of table salt free from lumps. Shake until the salt is nearly or quite dissolved and filter the solution if it is not perfectly clear. Pass a delivery tube through one hole of a double-boreal, tightly fitting stopper placed in a 300 cc. flask. Provide a plug for the other hole. Let the tube dip into the solution which is placed in the flask, and pass in carbon dioxide gas from a Kipp generator until all the air has been displaced from the flask ; then close the flask and allow the gas to pass in as fast as it will be absorbed. Occasionally, as the action seems to slacken, loosen the plug for a moment. Shake the flask frequently. It will take several hours for the solution to absorb sufficient carbon dioxide, and it may be left over night connected with the generator. When no more gas can be absorbed pour the mixture from the flask upon a Witt filter (see Note 4 ()). Apply suction to remove the liquid from the sodium bicarbonate. Wash the product three times with 15 cc. of ice water (see Note 5 (a), on page 12), sucking it free from liquid each time. Spread the preparation on an unglazed plate and leave it until it ceases to smell of ammonia. Test the preparation for chlo- rides (see Question i), of which it should not contain more than a trace. Questions 1. In testing the preparation for chlorides, the test solution must be acidified with nitric acid before silver nitrate is added. What other silver salt would otherwise be precipitated ? How does the presence of nitric acid prevent this ? 2. How is sodium carbonate prepared from sodium bicarbonate ? 3. What is an acid salt ? How does a solution of an acid salt such as KHSO 4 behave towards litmus ? Test the 34 ALKALI AND ALKALINE EARTH METALS behavior of solutions of NaHCO 3 and of Na 2 CO 8 towards litmus. Explain the cause of this behavior. 4. Why cannot potassium carbonate be prepared from potassium chloride by the ammonia process? (Look up the solubility of potassium bicarbonate.) What process may be used to obtain potassium carbonate from this source ? 4. CHEMICALLY PURE SODIUM CHLORIDE FROM ROCK SALT Common rock salt may contain other than sodium chlo- ride up to 10 per cent, of matter, which consists in the main of the sulphates and chlorides of potassium, calcium, and magnesium, not to mention a considerable amount of dirt and insoluble matter. For most commercial purposes these impurities are not harmful. By careful crystallization of the salt from solution, a product sufficiently free from these impurities can be obtained to be used as table salt. To obtain chemically pure sodium chloride, however, more elaborate precautions must be taken. A satisfactory method depends upon the insolubility of sodium chloride in a con- centrated solution of hydrochloric acid. A nearly saturated solution of the rock salt is prepared, and, without removing the dirt and insoluble matter, enough pure sodium carbonate is added to precipitate the calcium and magnesium in the solution as carbonates. Into the clear filtrate is then passed gaseous hydrochloric acid until the greater part of the sodium chloride is precipitated, while the small amounts of sulphates and of potassium salts remain in the solution. The precipitate is drained and washed with a solution of hydrochloric acid until the liquid clinging to the crystals is entirely free from sulphates. Procedure. Dissolve 25 grams of rock salt in 75 cc. of water, hastening the action with gentle heating. To the PURE SODIUM CHLORIDE 35 solution add about i gram of sodium carbonate dissolved in a few cubic centimeters of water. Stir, let settle, and add a few drops more of sodium carbonate solution, and if no fresh precipitate is produced in the clear part of the FIG. 6 solution no more need be added ; otherwise enough more must be added to produce this result. Filter the solution, hot, through an ordinary filter (Note 4 (V)). Prepare pure gase- ous hydrochloric acid as follows : Place 50 grams of dry rock 36 ALKALI AND ALKALINE EARTH METALS salt in a round-bottom liter flask provided with a rubber stopper with two holes, through which pass a thistle tube reaching to the bottom of the flask, and an exit tube just coming through the stopper. Provide a wash bottle for the gas as follows: A 300 cc. bottle is provided with a stopper with three holes. Through one passes a glass tube reaching to the bottom for entrance of the gas; through another a thistle tube reaching to the bottom for use as a safety tube; and through the third an exit tube just coming through the stopper. Pour into the bottle enough concentrated hydrochloric acid (sp. gr. 1.2) to come up three-quarters of an inch on the two lower tubes. Connect the exit tube with a 2-inch filter funnel for delivering the gas into the sodium chloride solution. Use entirely glass tubing, and, where connections must be made with rubber, bring the ends of the glass tubes close together. Before commencing to use this apparatus it must be approved by an instructor. Pour the sodium chloride solution into a beaker of 3 inches diameter, and insert the mouth of the funnel below the surface of the liquid. Pour gradually 95 cc. of concentrated sulphuric acid into the generating flask, and when the first action has ceased warm very gently. There will be a great deal of frothing, but the froth should at no time be allowed to get over into the wash bottle. 1 When hydrochloric acid gas ceases to be gen- erated, separate the precipitated sodium chloride from the mother liquor by pouring it upon a Witt filter. Suck the moisture from the crystals. Test the filtrate for sul- phate by adding a little barium chloride solution to a small sample of it diluted with water. A strong test will probably be obtained. Now wash the crystals with succes- 1 The melted sodium acid sulphate left in the generating flask is very hot and must not be poured into the sink. It may be poured into some dry receptacle specially provided, or it may be allowed to cool slowly and solidify in the flask and then be dissolved out with water. AMMONIUM BROMIDE 37 sive portions of 10 cc. of hydrochloric acid solution of 1.12 sp. gr. until the washings show no further test for sulphates. (See Note 5 (#).) Then transfer the crystals to a porcelain dish and heat gently, while stirring, until all decrepitation ceases. Questions 1 . Why must the hydrochloric acid gas be passed through a washing bottle ? Why is the safety tube necessary ? 2. Why, in the light of the Mass Law, should one expect the solubility of sodium chloride to be lessened by the pres- ence of hydrochloric acid ? [It may be stated that another effect also comes into play here which likewise tends to lessen the solubility of sodium chloride.] 3. Mention two possible causes for the very consider- able amount of heat produced when the hydrochloric acid gas is absorbed by the solution in the beaker. 5. AMMONIUM BROMIDE Ammonium bromide could be prepared by the neutral- ization of ammonium hydroxide with hydrobromic acid, NH 4 OH + HBr = NH 4 Br -f- H 2 O. Since, however, hydrobromic acid is a more expensive material than uncombined bromine, the latter would have the preference as a source of bromine, provided it yielded as satisfactory a product. Chlorine or bromine reacts as follows upon a cold solution of sodium hydroxide, as, for example, in the manufacture of bleaching liquors, Br 2 + 2 NaOH = NaBr + NaBrO + H 2 O, with the formation of sodium hypochlorite or hypobromite. Sodium hypobromite reacts with ammonium hydroxide ac- cording to the equation, 3 NaBrO + 2 NH 4 OH = 3 NaBr + 5 H 2 O + N 2 . 38 ALKALI AND ALKALINE EARTH METALS Thus the action of bromine upon ammonium hydroxide yields only ammonium bromide and nitrogen gas, because even if the primary effect were to yield bromide and hypobromite in equal quantities, as is the case if sodium hydroxide is used, the ammonium hypobromite would imme- diately react with fresh ammonia in the same manner as does sodium hypobromite. Procedure. Place 55 cc. of concentrated ammonia (sp. gr. 0.90), together with 50 cc. of water, in a flask, which should be set in a pan of ice water. Put 15.8 cc. of bromine in a small separatory funnel, 1 and add it a drop at a time to the ammonia, rotating the flask after each drop until the yellow color produced by the bromine has completely disap- peared. Do not allow the contents of the flask to become heated at any time, as a dangerously explosive compound might in that case be formed. As soon as a permanent yel- low color is produced, stop adding bromine and add at once a few drops of ammonia until the solution has again become colorless. Place the solution in an evaporating dish on top of a beaker of boiling water, and let the salt crystallize as the water evaporates. When only a little liquid remains, separate the crystals from it in a funnel, and dry them on a porcelain plate. Questions 1. What products would be formed if bromine were added to a solution of sodium hydroxide instead of am- monium hydroxide, (i) if the solution were kept cold ? (2) if it were heated? 2. Experiment. Add about 10 drops of bromine to x The stopcock of the funnel should first be lubricated with vaseline or grease and then fastened with a rubber band. Bromine produces very bad burns when it gets upon the hands. To avoid danger of accident ask an instructor to approve the apparatus before beginning actual operations with the bromine. STRONTIUM HYDROXIDE 39 10 cc. of a cold 10 per cent, sodium hydroxide solution. Add this gradually to a solution of ammonium hydroxide, made by diluting i cc. of desk reagent with 10 cc. of water. Determine what gas is given off. 3. Write the equation showing the complete reaction between bromine and ammonia. What fraction of the entire amount of ammonia used is lost through formation of nitrogen gas ? 4. Why cannot hydrobromic acid be prepared from potassium bromide by a method analogous to that used in the manufacture of hydrochloric acid ? 5. Explain why, from the standpoint of economy, the method of preparation above outlined is superior to the direct neutralization of ammonia with hydrobromic acid. 6. STRONTIUM HYDROXIDE FROM STRONTIUM SULPHATE One of the most important sources of strontium is the mineral celestite, SrSO 4 . By reduction with charcoal this can be converted into strontium sulphide, SrSO 4 + 4 C = SrS + 4 CO, and the strontium sulphide by treatment with copper oxide and water can be made to yield strontium hydroxide, SrS + CuO + H 2 O = Sr(OH) 2 + CuS. Copper oxide is in the ordinary sense insoluble ; neverthe- less in contact with water it does yield to an infinitesimal extent, first copper hydroxide, and then Cu ++ -ions, CuO -]- H 2 O ^ Cu(OH) 2 ^ Cu ++ + 2OH-. Therefore, since copper sulphide is a far more insoluble substance than copper oxide, it follows that the few 4O ALKALI AND ALKALINE EARTH METALS Cu ++ -ions from the latter unite with the S~~-ions from the strontium sulphide to form copper sulphide, which pre- cipitates continuously, while the copper oxide continuously goes into solution to resupply Cu ++ -ions, and this action continues until either the copper oxide or the strontium sulphide is exhausted. Strontium hydroxide crystallizes with 8 molecules of water, Sr(OH) 2 .8H 2 O. It is very soluble in hot water, but sparingly soluble in cold water. Procedure. Grind 50 grams of powdered celestite in a porcelain mortar until no more grit is felt under the pestle. Add 20 grams of powdered charcoal and continue to grind with the pestle until the two are thoroughly mixed. Place the mixture in a clay crucible, pack it firmly, and cover it with a layer of powdered charcoal \ inch deep. Cover the crucible with a close-fitting cover and heat it in a gas furnace for one hour, at a bright red heat (Instruc- tions as to regulating the flame). After the contents of the crucible have cooled, remove the layer of charcoal from the surface and bring the remainder, after crushing it to a powder, into an 8-inch porcelain dish ; add 300 cc. of water, bring the mixture to a boil, and while it is boiling add copper oxide, a little at a time, until all of the soluble sulphide has interacted with it, about 40 grams in all. So long as any unchanged strontium sulphide is present the solution will show a yellow color, which may be observed by letting the black solid settle for a moment, and then looking through the upper layers of the clear liquid at the background of the white porcelain dish. As soon as the yellow color has entirely disappeared, the strontium sulphide has all reacted. Crystals of strontium hydroxide separate rapidly from this solution when it cools. Hence it must be filtered quickly in order to avoid having the crystals form in the filter and clog it completely. Heat 50 cc. of STRONTIUM HYDROXIDE 41 water to boiling in a beaker, and keep it at this temperature until it is .required. Add hot water to the dish to replace any lost by evaporation, and pour (Note 2) the hot solution through a Witt filter (Note 4 ()), allowing the main part of the residue to remain in the dish. Add the 50 cc. of hot water to this residue, stir it thoroughly, heating it for a moment over the flame, and then pour solution and residue into the filter and drain out all of the liquid. Transfer the solution to an Erlenmeyer flask (or let it remain in the filter bottle), stopper the flask to exclude the air, and wrap it with a towel, so that the solution -may cool slowly and larger crystals may be formed. Finally, after several hours cool the solution with running tap water and then collect the crystals on a Witt filter. Drain the crystals for a moment, but do not draw too much air through them, as they retain all the carbon dioxide it contains. Spread the moist product on filter paper, and allow it to dry as quickly as possible by contact with the air. Stopper it in a sample bottle or tube as soon as it is dry. Questions 1. What constituent of the atmosphere must be ex- cluded from the solution while crystallizing and as much as possible from the crystals while drying? How would it contaminate the preparation? 2. A sample of the preparation should dissolve nearly clear in hot water. What will surely cause a slight cloudi- ness? 3. How could strontium chloride be prepared from strontium sulphide ? 4. Give some other method by which strontium hydrox- ide could be obtained from strontium sulphide without the use of copper oxide. 5. Starting with the mineral strontium carbonate, how 42 ALKALI AND ALKALINE EARTH METALS might strontium hydroxide be prepared ? Strontium oxide ? Strontium chloride ? 7. STRONTIUM CHLORIDE FROM STRONTIUM SULPHATE Strontium chloride might be prepared by treating strontium sulphide, the intermediate product in the last preparation, with hydrochloric acid. For the sake of illus- trating another method, however, a process which does not require the use of a furnace is here employed for decom- posing the strontium sulphate. The method consists in first converting the sulphate into the carbonate by boiling it with a concentrated solution of sodium carbonate, and then of dissolving the carbonate in hydrochloric acid, thereby yielding a solution of the chloride. The conversion of solid strontium sulphate into solid strontium carbonate furnishes an interesting illustration of the principle of mass action, for the solubility of these two salts in pure water is as follows : Solubility in grams Solubility in mols per liter per liter SrSO 4 o.on 0.0006 SrCO 3 o.oon 0*00007 Strontium sulphate would dissolve in the solution of sodium carbonate in the same manner as it would in pure water until it had saturated the solution and its solubility product, which is equal to 0.0006 X 0.0006, was reached, but for the fact that long before this could occur the solution would be supersaturated with respect to strontium carbon- ate, whose solubility product is only equal to 0.00007 X 0.00007. Thus strontium carbonate is precipitated con- tinuously as strontium sulphate dissolves ; and since the STRONTIUM CHLORIDE 43 solution cannot become saturated with the latter so long as there is a large excess of carbonate ions present, the solid salt finally remaining will consist entirely of strontium carbonate, provided a sufficient amount of sodium carbon- ate were employed. The reaction which takes place is, however, reversible, SrSO 4 + Na 2 CO 3 ^ SrCO 3 -f Na 2 SO 4 , and, if strontium carbonate were boiled with a solution of sodium sulphate, the solid would be converted into the sulphate. It is easy to deduce that if the ratio of the con- centration of the ions in solution ^ is greater cone SO 4 ~~ than T L, solid strontium sulphate will be converted into solid carbonate. Procedure. Take 50 grams of powdered celestite. Grind it in a mortar until it is so fine that it no longer feels gritty under the pestle. Cover it in an 8-inch dish with 300 cc. of water, add 60 grams of anhydrous sodium carbonate, and boil the mixture for 30 minutes, stirring it constantly at first. Transfer the solution and solid to a tall, narrow beaker, using 100 cc. of fresh water in rinsing out the last of the residue, and let the solid matter settle for 5 minutes. Decant off the liquid, which is still some- what cloudy, but from which the essential part of the solid has settled, and wash the residue three times by decantation with 400-500 cc. of water (see Note 5 (#), first paragraph). The residue is now sufficiently free from soluble sodium sulphate. Transfer about y 1 ^ of the moist strontium car- bonate to another beaker, to be used in a later part of the process. To the remaining T 9 3H 2 S -j- As 2 S 5 . Sodium sulphantimonate can be prepared from stibnite by the combined action of a solution of sodium sulphide and sulphur: 28 -f- Sb 2 S 3 - Sb 2 S 5 ; 3 Na 2 S + Sb 2 S 5 -* 2 Na 3 SbS 4 , and it crystallizes well with nine molecules of water. Procedure. To 50 grams of powdered stibnite, 107 grams of crystallized sodium sulphide (NasS.QHaO), 1 and 10 grams of powdered sulphur in a porcelain dish, add 100 cc. of water, bring to a boil, and keep at the boiling temperature for \ of an hour. Filter and rinse the residue in the dish and on the filter with hot water, bringing up the volume of the solution to 200 cc. While still hot put it away in a covered dish, with a towel placed over it, to crystallize. Drain the crystals ; evaporate the mother liquor somewhat to obtain a second crop of crystals. Spread the crystals on a porous plate, and stopper them tightly as soon as dry. Questions i. Experiment. Prepare a little precipitated antimo- nous sulphide. How? Treat this precipitate with a solu- tion of ammonium polysulphide, (NH 4 ) 2 S X . Discuss, with reactions, the nature of the soluble compound produced. Finally, acidify the solution with hydrochloric acid. What is the reaction? 1 Or use 35 grams of anhydrous sodium sulphide and an additional 72 cc. of water. IIO ELEMENTS OF THE FIFTH GROUP 30. ANTIMONY PENTASULPHIDE, Sb 2 S 5 This compound cannot be prepared directly from the trisulphide and sulphur, because it is decomposed at a temperature below that at which the latter substances would react. As has just* been seen, however, the higher sulpho-salt of antimony can be readily prepared in the wet way ; and this, on decomposition with a dilute acid, yields antimony pentasulphide. This substance is much used in vulcanizing rubber. Procedure. Dissolve 40 grams of the sodium sulphan- timonate obtained in the last preparation, and dilute with i liter of cold water. Mix 15 cc. of concentrated sulphuric acid with 350 cc. of water, cool, and place in a 2-liter, or, better, a 3 or 4-liter common bottle. To this add slowly, and with constant stirring, the solution prepared above. Fill the bottle with water and stir thoroughly. Let the precipitate settle, draw off the liquid, and wash by decan- tation until the wash water no longer gives, with barium chloride, the test for a sulphate. After the last washing let settle for some time, draw off as much as possible of the clear liquid, and transfer the slime to a large filter (Note 4 (c] ; do not omit to reenforce the point of the filter) to drain for 12 hours or longer. Without removing the pasty antimony sulphide from the filter, open out the latter on an unglazed plate, and leave it on a shelf above the steam table where the temperature does not rise above 50. When the product is completely dry, detach the hardened lumps from the paper and pulverize them in a mortar. Questions i. Write all the reactions involved in the preparation of antimony pentasulphide from stibnite. METALLIC ANTIMONY I I I 31. METALLIC ANTIMONY This metal is obtained on a commercial scale both by reducing antimony oxide with carbon and by reducing antimony sulphide by means of metallic iron. The latter method possesses the advantage that antimony sulphide, a natural product, is used directly and does not need to be first converted into the oxide. The iron sulphide formed by this method is fusible and forms a slag ; but the slag is made more fusible by the admixture of sodium sulphate as directed, and thus the globules of melted antimony are allowed to sink more easily to the bottom of the crucible and form a metallic regulus. The slag furthermore covers the surface of the metal and hinders its volatilization and oxidation. Procedure. Mix 100 grams of stibnite with 42 grams of iron filings, 10 grams of anhydrous sodium sulphate, and 2 grams of powdered charcoal, and place the mixture in a clay crucible. Cover the crucible tightly, and heat it in the gas furnace for one hour at a bright red heat. The temperature should not be high enough to volatilize the antimony, which would in that case escape as a white smoke consisting of antimony oxide, yet the slag of iron sulphide must be completely softened, although it should not melt to a thin liquid. After about half an hour test the conditions by removing the cover a moment and stirring the slag with an iron rod to see whether it is in the proper semi-fluid condition. After the reaction is complete, allow the crucible to cool, break it and separate the regulus of antimony from the slag. Questions i. Experiment. Warm a piece of metallic antimony with hydrochloric acid. Where does antimony stand in the electromotive series? 112 ELEMENTS OF THE FIFTH GROUP 2. Experiment. Warm a little powdered antimony with nitric acid in a casserole. What is the product? Compare it with the product obtained by treating tin in the same manner. 32. BISMUTH BASIC NITRATE (BISMUTH SUBNITRATE) Although bismuth is the most strongly metallic element of the fifth group, yet its salts in aqueous solution undergo partial hydrolysis very readily. In presence of a consider- able amount of free acid, the Bi + + + -ion is capable of exist- ence in solution; but with decreasing quantities of acid the tendency of water to produce hydrolysis increases, and the basic salt of bismuth, which is only slightly soluble, separates : OH _}- 2 H + + 2NO,f. On pouring a solution of bismuth nitrate into a considerable quantity of cold water the basic nitrate, according to the above formula, is precipitated. This salt, however, is not stable in contact with a solution which does not contain nitric acid of a concentration of at least about \ molal, but slowly changes over into some other more basic nitrate, and if washed repeatedly with pure water will finally go over completely into the hydroxide : Bi OH + H 2 O ^ Bi OH -f H + + NO-T . Under the conditions given in the following procedure, this production of a more basic salt will occur if the precipitate is allowed to stand in contact with the~solution for a con- siderable time; hence the directions to filter at once. BISMUTH BASIC NITRATE I I 3 The basic nitrate is by no means completely insoluble in water, and the filtrate contains considerable quantities of bismuth, which can be conveniently saved as oxide by precipitating with sodium carbonate. Procedure. Dissolve without heating 40 grams of crys- tallized bismuth nitrate, Bi(NO 8 ) 3 .5H 2 O, in 10 cc. of nitric acid (sp. gr. 1.2) and 20 cc. of water. Pour this into 2 liters of cold water and stir thoroughly for a few minutes. Let the precipitate settle completely, and as soon as this has occurred draw off and save the supernatant liquor; drain the precipitate on a suction filter, and wash it quickly with about 20 cc. of water. Dry the precipitate at the steam table, and preserve it as a powder. Bismuth Oxide. Combine all the liquors from the fore- going; add sodium carbonate until alkaline to litmus; let settle, and draw off the supernatant liquor ; boil the remain- ing suspension after adding to it about 20 grams more of sodium carbonate. Then wash the precipitate twice by decantation, drain on a suction filter, and wash with two or three portions of water. Dry and preserve. Questions 1. In accordance with the above directions, sodium carbonate is used to precipitate bismuth hydroxide. Why should not the precipitate be bismuth carbonate ? 2. If this precipitate is not finally boiled with an excess of sodium carbonate, it is likely to contain a certain amount of basic nitrate. Explain why this should be so and why the boiling will convert it completely into the hydroxide. 114 ELEMENTS OF THE FIFTH GROUP GENERAL QUESTIONS. V ELEMENTS OF THE FIFTH GROUP OF THE PERIODIC SYSTEM Experiments (The results observed are to be recorded in the laboratory note- book at the time the experiments are performed.) 1. Boil about \ gram of powdered metallic arsenic with nitric acid (1.2) until the metal is entirely dissolved. Evap- orate the solution just to dryness by heating it over a free flame in a casserole while holding the latter in the hand and rotating its contents. In this way all the unused nitric acid is expelled. When cool add 10 cc. of water and warm until the arsenic acid has all dissolved. Prove that the solution contains an acid (/. 3, instead of 40 grams of chromite. POTASSIUM BICHROMATE 121 potassium nitrate. Place the mixture in a cast-iron crucible, which it must on no account fill more than two-thirds full, else when melted it will run over. Heat in a gas furnace to a white heat (but using care not to reach the very highest heat, which might melt the crucible) until the melted charge has ceased to effervesce. Pour the molten mass out onto a dry 1 iron plate. When cool crack it up and dissolve it, together with what still adheres to the crucible, in boiling water. Filter the solution, and extract the residue with a little more boiling water and pour through the same filter. Add glacial acetic acid (cautiously) to the filtrate until it has become acid. Boil down the solution to 300 cc., or to even a less volume if no solid salt begins to separate. Add 25 cc. more of glacial acetic acid, let stand for some time, and finally cool to o before separating the crystal meal of potassium bichromate from the mother liquor. Purify the product by recrystallization. Questions 1. Mention at least three oxidizing agents which might have been used instead of potassium nitrate in this preparation. 2. How might the oxidation of chromic hydroxide, Cr(OH) 8 , be accomplished in the wet way? Experiment. To 5 cc. of a cold solution of a chromic salt add about i gram of sodium peroxide, agitate for a few moments, and then warm until effervescence ceases. Formulate the equations for the intermediate reactions in such a way as to show the state of oxidation of chromium in each com- pound involved, and then add the separate equations to give one for the complete reaction. 3. Experiment. To a solution of potassium bichro- 1 See footnote 2, page 60. 122 HEAVY METALS mate add potassium carbonate until no more effervescence occurs. Observe and explain any change in color. To a solution of potassium chromate add an acid and observe as before. Explain fully the relation between chromates and bichromates. 34. POTASSIUM CHROMATE FROM POTASSIUM BICHROMATE Dissolve 50 grams of potassium bichromate in water and add the calculated amount of potassium carbonate dissolved in water. The color should just change to clear yellow, and no trace should be left of the reddish hue characteristic of the bichromate. Crystallize the product from the solution (see solubility table on page 120). Answer the questions given under Potassium Bichromate. 35. CHROMIC ANHYDRIDE, CrO 8 When a chromate or a bichromate is treated with a strong acid, chromic acid is formed in the solution. The affinity of chromic anhydride for water is far less than that of sulphuric anhydride for water ; and chromic acid, there- fore, instead of existing in solution entirely in the form H 2 CrO 4 , is broken down to a great extent into H 2 Cr 2 O 7 (i.e., H 2 O.2CrO 3 ) and even to CrO 8 . Especially in the presence of a large amount of sulphuric acid, the last form is produced so freely that it crystallizes out in the shape of red needles. Commercially, chromic anhydride is most often prepared by the action of sulphuric acid on potassium bichromate. Potassium acid sulphate is first crystallized from the mixture and after that the chromic anhydride; but a good deal of care is necessary to obtain the product uncontaminated with CHROMIC ANHYDRIDE 123 potassium salt. When, however, lead chromate is used as the source of the chromic acid, the lead can be completely removed, since with sulphuric acid the extremely insoluble lead sulphate is formed. The remaining solution then con- tains nothing but chromic acid and an excess of sulphuric acid. The chromic anhydride can then be crystallized out, it being least soluble in a mixture containing in the neigh- borhood of 75 per cent, of sulphuric acid. Procedure. Lead Chromate. Dissolve 100 grams of lead acetate in i liter of water, and add a few drops of acetic acid if necessary to clear up any turbidity. Dissolve 39 grams of potassium bichromate in i liter of water, and add this solution to the first, while stirring. Wash the pre- cipitate by decantation until less than 0.5 per cent, of the soluble salt remains (Note 5 (b) on page 13); then collect the lead chromate on an ordinary filter (Note 4 (<:)), and after draining dry it thoroughly. Chromic Anhydride. Take the lead chromate prepared above (or 100 grams of a commercial sample), reduce all lumps to a fine powder, add 200 grams of concentrated sul- phuric acid (in cc.), and stir the mixture with a pestle until a perfectly smooth paste is produced. Allow the mixture to digest 24 hours in a warm place, as on the shelf above the steam table. Dilute to a liter with water, and filter the solu- tion through asbestos felt (Note 4 (d) on page 10). Wash the lead sulphate until it is nearly or quite white, separate it from the asbestos as completely as possible, and preserve it as a by-product. Evaporate the filtrate in a porcelain dish until crystals of chromic anhydride begin to form a scum on the surface of the liquid. Let the solution cool slowly ; then collect the crystals in a funnel in which a per- forated plate or a glass marble is placed, and drain out all the liquid with suction. Cover the funnel with a watch glass while evaporating the liquor to obtain a second crop 124 HEAVY METALS of crystals. Collect this crop in the same funnel together with the first crop, and wash the product twice with con- centrated nitric acid to remove the sulphuric acid adhering. Use sufficient nitric acid each time to just wet the entire mass of crystals, and then drain it off as thoroughly as pos- sible with the suction (Note 5 (a) on page 12). Finally, drive off the nitric acid by heating the crystals very cau- tiously in a small porcelain dish placed on a sand bath. Keep turning over the mass of crystals with a glass spatula to avoid heating the lower layer too strongly. Chromic anhydride melts at 192, and care must be taken to avoid reaching this temperature. When the crystals appear dry and no more white vapor can be detected on breathing across them, preserve them in a glass-stoppered bottle. Questions 1. When chromic anhydride is dissolved in water, what components are produced in the solution ? What salt is precipitated if lead acetate or barium acetate is added to this solution? Why is it not the bichromate which is obtained ? 2. Experiment. Heat a little chromic anhydride strongly on a bit of porcelain. What color change occurs ? (The color of the product can be better observed if a par- ticle is pulverized in a white mortar.) Is the product soluble in water ? In hydrochloric acid ? What relation does it bear to the mineral chromite? 36. CHROMIC ALUM The preparation of potassium bichromate illustrated how chromic oxide, Cr 2 O 3 , as it exists in nature as a constituent of the mineral chromite, can be oxidized to a chromate in which chromium exists as CrO 3 . For the preparation of CHROMIC ALUM 125 chromic alum, K 2 SO 4 .Cr 2 (SO^) B .24H 2 O, it might seem as if chromite should yield chromic sulphate directly on treat- ment with sulphuric acid. This is, however, impossible, because the natural material is, as already stated, very resistant to the action of acids,. It yields only to the action of powerful oxidizing agents, which convert it into a chro- mate, and therefore potassium, or sodium, bichromates are always the products made directly from the mineral, and these serve as the materials from which other compounds of chromium are prepared. To make chromic alum from potassium bichromate it is necessary to reduce the chro- mium to the same state of oxidation in which it originally existed in the mineral, and to add sufficient sulphuric acid to form the sulphates of potassium and chromium. Alcohol may be used as the reducing agent, it being itself oxidized to aldehyde, a body whose presence is made very evident by its penetrating odor. Chromic alum is isomorphous with common alum and can easily be obtained in large and beautiful deep purple crystals. Care must, however, be exercised not to allow the temperature of its solution to rise above 50 during the preparation, for when heated beyond this point it under- goes a change into a green noncrystallizable body. This green body is not stable at the ordinary temperature, and after cooling it will change slowly back into the ordinary crystallizable chromic alum ; but so slowly, however, that if once it is formed the preparation is practically spoiled. At 25, 24 grams of K 2 SO 4 .Cr 2 (SO 4 ) 3 .24H 2 O will dissolve in 100 grams of water, and the solubility increases very rapidly with the temperature. Procedure. Pulverize 100 grams of potassium bichro- mate, and cover it in an 8-inch evaporating dish with 400 cc. of water. Add 78 cc. of concentrated sulphuric acid, and stir until the bichromate is all dissolved. Adding 126 HEAVY METALS the sulphuric acid should produce enough heat to dissolve the bichromate, but if it is necessary heat the mixture a little more. Be sure that the last trace of solid is dissolved. Allow the solution to cool to 40 ; then add alcohol, a drop at a time, while stirring constantly with the stem of a ther- mometer until the temperature commences to rise. Then place the dish in a pan of ice and water and add alcohol, 65 cc. in all, at first very slowly, endeavoring to keep the temperature between 35 and 40, and finally more rapidly. Keep the temperature at all times well below 50; and if it should start to rise suddenly, due to too large an addition of alcohol, and get as high as 50, drop a piece of ice directly into the solution. Finally, let the solution cool completely in the bath of ice water, or, still better, let it stand over night. Collect the crystal meal on a Witt filter and suck it free from liquid. Recrystallize so as to obtain large, well-shaped crystals, following a similar procedure and observing the same precautions as with common alum (see page 54). A saturated solution of this salt should be .pre- pared at 35. After freeing it of any undissolved particles of the crystal meal, warm it to 40, and set it to crystallize, with the addition of about ten very small crystals to serve as nuclei. Dry with filter paper the crystals so obtained, and stopper them at once in a bottle, since they are quite efflorescent. Questions 1. Formulate the equations for the separate reactions involved in the reduction of the bichromate, in such a way as to show the changes occurring in the state of oxida- tion of the chromium ; the alcohol, C 2 H 6 O, is oxidized to aldehyde, C 2 H 4 O. 2. Sulphur dioxide might serve as the reducing agent. Give equations for the partial and complete reactions. CHROMIUM 127 3. Dissolve -J- gram of potassium bichromate in 10 cc. of water and add 10 cc. of dilute sulphuric acid. Heat to boiling, and pass in hydrogen sulphide until the color is changed completely to green. To what is the green color due ? What is the precipitate ? Formulate equations also for this reaction. 37. CHROMIUM METAL BY THE GOLDSCHMIDT PROCESS The readiest method of obtaining the metal chromium from its oxide, and one which yields it in a high state of purity, is the so-called Goldschmidt Process, in which use is made of metallic aluminum as the reducing agent according to the reaction, 2A1 + Cr 2 O 3 = A1 2 O 3 -f- 2Cr. The heat produced by the oxidation of aluminum is so great that it is sufficient to effect the decomposition of the chromic oxide with still enough surplus heat to produce a tempera- ture high enough to melt the metallic chromium. It is evident that before this reaction can be made to progress spontaneously a sufficient temperature must be developed to decompose the chromium oxide. This necessary tem- perature is a good deal higher than that of a Bunsen flame or of a common furnace, but can be obtained by use of the fuse powder described below. When once started in this way the reaction itself produces a temperature high enough to insure its continuance. Carried out on the small scale of a laboratory prepa- ration, the heat produced is not quite sufficient to melt the metal and slag so thoroughly that the metal can settle out to form a compact regulus at the bottom of the crucible. By adding a small amount of potassium bichromate to the 128 HEAVY METALS charge, however, the reaction becomes more energetic, owing to the more available supply of oxygen. Procedure. Heat some powdered chromic oxide in an iron pan over a Bunsen burner. Melt some potassium bichromate in a clean iron pan and pulverize it in a mortar after it has again solidified. It is necessary for the mate- rials used to be entirely free from moisture. Mix 210 grams of the ignited chromic oxide, 60 grams of the fused potas- sium bichromate, and 96 grains of granulated aluminum (not the powder which is used for a pigment), pack the mixture closely into a clay crucible, and embed the latter in a pail of sand. Make a hole about 4 cm. deep in the middle of the charge, and fill it with about 10 grams of a fuse powder made from 10 parts of barium peroxide and i part of granulated aluminum. Insert a strip of mag- nesium ribbon into the fuse powder. Place the whole under the hood at a distance from any woodwork, and start the reaction by igniting the end of the magnesium ribbon with a Bunsen flame. It is advisable for the operator to wear colored glasses while watching the reaction, and to keep at a little distance to be out of the way of flying sparks. When the crucible has cooled, break it and sepa- rate the regulus of metallic chromium from the slag of fused aluminum oxide. 38. MANGANESE CHLORIDE FROM WASTE MANGANESE LIQUORS The waste liquors left after the generation of chlorine from manganese dioxide and hydrochloric acid contain principally manganous chloride. Besides this, however, there is always some free acid and almost always a con- siderable amount of ferric chloride present. The greater part of the free acid can be removed by evaporating the MANGANESE CHLORIDE solution until a pasty mass is left which will solidify on cooling. The iron can be removed from the solution of this residue in virtue of the ease with which ferric salts hydrolyze. The nearly neutral solution is treated with suspended manganous carbonate (obtained by treating a part of the solution itself with a soluble carbonate). Ferric chloride hydrolyzes according to the reversible reaction, FeCl 3 + 3 H 2 O ^ 3 HC1 + Fe(OH) 8 . In the -presence of manganous carbonate the small amount of free acid thus formed is continuously used up according to the reaction, MnCO 3 + 2HC1 -* MnCl 2 + H 2 O + CO 2 . Thus the re- action of hydrolysis is enabled to run to completion. The remaining solution, which is almost absolutely neutral and entirely free from iron salts, yields crystallized manganous chloride, MnCl 2 .4H 2 O, upon evaporation. * Procedure. Boil 500 cc. of waste manganese liquor in a 6-inch evaporating dish under the hood until the residue becomes pasty. After a scum begins to form on the surface of the liquid, there is danger of spattering and the mixture should be stirred with a glass rod until it becomes semi- solid. Heat the residue to boiling with 1,000 cc. of water- without filtering the solution obtained, take one-tenth of it, dilute this portion to 1,000 cc., and add a solution of sodium carbonate to it until all of the manganese is precipitated as carbonate (test for complete precipitation). Transfer the pre- cipitate to a tall, common bottle and wash it by decantation at least four times. Add the slime of manganous carbonate to the remaining nine-tenths of the manganous chloride solution, and boil the mixture in a casserole until a few drops of the filtered liquid give no red color when tested with potassium sulphocyanate. Filter the solution and evaporate it in an 8-inch dish until a crystal scum forms on blowing across the surface. Then allow the solution to cool slowly and crystallize, leaving it for at least 12 hours I3O HEAVY METALS uncovered in a place protected from dust Collect the crystals and evaporate the mother liquor to obtain further crops of crystals until practically all of the salt has crystal- lized. 1 Spread the light pink crystals on an unglazed plate to dry. Questions 1. Explain the purpose of the test with potassium sulphocyanate. 2. Explain the action of manganese dioxide in the generation of chlorine gas from hydrochloric acid. In what state of oxidation does manganese exist in the salt manganous chloride ? 3. If iron were in the ferrous condition, it would not be removed from the solution by the above procedure. Explain why iron is necessarily in the ferric condition in the liquors used. 4. Experiment. Dissolve a small grain of manganous chloride in a half test tube of water. Test the solution with hydrogen sulphide ; then add a few drops of ammonia, and if necessary pass in a little more hydrogen sulphide. Then add acetic acid (a weak acid) until the solution is again faintly acid. Does the manganous sulphide dissolve ? Com- pare the solubility of manganous sulphide with that of copper sulphide ; of zinc sulphide. 5. In testing for the complete precipitation of iron from the manganese chloride solution, what would be the effect observed on adding ammonium sulphide (a) before, and (b) after, all the iron has been precipitated ? J The crystals of manganous chloride are deliquescent when the temperature is low and the atmosphere charged with moisture. If the product cannot be obtained satisfactorily by the above directions, carry out the crystallization and drying in a place at a slightly elevated temperature, 25-30; or cool the saturated hot solution rapidly by stirring or shaking, and dry the crystal meal so obtained by rinsing it with alcohol and then letting the latter evaporate rapidly. POTASSIUM PERMANGANATE 13! 6. Explain how facts involved in the foregoing prepa- ration show that Mn(OH) 2 is more strongly basic than Fe(OH) 3 . 39. POTASSIUM PERMANGANATE FROM MANGANESE DIOXIDE Although manganese dioxide is a powerful oxidizing agent, it is nevertheless capable of being itself oxidized when it is fused with a basic flux. The trioxide of manga- nese is acidic in nature and combines with the base to form a salt. Thus it is evident that the presence of a base favors the oxidation. The dioxide of manganese is neither strongly basic nor acidic in nature and shows no marked tendency to form salts. The monoxide is distinctly basic and the trioxide is distinctly acidic, so that the former forms salts with acids and the latter with bases. It follows, therefore, that in the presence of acids the dioxide has a tendency to produce salts of man- ganous oxide whereby an atom of oxygen is set free (see No. 38, Manganous Chloride), and that in the presence of bases manganese dioxide has a tendency to take on another atom of oxygen in order to produce a salt of the trioxide. Thus when manganese dioxide is fused with potassium hydroxide and an oxidizing agent, the salt potassium man- ganate is formed. This salt is soluble in water and is fairly stable so long as a considerable excess of potassium hydrox- ide is present ; but in presence of an acid even so weak a one as carbonic acid the manganate decomposes spon- taneously, two-thirds being oxidized to permanganate at the expense of the other one-third, which is reduced again to manganese dioxide : 3 H 2 MnO 4 -* 2HMnO 4 + MnO 2 -f- 2H 2 O. The permanganate (or permanganic acid) corresponds to 132 HEAVY METALC the heptoxide of manganese, Mn 2 O 7 , which is the most strongly acid-forming of the oxides of manganese. Perman- ganic acid is a strong and very soluble acid, it being of approximately the same acid strength as nitric or hydro- chloric acids. It is in addition a very powerful oxidizing agent. Procedure. Place 50 grams of potassium hydroxide and 25 grams of potassium chlorate in an 8 cm. sheet iron crucible. Heat the mixture carefully until it is just melted. Meantime grind 50 grams of pyrolusite to as fine a powder as possible (the finer it is ground, the more successful the preparation). Remove the flame from under the crucible and add the pyrolusite, a little at a time, stirring vigorously with an iron spatula (an old file with a wooden handle) all the while. 1 After all is added, place a small flame below the crucible, and keep stirring the charge. Gradually in- crease the strength of the flame, and stir continuously until the mass stiffens completely. Then cover the crucible and heat it 5 minutes longer at a dull red heat. When the mass has cooled, place crucible and all in i liter of water in an 8-inch porcelain dish. After the solid- has entirely disin- tegrated, remove the crucible and rinse it off with a little water from the wash bottle. Boil the solution in the dish, and at the same time pass in carbon dioxide generated from marble and hydrochloric acid until the green color of the manganate has entirely changed to the violet-red color of the permanganate. Test the color by touching a drop of the solution on a stirring rod to a piece of filter paper. If the spot is violet with no trace of green and only a fleck of brown manganese dioxide in the center, the change to 1 Since the charge in the crucible effervesces and spatters particles of melted salt, great care should be taken to keep the eyes at a safe distance. The hand holding the stirrer should be protected with a thick glove or with a towel, and with the other hand the crucible should be held firmly by means of long iron tongs. POTASSIUM PERMANGANATE 133 permanganate is complete. Remove the lamp ; let the sludge settle in the dish for 5 minutes ; then pour the solution through an asbestos filter (see Note 4 (//) on page 10), being careful to avoid stirring up the sludge until the very last, since the slimy precipitate of manganese dioxide would so clog the filter as to nearly stop the flow. Lastly, with the aid of a jet of water from the wash bottle, transfer all the sludge to the filter and drain it free from liquid. Evaporate the solution in a clean dish to a volume of 300 cc. Let it settle a moment and filter it through asbestos as before. Pour the filtrate into a 6-inch dish, and allow it to cool slowly in a place protected from the dust. When cold, col- lect the crystals of potassium permanganate on a perforated plate placed loosely in a filter funnel. Evaporate the mother liquor to 100 cc., filter it through asbestos, and obtain a second crop of crystals. Discard the remaining liquid, since it cannot contain more than about 6 grams of potassium permanganate and to evaporate it further would cause po- tassium chloride also to crystallize out. Weigh all the crys- tals, dissolve them in eight times their weight of water (to give a saturated solution at about 40), filter the solution through asbestos at near the boiling temperature, and let it cool slowly and crystallize in a small porcelain dish cov- ered with a watch glass. Recover another crop of crystals in the same way from the mother liquor, after evaporating it to a volume of 60 cc. Allow the crystals to dry on a clean unglazed plate. Questions 1. Name and give the symbols of all the oxides of manganese. 2. From which oxide is K 2 MnO 4 derived ? KMnO 4 ? 3. Write the reactions involved in the above preparation. 4. How could KMnO 4 be converted back into K 2 MnO 4 ? Reaction ? 134 HEAVY METALS 5. Does it frequently happen that, with an element which can exist in several states of oxidation, a compound derived from one oxide is stable in an alka'line solution but unstable in an acid solution, while in the latter solution the compound derived from another oxide is the stable one ? What other preparation besides the present one illustrates this point? 40. MANGANESE METAL BY THE GOLDSCHMIDT PROCESS The principle of the production of manganese by this process is exactly the same as that of the production of chromium in Exercise 33. On account of the violence of the reaction between the oxide of manganese and aluminum it is not advisable to ignite the whole charge at once in the crucible ; yet on account of the high melting point of man- ganese a considerable quantity of charge must be used in order to produce heat enough to obtain the metal melted together in a uniform lump, instead of distributed in small globules throughout the mass of the slag. Before mixing up the charge, the pyrolusite which is used must be first heated by itself in order to drive off any water which it may contain and to convert it to the lower oxide, Mn 2 O 3 . Procedure. Place i kilogram of finely powdered pyro- lusite in a Hessian crucible and heat to a bright heat in a gas furnace. To prepare the charge, mix 750 grams of this material, when it is cooled sufficiently, with 250 grams of granulated aluminum. Heat the empty crucible again in the furnace, and while still hot place it in a pail of sand, as in the preparation of chromium. Place about 20 grams of the charge in the bottom of the hot crucible. Put on colored glasses and a heavy glove ; start the reac- tion with a little fuse powder and a magnesium ribbon (see FERROUS AMMONIUM SULPHATE 135 Chromium), and then add fresh portions of the charge rapidly but without allowing the reaction to become too violent. When the crucible has cooled, break it, and sepa- rate the regulus of metallic manganese from the slag of fused aluminum oxide. Questions 1. If pyrolusite containing water were used without previous heating, what disadvantage would result during the process? 2. What economy of materials is effected by converting the manganese dioxide into the lower oxide ? 41- FERROUS AMMONIUM SULPHATE FeSO 4 .(NH 4 ) 2 SO 4 .6H 2 Corresponding to the two most important oxides of iron, FeO and Fe 2 O 3 , the two sulphates, FeSO 4 and Fe 2 (SO 4 ) 8 , can be prepared. By dissolving iron in sulphuric acid a solution of ferrous sulphate is obtained. This, however, is readily oxidizable, slowly even by the oxygen of the air, to the higher sulphate, and ferrous sulphate can only be preserved free from ferric salt when all oxygen is excluded, or when it is kept in contact with an excess of metallic iron in an acidified solution. Dry crystallized ferrous sul- phate or green vitriol, FeSO^yHoO, can be preserved fairly well without becoming oxidized ; but the double ferrous and ammonium sulphate is not only more easily prepared on account of the readiness with which it crystallizes, but it is also much less easily oxidized by contact with the air. Procedure. Prepare crystallized ferrous ammonium sul- phate from equimolal quantities of crystallized ferrous sulphate and ammonium sulphate, using 70 grams of the former and 136 HEAVY METALS 33 grams of the latter. In crystallizing the product, observe the solubilities given in the following table, and make use of suggestions given in Note 8, page 15 and under Alum, page 54: A saturated solution contains for each 100 grams of water the given number of grams of the anhydrous salt. 10 20 30 40 50 70 90 FeSO 4 16 26 33 44 48 56 43 (NtDoSO.. 71 73 75 78 81 84 92 99 FeSO 4 (NH 4 ) 2 SO 4 12 17 22 28 33 40 52 Questions i. Dissolve a little of the preparation in water and test it with potassium ferrocyanide. If the precipitate is white or only a pale blue, of what does it consist ? If it is deep blue, what is shown? 42. FERRIC AMMONIUM ALUM In this preparation ferrous sulphate is converted into ferric sulphate under the oxidizing action of nitric acid in the presence of the amount of sulphuric acid theoretically necessary to form this salt. By the addition of ammo- nium sulphate the double salt, ferric ammonium sulphate, Fe 2 (SO 4 ) 8 ,(NH 4 ) 2 "SO 4 .24H 2 O, crystallizes, this being one of the isomorphous series of alums (see Alum). At 25, 100 grams of water dissolve 44 grams of the anhydrous or 124 grams of the hydrated ferric ammonium sulphate. Procedure. Heat together 100 grams of crystallized ferrous sulphate, 100 cc. of water, and 12 cc. of concen- trated sulphuric acid until the salt is dissolved. While the FERRIC ALUM 137 solution is boiling, add concentrated nitric acid, a little at a time, until the iron is completely oxidized to ferric sul- phate and a few drops of the solution diluted with a few cubic centimeters of water give no blue precipitate with potassium ferricyanide. Evaporate the solution until it is thick and sticky and most of the excess of nitric acid has been driven off. Dissolve this in water, making up to a volume of 125 cc. ; heat to boiling and add 25 grams of ammonium sulphate dissolved in 100 cc. of hot water. Allow the solution to cool slowly and crystallize. Collect the crystals in a funnel ; wash with a very little water and allow to dry on an unglazed plate. Obtain a second crop of crystals from the mother liquor. Questions 1. Write the reaction involved in the oxidation of ferrous sulphate as carried out in this preparation. If an unacidified solution of ferrous sulphate is oxidized by the oxygen of the air, what products are formed ? 2. Write the reaction involved in the test for ferrous salt with potassium ferricyanide. 3. Experiment. Prepare a solution of a ferrous salt by dissolving 2 grams of ferrous ammonium sulphate in 20 cc. of water, adding a little dilute sulphuric acid and a piece of iron wire. Test both this solution and a solution of a ferric salt (nitrate or chloride) with potassium ferrocya- nide, potassium ferricyanide, and potassium sulphocyanate. Tabulate the results. These constitute the standard tests for ferrous and ferric salts. 138 HEAVY METALS GENERAL QUESTIONS. VI HEAVY METALS OF THE SIXTH, SEVENTH, AND EIGHTH GROUPS OF THE PERIODIC SYSTEM Experiments (The results observed are to be recorded in the laboratory note- book at the time the experiments are performed.) 1. Test the stability of nickel carbonate by heating i gram of it gently in a test tube while shaking it in a Bunsen flame. Test the gas evolved for carbon dioxide; and compare the action of the remaining solid, when treated with hydrochloric acid, with that of the original carbonate. The carbonates of divalent iron, cobalt, manganese, and chromium are all of approximately the same degree of stability as nickel carbonate, so that this one experiment may be taken as typical of this class of carbonates. 2. To Show Whether the Carbonate of a Trivalent Metal Can Exist. Dissolve 2 grams of ferric alum in i o cc. of water (this gives a trivalent iron salt in a solution that con- tains no free acid). Add a 10 per cent, sodium carbonate solution slowly until no more action takes place. What is the gas evolved ? What is the precipitate ? In this ex- periment the ions Fe + + + and CO 3 ~ ~ are brought together; the other ions, Na + and SO 4 ~ ~, could not react together to give any visible effect. If, therefore, ferric carbonate is stable in contact with water, it will either form a precipitate if it is insoluble, or if it is soluble it will simply stay in solution and no effect will be observable. The gas given off shows that the carbonate is unstable. Write the equation for the reaction. None of the carbonates of the metals of this group, when they are in the trivalent condition, are any more stable than ferric carbonate. A salt of chromium, such as chromic alum, might be used instead of ferric alum in the above experiment. GENERAL QUESTIONS VI 139 3. Oxidation of a Divalent Oxide. Heat \ gram of cobalt carbonate in an open porcelain dish, holding the dish with tongs and keeping it rotating in the flame, but not allowing the porcelain to even approach a visible red heat. Heat until the color of the cobalt carbonate has completely changed. Like nickel carbonate, cobalt car- bonate is decomposed by heat into cobaltous oxide, CoO, and carbon dioxide. If the cobaltous oxide is readily oxidized by the oxygen of the air, it may at once be changed into Co 2 O 3 or Co 3 O 4 . To test for this, treat a little of the product with hydrochloric acid in a test tube. From CoO the salt CoCl 2 would be obtained. From Co 2 O 8 or Co 3 O 4 the same salt, CoCl 2 , would be obtained not CoCl 3 and chlorine would be liberated (action similar to that of MnO 2 with hydrochloric acid). Test for the forma- tion of chlorine by means of the odor or by using iodide starch paper. NOTE : In Experiment i a higher oxide of nickel was probably formed in the same manner, although to a considerably less extent. 4. Acidify solutions of potassium permanganate and potassium bichromate each with sulphuric acid, warm them and treat them with sulphur dioxide (sulphurous acid), and note any change in color. The change is due to the reduc- tion of the given salts which are derived from the oxides Mn 2 O 7 and CrO 3 , respectively, to salts derived from the oxides MnO and Cr 2 O 3 , respectively. Questions In the group of elements discussed under General Questions, I, changes of valence do not occur, but the metals of the alkali and alka- line earth families show the same valence in all their compounds. Proceeding in the order in which the elements have been taken up in this book, a constantly increasing tendency is observed for the ele- ments to display different valences, until in the group under consid- eration the most important chemical characteristics of the elements I4O HEAVY METALS are connected with their changes from one state of valency to another. NOTE : The terms state of valency and state of oxidation can in most cases be used interchangeably. 1. In which groups of the periodic system do the elements chromium, manganese, iron, nickel, and cobalt fall ? What is peculiar about the position of the last three ? What other metals belong to the same family as chromium ? In what relation do they stand to sulphur, selenium, and tellurium ? In what relation does manganese stand to the halogens ? What other elements occur in the eighth group in triads similar to iron, nickel, and cobalt ? 2. How do the monoxides of chromium, manganese, iron, cobalt, and nickel compare in basic strength with the oxides of copper and zinc and with the oxides of the alkali and alkaline earth metals ? How do the sesquioxides, R 2 O 3 , compare with the monoxides of this group as regards basic strength (see Experiment 2) ? What is true as regards the base- or acid-forming prop- erties of the oxides higher than the sesquioxides, e.g., of CrO 3 , MnO 3 , Mn 2 O 7 ? 3. Give the symbols and names of salts derived from each of the three oxides of chromium, CrO, Cr. 2 O 3 , CrO 3 . In which of its compounds does chromium most resemble sulphur ? iron and aluminum ? nickel, cobalt, copper, and zinc ? 4. Give the symbols and names of salts derived from each of the oxides of manganese, MnO, Mn 2 O 3 , MnO 2 , MnO 3 , Mn 2 O 7 . In which of its compounds does manganese most resemble chlorine ? aluminum ? cobalt, nickel, copper, and zinc ? sulphur ? lead in the dioxide ? 5. Formulate the reaction between sulphurous acid and potassium permanganate in acid solution (see Experiment 4). By means of partial equations resolve the compounds into their constituent oxides ; show the simple oxidation and GENERAL QUESTIONS VI 14! reduction, and then show how the new oxides combine to give the salts which actually result. Then add the separate equations to give the total equation for the complete reaction. 6. In the same manner formulate the reaction between sulphurous acid and potassium bichromate. CHAPTER VII NON-METALLIC ELEMENTS OF THE SIXTH AND SEVENTH GROUPS OF THE PERIODIC SYSTEM The elements which are distinctly and invariably non- metallic in character are boron in the third group, carbon and silicon in the fourth group, nitrogen and phosphorus in the fifth group, oxygen, sulphur, selenium, and tellurium in the sixth group, and fluorine, chlorine, bromine, and iodine in the seventh group. It has been assumed that these elements, or at least the most important of them, were studied before entering on the course of study outlined in this book. Indeed no knowledge of the chemistry of the metallic elements would be possible without a certain knowl- edge of the non-metallic elements with which they form compounds. By turning to the table of the periodic arrangement of the elements, it is at once seen that the non-metals do not occur at all in the first and second groups ; that they occur only at the top in the third, fourth, and fifth groups ; and that in the sixth and seventh groups they comprise all the members of the B families. It is true in these families, as might be expected by recalling characteristics of preceding groups, that the strength of the non-metallic character grows weaker and that the approach towards metallic character grows more evident, as the atomic weight increases ; indeed it is probable that if the elements which should fit into the places below tellurium and iodine, respectively, are ever found, they will display quite as marked metallic properties as non-metallic. H3 144 NON-METALLIC ELEMENTS The characteristic valences of the sixth and seventh groups are VI and VII, respectively, and the corresponding oxides are EO 8 and E 2 O 7 . In these oxides and in the compounds derived from them, there is little dissimilarity between the A and B families. Thus perchlorates and permanganates are in every way analogous to each other, as are also sulphates and chromates. In any lower state of valence, the elements of the B families are entirely dif- ferent from those of the A families, and the most striking non-metallic properties of the former are exhibited in their ability to combine directly with metallic elements, forming oxides, sulphides, chlorides, bromides, etc. It will be noticed that the negative valence, that is, the valence exhibited towards hydrogen or a metal, is I for the non-metals of the seventh group and II for the sixth group. 43. POTASSIUM IODIDE Of the two most obvious possibilities for making this salt, the direct synthesis from the elements is quite imprac- ticable, both because potassium metal is expensive and because the action would be violent and difficult to control. The method of neutralizing hydriodic acid with potassium hydroxide presents no chemical difficulties, but the materials are more costly than the iodine and potassium carbonate which are employed according to the following procedure (compare with the preparation of ammonium bromide, No. 5). Iodine, when brought together in the presence of water with an excess of iron, reacts to form soluble ferrous iodide. By treating this solution with potassium carbonate a metath- esis takes place, yielding insoluble ferrous carbonate and soluble potassium iodide. The desired product should then be obtained by filtration, except that the ferrous carbonate forms a slimy, bulky precipitate which clogs the filter; but this difficulty may be overcome if the ferrous salt is partially oxidized by adding more iodine before throwing out the iron, and the mixture of ferrous carbonate and ferric hydrox- ide that is obtained in this way will filter very readily. Procedure. Place 7 or 8 grams of iron filings and jo cc. of water in an Erlenmeyer flask and add 25 grams of iodine, a small portion at a time, while shaking continu- ously. When all is added warm the mixture somewhat until ail of the iodine has combined, and the dark brown color of free iodine gives place to the yellow of ferrous iodide. Filter off the excess of iron and add 5 grams more of iodine to the solution. Warm the solution until the iodine is dis- solved, and then pour it into a boiling solution of 17 grams of potassium carbonate in 50 cc. of water in a good-sized MS 146 NON-METALLIC ELEMENTS flask. Warm the solution, which is at first very thick with the- gelatinous precipitate, until the latter becomes more compact and the mixture thus becomes more fluid. Filter a little of the liquid ; it should be colorless and should con- tain no iron; otherwise a little more potassium carbonate should be added. Filter the whole solution and wash the precipitate with hot water to save all the soluble salt. Evaporate the filtrate to a small volume in a porcelain dish, filter again if necessary, and evaporate further until crystals begin to form. Then leave the solution to evaporate spon- taneously in a place protected from dust (best in a somewhat warm place). Collect the crystals and recover another crop from the mother liquor. Questions 1. What proportion of the ferrous compound is oxidized to ferric by the addition of the 5 grams of iodine to the filtered solution in the foregoing procedure ? 2. On treating the mixture of ferrous iodide and iodine with potassium carbonate, state reasons why ferrous carbon- ate rather than ferrous hydroxide and why ferric hydroxide rather than ferric carbonate should precipitate. 3. Starting with iodine and potassium hydroxide, devise a process for preparing potassium iodide without the use of iron or similar metal. Compare No. 5, Ammonium Bromide, and No. 46, Potassium Bromide and Potassium Bromate. 44. HYDRIODIC ACID The direct synthesis of hydrogen iodide from the ele- ments is impracticable, because the chemical affinity between hydrogen and iodine is so small that at any temperature sufficiently elevated to make them react at all, they would combine only very incompletely. By the interaction of an HYDRIODIC ACID 147 iodide with a non-volatile acid, such as sulphuric acid, hydrogen iodide gas could, it is true, be formed : KI -f H 2 S0 4 - KHS0 4 + HI ; but the hydrogen iodide so formed acts as a reducing agent upon the sulphuric acid, whereby it is itself oxidized to free iodine and water. Hydrogen iodide, then, cannot be satisfactorily prepared by the direct union of hydrogen and iodine nor by the me- tathesis of an iodide with the non-volatile sulphuric acid. The most convenient method of preparing it is by means of the action of hydrogen sulphide with iodine in aqueous solution : H 2 S+ I 2 -> 2HI -f S. The affinity between hydrogen and sulphur is very small, as is also that between hydrogen and iodine ; but this fact, combined with the fact that free iodine is appreciably solu- ble while free sulphur is exceedingly insoluble in water, makes it possible for the reaction to run to completion. The rather dilute solution of hydriodic acid which is obtained in this manner can be concentrated by distillation ; at first nearly pure water passes off, then the- quantity of acid in the distillate increases until an acid of specific gravity 1.7, containing 57 per cent, of HI, comes over. At this point the distillate has the same composition as the residual liquid, and the remainder of the acid can be dis- tilled with a constant composition. By this method it is not possible to obtain an acid of higher concentration, but by allowing this acid to absorb hydrogen iodide gas until it is saturated at o, an acid of specific gravity 2.0, containing 90 per cent, of HI, can be obtained. Hydrogen iodide gas can be prepared by the interaction of red phosphorus and iodine with a little water. Procedure. Grind 30 grams of iodine to a fine powder 148 NON-METALLIC ELEMENTS and add gram of it to 100 cc. of water in a small Erlen- meyer flask. Pass hydrogen sulphide into the solution until the brown color of iodine has disappeared. Add about i gram more of iodine, and again pass hydrogen sulphide until the iodine is used up. After 10 grams of iodine have reacted in this way, add the remaining 20 grams and allow the mixture to stand, with repeated shaking, until the iodine is entirely dissolved (half an hour or more). Then pass hydrogen sulphide slowly until the solution is decolor- ized. Pour the solution into another flask, leaving the clotted lumps of sulphur behind, and rinse the first flask and the residue with a few cc. of water. Pass a current of carbon dioxide through the solution until the excess of hydrogen sulphide is entirely removed ; then shake the flask vigor- ously to cause the suspended sulphur to clot together, and filter the solution. In this way a rather weak solution of hydriodic acid is obtained. Fit a distilling flask with a thermometer and an inlet tube for hydrogen, and pass the side arm of the flask into a condenser. After introducing the hydriodic acid solu- tion, fill the whole apparatus with hydrogen, and keep a slow current of this gas passing during the distillation. On distilling, nearly pure water passes over at first and the thermometer does not register appreciably above 100. When the thermometer rises to 105 change the receiving vessel and collect the distillate until the temperature has risen to 120. Change the receiver again and collect the rest of the distillate. The temperature rises quickly to 126, and remains very close to this point until practically all of the acid has passed over. This last fraction is the desired concentrated acid. Questions i. Compare the stability of the hydrogen halides as judged (a) by the heat produced when they are formed POTASSIUM CHLORATE 149 from the elements; and (b) by the ease with which the compounds are decomposed. 2. Compare the action of the hydrogen halides as reducing agents. Does hydrofluoric acid behave as a re- ducing agent towards any substance ? What substances are reduced by hydrochloric acid? What substances can be reduced by hydriodic acid that are not reduced by hydrochloric acid ? 3. What is the common commercial method for prepar- ing hydrochloric acid from sodium chloride ? Why could not hydrobromic or hydriodic acids be prepared in a similar manner ? 4. Describe and explain the method of preparing hydro- gen iodide from red phosphorus and iodine. 5. As illustrated in the foregoing preparation hydriodic acid is extremely soluble in water. A great deal of heat also is liberated when hydrogen iodide gas dissolves (19,200 calories for each mol of HI). Compare this heat effect with that of some well-known chemical reaction, for example, the neutralization of a strong acid with a strong base. In the absence of water the reaction 2 HI -|- S -* H 2 S -f- ^2 takes place mainly in the direction indicated. What is the connection between the large heat effect just mentioned and the fact that the direction of the reaction is reversed when it takes place in aqueous solution ? 45. POTASSIUM CHLORATE When chlorine dissolves in water it hydrolyzes to some extent : C1 2 + HOH ^ HC1 + HOC1 (i) The presence of a base causes this hydrolysis to run to completion because the two acids produced by the reaction ISO NON-METALLIC ELEMENTS are immediately neutralized; thus by passing chlorine into a solution of sodium hydroxide a mixture of chloride and hypochlorite is obtained : C1 2 + 2 NaOH - NaCl + NaOCl + H 2 O. (2) Sodium hypochlorite is fairly stable in a cold solution, but in a warm solution it is less so ; it gives up its oxygen, and if no more easily oxidizable substance is present it will oxi- dize either chloride or hypochlorite ions to chlorate ions : 3 NaOCl -* 2 NaCl + NaClO 3 (3) or 3 NaOCl + KC1 - KC1O 3 -f 3 NaCl. (4) But compared with a hypochlorite (that is, the OCl~-ion), free hypochlorous acid, HOC1, is a far stronger oxidizing agent, and therefore the formation of chlorate takes place more readily when the solution contains a trace of acid: 3 HC1 + 3 NaOCl - 3 HOC1 + 3 NaCl 1 3 HOC1 + KC1 -> KC10 3 + 3 HC1 j Too much acid, however, causes a reversal of reaction (i), HOC1 -f HC1 - H 2 O + C1 2 , and the action either of chlorine or of hypochlorous acid on a strongly acidified solution cannot produce any chlorate. One method for preparing potassium chlorate is to pass chlorine gas into a hot concentrated solution of potassium hydroxide until the alkali has been entirely neutralized and a small amount of free acid has been formed. This point is recognized by the solution assuming the permanent yellow tint due to free chlorine. This method is the one that will actually be used in the preparation of potassium bromate, No. 46 ; but in the present preparation, to avoid the use of chlorine gas, which is very objectionable in the laboratory, POTASSIUM CHLORATE 151 bleaching powder, or calcium hypochlorite, is employed as the oxidizing agent, and potassium chloride is used in amount only sufficient to supply the necessary potassium ions. Bleaching powder is never obtained as the pure sub- stance represented by the chemical symbol ; it always con- tains a considerable amount of unchanged calcium hydroxide as well as calcium carbonate. The oxidation to chlorate cannot be accomplished rapidly so long as there is any con- siderable amount of calcium hydroxide present, and this is therefore neutralized with hydrochloric acid. After the base is neutralized, any further addition of acid begins to react with the calcium carbonate and the escape of carbon dioxide is observed. The carbonic acid thus formed gives just about the right degree of acidity to the solution to produce the best yield of chlorate. If the hydrochloric acid is used incautiously, and more than enough to react with all the carbonate is added, a corresponding amount of chlorate will be lost, for hydrochloric acid reduces chloric acid approxi- mately according to the reaction HC1 + KC1O 3 -> KC1 + HC1O 3 HC1O 3 + HC1 - H 2 O + Cl + C1O 2 . Procedure. Wet a mixture of 175 grams of bleaching powder and 20 grams of potassium chloride, and rub it with a pestle until a smooth, thin paste is obtained. Transfer this paste to a tall bottle or jar and add hydrochloric acid (at the hood) as follows: Dilute one volume of acid of 1.12 sp. gr. with three volumes of water, and add this diluted acid very cautiously through a thistle tube, the stem of which reaches to the bottom of the liquid. Rotate the bottle or stir the contents vigorously throughout the process until carbon dioxide begins to be evolved freely. This point is recognized when with each fresh addition of \ to i cc. of acid the gas that is produced at the point where the 152 NON-METALLIC ELEMENTS acid enters the liquid rises to the top without being absorbed and causes considerable frothing. Probably about 250 cc. of the 1. 1 2 acid will be required, but this depends very largely on tie quality of the bleaching powder. The reac- tion should now have warmed the solution to about 40, and the formation of chlorate is nearly completed. The solution cannot be filtered at this point, as it still contains so much hypochlorous acid that it would disintegrate the filter paper. Pour it, therefore, into an evaporating dish and boil it (at the hood) until it is concentrated to a volume of about 400 cc. By this time all of the hypochlorous acid has either reacted to form chlorate or has been otherwise decomposed, and the solution may be poured through a large, common filter. Allow the filtrate to cool completely, and collect the crystals of potassium chlorate. Dissolve the crystals in hot water (see solubility table in appendix) and recrystallize twice or three times, as may be necessary to obtain the product entirely free from chloride (test with silver nitrate solution). Questions 1. How is bleaching powder prepared ? What is its formula and its chemical name? 2. Explain why the odor of chlorine becomes very noticeable when, according to the foregoing procedure, the requisite amount of hydrochloric acid has been added. Explain why addition of any further amount of acid 'would cause an appreciable loss of chlorine and a corresponding diminution of product. 3. How many mols of bleaching powder (assuming it to be the pure compound whose formula is given in Ques- tion i) would be necessary to convert i mol of potassium chloride into chlorate? Calculate the weight of this sub- stance that would react with the 20 grams of potassium POTASSIUM BROMATE AND BROMIDE 153 chloride taken, and compare this amount with the amount of bleaching powder actually taken. How much hydrochlo- ric acid would it be necessary to use if the bleaching powder were actually this pure substance ? 4. The modern commercial method of making potas- sium chlorate is by the electrolysis of a potassium chloride solution. What are the primary products formed at the two electrodes ? Explain how the secondary reactions are sim- ilar to those outlined in the introductory discussion. 46. POTASSIUM BROMATE AND POTASSIUM BROMIDE The reaction of bromine on solutions of caustic alkalies is almost identical to that of chlorine, and in this connection the discussion of the preparation Potassium Chlorate, No. 45, should be read. The reaction of bromine on ammonium hydroxide should also be referred to under the discussion of Ammonium Bromide, No. 5. Bromine itself is used in this preparation, as it is not so difficult or disagreeable to handle as chlorine. By its action on concentrated potassium hydroxide solution a great deal of heat is produced, and in this hot solution any hypobromite at first formed is rapidly converted into bromate, so that as a final result i molecule of potassium bromate to 5 molecules of potassium bromide is obtained. By taking advantage of the great difference in the solubility of these salts, the former may be crystallized pure from the solution while the mother liquor contains all of the latter, in addition to the small amount of bromate that is soluble. The potassium bromide could not well be crystallized pure from this solution, but it is possible to reduce the bromate present to bromide by heating with charcoal and then to crystallize pure potassium bromide. 154 NON-METALLIC ELEMENTS Procedure. Dissolve 3 1 grams of potassium hydroxide in 31 grams of water in a 250 cc. Erlenmeyer flask. Place 40 grams of bromine (12^ cc.) in a small separatory funnel, and clamp the latter firmly in a vertical position. Place the flask in a pan of cold water, and lower the stem of the sepa- ratory funnel into the flask until it nearly reaches the surface of the solution. The funnel should now be fastened rigidly, and the flask should be floating on the surface of the bath, so that it may be held by the hand and constantly rotated. Open the stopcock of the funnel cautiously, and allow the bromine to run into the solution at the rate of 2 or 3 small drops per second. The solution should grow hot, but if the reaction becomes violent and red vapors escape from the flask, stop the flow of bromine for a few moments. The reac- tion is complete when the solution has acquired a permanent reddish yellow tint, due to a small excess of bromine. Cool the solution completely, collect the crystals of potassium bromate on a filter, and recrystallize them once or twice from a small amount of hot water until free from bromide (test with silver nitrate). Combine all of the mother liquors, evaporate until a pasty mass is obtained, mix this thoroughly with 5 grams of powdered charcoal, and dry the mass com- pletely. Pulverize the dry mixture in a mortar, and then heat it to redness for an hour in a large porcelain crucible surrounded with a funnel of asbestos. Extract the product with 60 cc. of hot water, filter, wash the residue and the filter with an additional 15 cc. of hot water, and evaporate the solution to obtain crystals of potassium bromide. Questions 1. Write the reaction between bromine and sodium hydroxide in a cold dilute solution; in a hot concentrated solution. 2. Why is it impossible to obtain a mixture of 5 mole- POTASSIUM IODATE 155 cules of ammonium bromide and i molecule of ammonium bromate by the action of bromine on a warm solution of ammonium hydroxide ? 3. Test the Purity of the Potassium Bromide. Dissolve some of the salt in water, and acidify the solution with sul- phuric acid. Appearance of the red color of free bromine indicates that the bromate was not all decomposed by the heating with charcoal. Explain and give reactions. 4. State reasons why it would not be feasible to purify the by-product, potassium bromide, by recrystallization with- out first decomposing the bromate. 47. POTASSIUM IODATE As is well known, the chemical affinity of the halogens for hydrogen or positive elements decreases in passing from fluorine to iodine ; but the affinity for oxygen increases in this order, so that iodates and iodic acid (I 2 O 5 ) are much more stable than chlorates and chloric acid (C1 2 O 5 ). Use is made of this fact in the following preparation, in which the total change is represented fairly closely by the equation, KC1O 3 + I -> KIO 3 + Cl. The actual reaction, however, is not so simple as this. The presence of a small amount of acid is necessary to make it take place. This acid gives rise to a little free chloric acid, which is a far stronger oxidizing agent than potassium chlo- rate, and oxidizes the iodine to iodic acid (I 2 O 5 ). By this reaction more acid (HIO 8 or HC1) is generated, and thus the reaction when once started proceeds to completion. It will be noticed that in carrying out the following directions more iodine is taken than is necessary to react with the potassium chlorate according to the equation given above. This excess of iodine is oxidized to iodic acid by a part of 156 NON -METALLIC ELEMENTS the free chlorine which is represented in the equation as escaping. Procedure. Dissolve 30 grams of potassium chlorate by warming it with 100 cc. of water in an 800 cc. flask. Add 35 grams of powdered iodine and hang a small funnel in the neck of the flask to prevent, to some extent, the escape of iodine vapor. Place a pan of cold water close at hand ; then add i cc. of nitric acid (1.2) to the flask, and warm rather carefully until a brisk reaction commences. Then allow the reaction to proceed so that violet vapors fill the flask, but no appreciable quantity of iodine escapes through the funnel. If the reaction grows more violent than this, check it by dipping the flask for a moment in the cold water. When the reaction is complete, boil the solution until the last trace of iodine has disappeared. Then add i gram more of iodine, and boil, first in the flask and then in a beaker, until the odor of chlorine can no longer be detected. The solution now contains a considerable quan- tity of iodic acid in addition to the potassium iodate. Add a solution of potassium hydroxide until the neutral point is just reached (test by dipping a stirring rod in the solution and touching it to litmus paper). Allow the solutior to cool, collect the crystals of potassium iodate, and evaporate the mother liquor to obtain another crop of crystals. Purify the entire product by recrystallizing once from hot water. Questions i . Experiment. To 3 drops of potassium iodide solu- tion in 10 cc. of water add freshly prepared chlorine water, drop by drop, until the iodine color which at first appears has been bleached. What change takes place in the state of oxidation of the iodine, first when it is liberated from the potassium iodide, and second when it is further oxidized to iodic acid ? Write the equation for the latter action, IODIC ACID 157 2. Dissolve a few small crystals (0.05 gram) of potas- sium iodate in 3 cc. of warm water, and add sulphurous acid, drop by drop, to this solution, noting the successive changes that occur until the solution again becomes clear and colorless. Trace the changes in the state of oxidation of the iodine, giving reactions, and compare with the changes in Experiment i. 48. IODIC ACID; IODINE PENTOXIDE Iodine pentoxide is a white solid substance that at ordinary temperatures is entirely stable. It cannot be pre- pared by direct synthesis from iodine and oxygen, because when cold the elements combine too slowly, and when heated the compound is decomposed into the elements. It may be readily prepared by the direct oxidation of iodine by means of strong oxidizing agents, such as concentrated nitric acid or chlorine. One method for the oxidation of iodine has already been illustrated under the preparation of Potassium Iodate, No. 47, but there the conditions were such that a salt of iodic aoid was obtained rather than the free acid or its anhydride. Starting with this salt, however, the free acid is easily obtained by metathetical reactions which depend on the insolubility of barium iodate and the still greater insolubility of barium sulphate. Procedure. Dissolve 43 grams of potassium iodate and 26 grams of barium nitrate, separately, each in 250 cc. of hot water, and mix the two solutions at the boiling tempera- ture while stirring well. Cool the mixture, let the heavy precipitate settle, decant off the clear liquid, and wash the salt twice by decantation with pure water. Drain the barium iodate on a Witt filter, and wash it on the filter with cold water. Then remove it to a porcelain dish, sus- pend it in 250 cc. of water, heat to boiling, and stir in a 158 NON-METALLIC ELEMENTS solution of 15 grams of concentrated sulphuric acid (8 cc.) in 100 cc. of water. Keep this mixture well stirred at the boiling temperature for at least 10 minutes, since the con- version of solid barium iodate into solid barium sulphate is a reaction that requires some time. Filter the solution and rinse the last of the iodic acid from the solid barium sulphate by washing two or three times on the filter with small portions of water. Evaporate the solution in a casse- role to a small volume, and finally, holding the casserole in the hand, keep the contents rotating, so that the whole inside of the dish is continually wet, and evaporate until solid iodic acid separates in some quantity. Cool com- pletely and rinse the crystals with three successive portions of 10 cc. each of concentrated nitric acid (sp. gr. 1.42), tritu- rating the crystals thoroughly with each portion of the acid. Warm the casserole carefully until the product is perfectly dry and ceases to give off acid vapors. This warming will convert the iodic acid to a large extent into the anhydride I 2 O 5 . Place the iodine pentoxide at once in a sample bottle or tube. . To obtain surely anhydrous iodine pentoxide, the prod- uct could be heated for some time in an oven at about 200. Crystallized iodic acid could be obtained by dissolving the product in a very little water, in which it is extremely soluble, and allowing the solution to evaporate slowly. Questions 1. Experiment. Dissolve a little of the iodine pen- toxide in water. Test the solution in such a manner as to show whether it contains a strong acid. NOTE : A test with litmus is not conclusive, for the preparation may still con- tain a trace of nitric or sulphuric acid which has not been completely removed. 2. Experiment. Heat \ gram of iodine pentoxide in POTASSIUM PERCHLORATE 159 a dry test tube. Insert a glowing splinter in the tube. Note whether the entire substance can be volatilized ; also if any of the original substance deposits in the cooler part of the tube. 49. POTASSIUM PERCHLORATE When potassium chlorate is heated to about 400 it may decompose according to either of the following independent reactions : 4 KC1O 3 = KC1 + 3KC1O 4 , (i) KC1O 3 = KC1 -f 3 O. (2) The second reaction is accelerated by catalyzers, such as manganese dioxide or ferric oxide, or in fact any material with a rough surface. Too high a temperature also causes reaction (2) principally to take place. On the other hand, if the temperature is maintained at the right point, and the salt is free from dirt, and the inside of the crucible is per- fectly clean and free from roughness, the decomposition proceeds mainly according to reaction (i). Potassium per- chlorate, being very sparingly soluble in cold water, may easily be separated from potassium chloride and any unde- composed potassium chlorate by means of crystallization. Procedure. Place 50 grams of potassium chlorate in a dry, clean 100 cc. porcelain crucible, the glaze of which is in perfect condition. Place a small watch glass over the crucible to prevent loss of particles of the salt by decrep- itation, and heat gently until the charge just melts. Then remove the watch glass and keep the melt just hot enough to keep up a gentle evolution of oxygen, but do not increase the temperature when the mass shows a tendency to grow solid. At the end of about 20 minutes the melt should begin to solidify around the edges and should become more l6o NON-METALLIC ELEMENTS or less pasty or semi-solid throughout; when this point is reached, let the contents of the crucible cool completely, then cover it with 50 cc. of water, and let it stand until it is entirely disintegrated. Collect the undissolved potassium perchlorate on a Witt filter and wash it with two successive portions of 15 cc. of cold water (see Note 5 (a) on page 12). Redissolve the salt in hot water (see solubility table) and allow it to recrystallize. About 30 grams of potassium per- chlorate should be obtained. A few crystals of the product should give no yellow color (C1O 2 ) when treated with a few drops of concentrated hydrochloric acid. The product should be entirely free from chloride (test with silver nitrate). Questions 1. Why is manganese dioxide added when oxygen is prepared by heating potassium chlorate ? 2. What is the reaction of hydrochloric acid with hypo- chlorous, chloric, and perchloric acids, respectively ? 3. What are the four oxyacids of chlorine? Compare their stability. 4. To what extent are hydrochloric, hypochlorous, chlo- ric, and perchloric acids dissociated electrolytically in dilute solution ? 5. How could pure perchloric acid be prepared from potassium perchlorate? 6. What is the solubility of silver chlorate and of silver perchlorate ? How may preparations of chlorates and per- chlorates be tested for the presence of chlorides ? 50. SODIUM THIOSULPHATE (Na 2 S 2 O 8 .5H 2 O) Sodium sulphite is a salt of the lower oxide of sulphur, and may thus be regarded as unsaturated with respect to oxygen ; it is, in fact, capable of slowly absorbing oxygen SODIUM THIOSULPHATE l6l from the air and thereby going over into sulphate. If it is allowed to react with sulphur, the latter enters into the compound in much the same way as does oxygen, and //&z'0sulphate instead of sulphate is formed. The sulphur so taken up certainly plays a different function from the sulphur already contained in the compound, although it is perhaps a question whether the thiosulphate is exactly the same compound as sulphate, except that one oxygen atom is replaced by a sulphur. Sodium sulphite is conveniently prepared by allowing sulphur dioxide (sulphurous acid) to react with sodium carbonate. It is practically impossible, however, to dis- tinguish the exact point at which the normal sulphite (Na 2 SO 3 ) is formed; therefore it is more expedient to divide a given amount of sodium carbonate into two equal parts, to fully saturate one part with sulphur dioxide, whereby sodium bisulphite, NaHSO 8 , is formed, and to add the other half of the sodium carbonate, thereby obtaining the normal sulphite, Na 2 SO 3 . Procedure. Dissolve 100 grams of sodium carbonate (anhydrous) in 300 cc. of hot water, and divide the solution into two equal parts. Reserve one part and place about five-sixths of the other half in one flask and the remainder in another flask. Connect these flasks in series so that sulphur dioxide gas may be passed first into the larger volume of solution, and what is there unabsorbed may pass on through the second flask. Draw the gas from a steel cylinder of liquefied sulphur dioxide, if one is available, otherwise generate it by the action of copper with concen- trated sulphuric acid, and pass a vigorous stream of the gas into the solutions. After a short time a marked frothing occurs in the first flask, due to the escape of carbon dioxide, and after this frothing ceases a similar frothing soon com- mences in the second flask. When the latter ceases, pass 1 62 NON-METALLIC ELEMENTS the gas a little while longer until sulphur dioxide escapes freely from the second bottle. Then place the solution of sodium bisulphite in a 600 cc. beaker, and stir in rather slowly the remaining sodium carbonate. Add 45 grams of flowers of sulphur, cover the beaker with a watch glass, and keep the mixture just barely boiling for an hour or longer. Filter the solution, concentrate it to a volume of about 200 cc., and leave it uncovered over night to crystal- lize in a place free from dust. Collect the crystals and obtain further crystals from the mother liquor. Questions 1. Experiment. Dissolve \ gram of the product in 5 cc. of water and add 2 cc. of hydrochloric acid. Observe the odor and the precipitate. What is the free acid corre- sponding to the salt, sodium thiosulphate ? What can be said regarding the stability of this acid? 2. What is the valence of sulphur in each of the salts, sodium sulphide, sodium sulphite, and sodium sulphate? State in each case whether the sulphur plays the part of a positive or negative element. 3. Distinguish between the parts played by the two atoms of sulphur in sodium thiosulphate. 4. Give equations to represent the successive reactions that take place when sulphur dioxide is passed into a sodium carbonate solution. What stage of the process is indicated by each of the succeeding phenomena ? (a) The gas passes into the solution in distinct bubbles and is in large part absorbed, (b) Effervescence takes place with minute bub- bles arising from every part of the solution, (c) Efferves- cence ceases, and the- gas enters the solution again in clear, distinct bubbles, but still it is for the most part absorbed. (d] The gas passes through the solution in distinct bubbles and is entirely unabsorbed. GENERAL QUESTIONS VII 163 GENERAL QUESTIONS. VII NON-METALLIC ELEMENTS OF THE SIXTH AND SEVENTH GROUPS OF THE PERIODIC SYSTEM i. What is the valence of oxygen and sulphur towards hydrogen and metallic elements? What is the valence of the halogens? ?. Give the symbols of the oxides of sulphur and of the halogens which can actually be made. State which of them are salt-forming oxides. Give the symbols of the most important oxyacids of sulphur, chlorine, bromine, and iodine, and state in each case what oxide (actual or hypothetical) is to be regarded as its anhydride. 3. Give data regarding the ease with which the hydro- gen compounds of oxygen, sulphur, and the halogens can be formed ; also state how readily these compounds are decomposed by heat. Draw conclusions as to the relative chemical activity of these elements when they act as negative elements. 4. Compare these hydrogen compounds with regard to their acid strength when they are dissolved in water. 5. Compare the same elements with regard to the ease with which they form compounds with oxygen, and the ease with which these compounds can be decomposed. Draw conclusions as regards the chemical activity of these elements when they act as positive elements. 6. Compare the degree of electrolytic dissociation in aqueous solution of the acids derived from the oxides, SO 2 , SO 3 , C1 2 O, C1 2 O 5 , C1 2 O 7 . When an element forms more than one oxide, which, as a general rule, has the more pronounced acid character? APPENDIX ADDITIONAL GENERAL QUESTIONS I. ALKALI AND ALKALINE EARTH METALS 1. What are the symbols of the oxides of sodium and potassium in which these metals undoubtedly display their ordinary valence of one ? How can these oxides be pre- pared, using the peroxides as a starting point ? How can they be prepared from the hydroxides or from the nitrates ? 2. What oxide is obtained by burning sodium in the air? What relation does this oxide bear to hydrogen perox- ide? What is supposed to be the molecular structure of a peroxide ? 3. What is the difference in meaning between the terms base and alkali? 4. Distinguish between a mild alkali and a caustic alkali. 5. Explain why solutions of the carbonates of sodium and potassium behave as mild alkalies. 6. Judged by its composition, sodium bicarbonate is an acid salt. Explain why in spite of this fact its solution behaves as a mild alkali. 7. What is the difference between ammonia and am- monium ? How can ammonia be set free from ammonium salts ? 8. In its compounds the ammonium radical, NH 4 , behaves similarly to the metals of the alkali group. It is not unreasonable to suppose that if ammonium could be isolated, it would show properties similar to those of sodium 165 1 66 APPENDIX and potassium when in the metallic state. To what extent is this supposition borne out by the facts ? 9. Why is the alkaline strength of an ammonium hydroxide solution greatly reduced by the addition of an ammonium salt ? 10. Give an outline of the modern method for preparing the alkali metals and the alkaline earth metals. What are the most characteristic properties of these metals ? 11. What are the chief differences between the prop- erties of the alkali and alkaline earth metals (compare potassium with calcium) ? 12. Tabulate the solubilities of the hydroxides, car- bonates, and sulphates of the alkaline earth metals, and observe in what sense the solubility changes with increasing atomic weight of the metal in each series of salts. II. ELEMENTS OF THE THIRD GROUP OF THE PERIODIC SYSTEM 1. Give the formulae of normal boric and aluminic acids; of the meta acids; of the anhydrides; and of the sodium salts of tetraboric acid, aluminic acid, and met- aluminic acid. 2. All of the alums are isomorphous. What is iso- morphism ? Give the name and symbol of at least two alums which contain neither aluminum nor potassium. It is not possible to prepare an alum from boron in which this element plays the part of aluminum in common alum. What is the fundamental difference between boron and aluminum which accounts for this difference in behavior ? 3. Before the use of the electrolytic process, what metal was used to reduce the oxide of aluminum in the preparation of metallic aluminum? 4. Describe the principle of the alumino-thermic processes. APPENDIX 167 5. Why would it not be possible to prepare metallic calcium by means of the aluinino-thermic process ? 6. What other metals could be substituted for aluminum in the alumino-thermic process ? 7. Experiment. To a solution of any aluminum salt (use common alum), add a solution of sodium carbonate. What is the gas evolved, and what is the precipitate ? Why is it impossible for aluminum carbonate to form under these conditions ? III. HEAVY METALS OF THE FIRST Two GROUPS OF THE PERIODIC SYSTEM 1. Cite facts which show the difference in chemical activity between the elements of Family A and of Family B in Group I. 2. Compare the chlorides of copper, silver, and gold (copper and gold each possess two chlorides) with regard to their solubility and the valence of the metal. Which of the chlorides are characteristic of the position of these elements in Group I ? 3. Starting with an alloy of copper and silver, devise a method for obtaining the two metals separately in the metallic condition. 4. Experiment. Prepare cuprous oxide by reducing a cupric salt with grape sugar in an alkaline solution. To 20 cc. of a 10 per cent, solution of copper sulphate add 50 cc. of water, and then stir in 20 cc. of a 10 per cent, solution of sodium hydroxide ; add a solution of grape sugar until a deep blue color appears and the precipitate of cupric hydroxide has nearly or quite dissolved; lastly, warm the mixture gently until a yellow and then a clear, bright red precipitate is obtained. Collect this precipitate of cuprous oxide on a filter and wash it with water. Treat the cuprous I 68 APPENDIX oxide on the filter with dilute sulphuric acid, and examine the solution formed and the insoluble residue to determine what products are formed in the reaction. Explain how one-half of the cuprous oxide is oxidized at the expense of the other half which is reduced. 5. Experiment. Add a little cupric salt solution to a potassium iodide solution, and test a drop of the resulting liquid for the presence of free iodine. What is the precipi- tate ? Rinse it free from the dark-colored solution in order to observe its own color. Write the reaction, and explain how the copper salt has acted as an oxidizing agent. 6. Compare the chlorides of zinc, cadmium, and mer- cury (mercury has two chlorides) as regards their solubility and the valence of the metal. 7. In which of their compounds do zinc, cadmium, and mercury most resemble each other, and thus justify their classification in the same family of the periodic system ? 8. In what important respects do these metals differ from the alkaline earth metals which form the other family in the second group of the elements ? 9. Given an alloy of zinc cadmium and mercury, how might one proceed to obtain the three separately in the metallic condition ? i o. Experiment. Prepare a dilute acidified solution of zinc sulphate, using 10 cc. of a 10 per cent, solution and adding 20 cc. of water and 5 cc. of sulphuric acid (i : 4). Divide the solution into two parts, and treat one part with hydrogen sulphide and the other part with ammonium sulphide. Explain why the solubility product of zinc sulphide is exceeded in the one case and not in the other. n. Experiment. Saturate with hydrogen sulphide an unacidified solution of zinc sulphate (made by diluting 5 cc. of the 10 per cent, solution with 10 cc. of water), when APPENDIX 169 saturated the solution will continue to smell of hydrogen sulphide after closing the mouth of the tube with the thumb and shaking. Filter off the precipitate, and test for zinc ions in the filtrate by adding ammonium sulphide. Explain the cause of the incompleteness of this reaction. 12. Experiment. Modify the last experiment by dis- solving 3 grams of solid sodium acetate in the solution before saturating with hydrogen sulphide. Do any zinc ions pass into the filtrate in this case ? How is the differ- ence in the results explained ? 13. Experiment. Dilute 5 cc. of a 10 per cent, solution of cadmium chloride or cadmium sulphate with 10 cc. of water, and saturate with hydrogen sulphide. Filter off the precipitate, and test for cadmium ions in the filtrate by adding ammonium sulphide. Saturate with hydrogen sulphide a similar solution of cadmium salt after first adding 10 cc. of hydrochloric acid (1.12 sp. gr.). How do the solubility products of cadmium and zinc sul- phides compare with each other ? 14. Find out from the text-book, or by experiment, whether the sulphides of mercury, copper, and silver can be precipitated from moderately acidified solutions. Com- pare these sulphides with those of zinc and cadmium with respect to their solubility products. 15. Solutions of mercuric chloride and mercuric cya- nide are poor conductors of electricity. To what general rule do these salts thus form an exception ? IV. ELEMENTS OF THE FOURTH GROUP OF THE PERIODIC SYSTEM i. Why is carbon of importance in the organic world ? With what other elements is carbon found in combination in organic compounds ? I/O APPENDIX 2. In what compounds does carbon most frequently occur in rocks ? 3. What is the relation between an ortho and a meta acid ? Give symbols of ortho and meta silicic, stannic, and plumbic acids. 4. Experiment. Add hydrochloric acid to a solution of sodium silicate (water glass). What is the gelatinous substance obtained, and what change would it undergo if dried and baked ? 5. Name some of the common minerals that consist of some form of silicic acid or its anhydride. 6. With what bases is silicic acid most commonly found in combination in the minerals that constitute our common rocks ? 7. Place tin and lead approximately in their places in the electromotive series of the metals. Explain why either of these metals will react slowly with hydrochloric acid but rapidly with nitric acid. 8. By what oxidizing agents can a stannous compound be converted into a stannic ? By what means can a stannic compound be reduced to a stannous ? 9. Under what conditions can lead in the plumbous state be oxidized to the plumbic state? Is this as easily accomplished as the corresponding oxidation of stannous to stannic tin ? Under what conditions does lead dioxide act as an oxidizing agent? (See Experiment 4 under Lead Dioxide, page 95.) 10. Show how the action of the common lead storage battery depends upon changes in the state of oxidation of lead. 11. Show how the mixed oxides of lead, Pb 2 O 3 and Pb 3 O 4 , may be regarded as salts of ortho and meta plumbic acid. 1 2 . Find out the following facts either from a text-book APPENDIX I/I or by means of experiment : What is the color and what is the chemical formula of the precipitates formed when hydro- gen sulphide is passed into solutions of PbQ 2 , SnCl 2 , and SnCl 4 , respectively? Are all of the lead and tin thrown out in this way from faintly acidified solutions ? Does precipi- tation occur in solutions strongly acidified with hydrochloric acid ? Compare lead sulphide and stannous sulphide with the sulphides of other heavy metals in regard to their solubility product. 13. If a precipitate was obtained containing sulphides of both tin and lead, how might it be treated in order to separate the two metals ? (See Experiment 2 under Stannic Sulphide, page 89.) V. ELEMENTS OF THE FIFTH GROUP OF THE PERIODIC SYSTEM 1. Describe the physical properties of the pure trichlo- rides of phosphorus, arsenic, and antimony. Compare the properties of these compounds with those of the tetrachlo- rides of carbon, silicon, and tin. Are such chlorides of the metalloids to be regarded as salts? 2. What is the behavior of the trichlorides of phos- phorus, arsenic, antimony, and bismuth when treated with water ? Show how their varying behavior corresponds with the increasing metallic properties of the elements. 3. Give reasons why arsenic and antimony should be classed as metals ; as non-metals. Which one of the metals studied under Group IV also falls into the transitional class between metals and non-metals? 4. Distinguish between ortho, meta, and pyro phos- phoric acids. To which of these acids does nitric acid correspond in structure ? I7 2 APPENDIX Which of the antimonic acids yields a sparingly soluble sodium salt ? 5. Distinguish between phosphorous and phosphoric acids. Give equations to show reactions in which nitrous, phosphorous, and arsenious acids act as reducing agents. Give an example of the action of arsenic acid as an oxidizing agent. 6. What is meant by antimonyl and bismuthyl salts ? Give symbols and describe their properties. By what other names are bismuthyl nitrate and antimonyl chloride known ? 7. Arsenic pentasulphide can be slowly precipitated by leading hydrogen sulphide into a solution of arsenic acid containing a large amount of concentrated hydrochloric acid. Does this fact give any indication of the extent to which the ion As +H ++ is capable of existing in solution ? VI. HEAVY METALS OF THE SIXTH, SEVENTH, AND EIGHTH GROUPS OF THE PERIODIC SYSTEM 1. What is the valence of the metal in the oxides Mn 8 O 4 and Fe 3 O 4 ? Explain in what sense these oxides may be regarded as salts. 2 . Experiment. Oxidation of Manganous Salts to Permanganate in Acid Solution : To half a test tube of water add 3 drops of a manganous sulphate solution and 10 cc. of nitric acid. Then add \ gram of lead dioxide and boil the mixture. Let the solid settle, and observe the color of the clear solution. Explain why the presence of hydrochloric acid would prevent the formation of the red color. Recall the action of hydrochloric acid with potassium permanganate. 3. Experiment. Oxidation of a Chromic Salt in Alka- line Solution: To a little of a chromic salt solution add APPENDIX 173 three times its volume of 10 per cent, sodium hydroxide solution. Warm the mixture and pass in chlorine until the color has become a clear yellow. Compare this action with the action of hydrochloric acid with sodium chromate, and explain how the different conditions make these practically opposite reactions possible. 4. Write the reactions for the oxidation of a ferrous salt by (a) nitric acid ; (b) bromine ; (c) potassium perman- ganate. Write the reactions for the reduction of a ferric salt by (d) hydrogen sulphide; (e) sulphur dioxide; (/) stannous chloride. 5. From what oxide of iron are the ferrates derived, and is this oxide acidic or basic in character ? Can the oxide itself be prepared ? How is potassium ferrate pre- pared ? With what compounds of chromium and manganese is it analogous in composition ? 6. Show that chromium exists in the same state of oxidation in both chromates and bichromates. In what respect do these salts differ from each other ? Show that on the other hand manganates and perman- ganates are derived from different oxides of manganese. 7. What salt is formed by treating i mol of potassium sulphate with i mol of sulphuric acid ? How does this salt resemble and how differ from the salt obtained by treating i mol of potassium chromate with i mol of chromic acid ? VII. NON-METALLIC ELEMENTS OF THE SIXTH AND SEVENTH GROUPS OF THE PERIODIC SYSTEM i. The tension of the elements chlorine, bromine, iodine, and sulphur, to pass into the form of simple negative ions in aqueous solution, decreases in the order given. Describe simple experiments which show the correctness of this statement. APPENDIX 2. Experiment. Reduction of any Oxysalt of Sulphur. Mix a little sodium sulphate (or any oxysalt of sulphur) with twice its amount of sodium carbonate ; moisten the mixture j so that some of it may be made to adhere to the charred end of a burnt match. Heat the mixture on the match end in the reducing part of a Bunsen flame until the salt melts. Detach the end of the match with the fused salt, and place it on a bright silver coin, together with i drop of water. After a few moments observe the dark fleck on the coin. Explain the reactions involved in the test. 3. Would it be feasible to reduce sodium sulphate with charcoal as an industrial method of preparing sodium sul- phide ? Discuss the relative advantages of this and other possible methods. 4. If it were desired to reduce potassium chlorate to chloride, why would it be unsafe to mix it with charcoal and heat ? Why does not the same danger exist with sodium sulphate ? 5. Experiment. Agitate a few grains (0.05 gram) of iodine with 2 cc. of water. To what extent does the iodine dissolve ? Add a few crystals of potassium iodide (o.i gram), and observe whether any more of the iodine passes into solu- tion. Account for the increased solubility of the iodine. What are polyiodides? 6. Experiment. Treat gram of powdered sulphur with 5 cc. of a colorless solution of sodium or ammonium sulphide. What is a polysulphide ? To a solution of a colorless sulphide add hydrochloric acid until the solution is acid to litmus. Treat in the same way some of the polysulphide just prepared and compare the results. 7. When a solution of potassium iodide is exposed to the action of the oxygen of the air no change takes place. When, on the other hand, a solution of hydriodic acid is exposed to the same conditions, iodine is slowly displaced APPENDIX 175 and the solution becomes brown, due to the liberated iodine. Show by writing the ionic reactions that when iodine is dis- placed by oxygen, either hydrogen ions must be used up or hydroxyl ions must be produced, and that therefore the reaction is favored by the presence of the former and re- tarded by the presence of the latter. 8. Write reactions in which sulphuric acid acts as an oxidizing agent ; in which sulphurous acid acts as a reducing agent; in which sulphurous acid acts as an oxidizing agent; in which hydrogen sulphide acts as a reducing agent. 9. Write reactions in which hydriodic, hydrobromic, and hydrochloric acids act as reducing agents; in which the oxyacids of the halogens (take, for example, hypochlo- rous, chloric, and iodic acids) act as oxidizing agents. Ex- plain in what way these reactions are of the same character as these written under Question 8. 1 7 6 APPENDIX INTERNATIONAL ATOMIC WEIGHTS, 1915 Symbol. Atomic weight. Symbol. Atomic weight. . . . Al 27 1 Mo 96 . . . Sb 120.2 Nd 1443 A 39 9 Ne 20 ... As 74.96 Nickel . . . Ni 58.68 - Ba 137 37 N 14 01 Be 9 1 Os 190 9 . . . Bi 208 o 16 00 . . . B 11 . . . Pd 106 7 . . . Br 79 92 p 31 . . . Cd 112.40 . Pt 1950 . . . Cs 132.81 . . . K 39 10 . . .Ca 40.09 . . . Pr 1406 c 12 00 Ra 226 4 . . Ce 14025 Rh 102 9 Chi ' Cl 3546 Rb 85 45 mi m Cr 52.0 Ru 101 7 Cobalt Co 58.97 Sa 150 4 Cb 93.5 Sc 44 1 Cu 63.57 . Se 79 2 Dysprosium .... . . .Dy Er 162.5 167.4 Silicon Silver . . .Si . Ag 28.3 10788 Eu 152.0 . . . Na 23 00 Fluorine , . . F 19.0 . . . Sr 87.62 Gd 1573 s 32 07 . . . Ca 69.9 Tantalum . . . Ta 181.0 Ge 725 Tellurium Te 127 5 Gold . . Au 197.2 Terbium ..... . . .Tb 159.2 He 4 Thallium Tl 2040 H 1 008 Th 232.42 In 114.8 . . .Tm 1685 I 126.92 Tin . . . Sn 119.0 Ir 193.1 . . . Ti 48.1 . . . Fe 55.85 . . . W 184.0 Kr 83 u 238 5 L y h m La 1390 v 51 2 Lead . . . Pb 207 10 Xe 1307 Lithium ... Li 7.00 Ytterbium Lu 174.0 . Yb 1720 Me 2432 Yttrium . . . Yt 89.0 Mn 54.93 . . . Zn 6537 He 200.0 . . . Zr 90.6 APPENDIX 177 at d s o ^> St -Oo? 02 d jap 21 . a >-l I- OJ92, 3 It m wl a TABLE OF SOLUBILITIES i In the following tables are given data which should be useful in connection with the preparations and questions in this book. The formulae given are those of the crystallized compounds which most readily separate from aqueous solution at the laboratory temper- ature, but it should be remembered that many salts have several hydrates, and it has often been difficult to decide which one to place in the table. In the second column the behavior of the crystallized salt when it is exposed to the air of the laboratory is indicated : s = unchanged by exposure to atmosphere; e efflorescent; d deliquescent; d, e = deliquescent or efflorescent, according as to whether the humidity is above or below the average ; COo = absorbs carbon dioxide and falls to a white powder ; Ox = compound is oxidized, especially in presence of moisture. In the third column are given the figures for the solubility at o, 25, and 100, except in the cases in which other temperatures are indicated in parenthesis. Fractions have, as a rule, been dropped in giving the solubilities. 1 Much of the data in this table has been obtained from Seidell, Solubilities of Inorganic and Organic Substances. 178 APPENDIX 179 c 4J SOLUBILITY IN Wi LTER. c_u Formula ot crystallized salt. Behavior wh exposed to atmosphere Grams anhydrous salt per 100 grams water in a sat- urated solution at 25 100 Mols per liter of solution at laboratory temperature. Aluminum : chloride A1C1 3 6H 2 O (15) 70 . 4 nitrate, A1(NO 8 ) 3 .9H 2 O sulphate, A1 2 (SO 4 ) 3 .18H 2 O Ammonium : acetate, NH 4 C 2 H 3 O 2 .. d .. . . s . . d very soluble ..31 38 89 .. .... 0.8 bromide, NH 4 Br . . s (10) 66 (30) 81 (100) 128 6 chloride, NH 4 C1 nitrate, NH 4 NO 3 Antimony : chloride, SbCl 3 . . s . . ..d .. d ..29 39 77 .. . 118 214 871.. } hydrolyzes with water to 5 11 sulphate, Sbo(SO 4 ) 3 d (insoluble basic salt ; sulphide, SboS 3 . . s insoluble ; soluble in concen- Arsenic : sulphide, As 2 S 3 trated acids insoluble in water or acids ; Barium : carbonate, BaCO 3 chloride, BaCl 2 .2H 2 O chromate, BaCrO 4 . . . s . . . . s . . s soluble in alkalies 0.0023 ..32 37 59 .. 00004 . . 0.00011 .... 1.7 000015 hydroxide, Ba(OH) 2 .8H 2 O iodate, Ba(IO 3 ) 2 .H 2 O nitrate, Ba(NO 3 ) 2 CO 2 . . s . . . s, d . (0)1.7 (25) 4.7 (80) 101 .0.008.... 0.03 0.2. . 5 10 . . 34 .... 0.2 ....0.001 . . 03 sulphate, BaSO 4 sulohite BaSO 3 . . s . . s 0.00023 (20) 020 (80) 002 . . 0.000010 001 sulphide, BaS.6H 2 O . Ox . very soluble ; hydrolyzes to Bismuth : chloride BiCl 3 d Ba(SH) 2 and Ba(OH) 2 nitrate, Bi(NO 3 ) 3 .5H 2 O sulphate, Bi 2 (SO 4 ) 3 .. d .. d (hydrolyzes with water to insoluble basic salt ; very soluble in acids sulphide, Bi 2 S 3 . . s . . Cadmium : chloride, CdCl 2 .2iH 2 O nitrate Cd(NO 3 ) 2 4H 2 O . . e . . d (0) 90. (18) 110. (ioo)147 (0)110 (18) 127 (60) 326 5 43 sulphate, CdSO 4 2f H 2 O (0) 76 (40) 79 (100) 61 2 i8o APPENDIX c , Formula of crystallized salt. Behavior when exposed to atmosphere. SOLUBILITY IN WATER. Grams anhydrous salt per 100 grams water in a sat- urated solution at 25 100 Mols per liter of solution at laboratory temperature. Cadmium : sulphide, CdS . . S insoluble ; soluble in con- centrated acids 0.0013 ..60 88 159.. (18) 178 ...0.00013 5.2 53 Calcium : carbonate, CaCO 3 chloride, CaCl 2 .6H 2 O chlorate, Ca(ClO 3 ) 2 .2H 2 O .... . . S . . .. d .. .. d chromate, CaCrO 4 .2H 2 O fluoride CaF 2 . . e . . g ..11 12 3... 00016 0.8 0002 hydroxide Ca(OH) 2 C0 2 .. d .. . 0.19.... 0.16 ....0.08 . (18) 122 .... 0.02 52 nitrate Ca(NO 3 ) 2 4H 2 O .... oxalate, CaC 2 O 4 .H 2 O sulphate, CaSO 4 .2H 2 O sulphite CaSO 3 . . s . . . . s . . 0.0007 .0.18 .... 0.21 ... 0.16.. 004 . . 0.00004 . . . 0.015 . . .0.0003 Chromium : chloride CrCH 6H 2 O d 130 8 nitrate Cr(NO 3 )s 9H 2 O very soluble ; melts 36.5 120 insoluble ; soluble in acids ..42 53 104. . 4 3 sulphate, Cr 2 (SO 4 ) 3 .18H 2 O .... Cobalt : carbonate CoCO3 .... . s, e . . . s . . chloride CoCl 2 6H 2 O . . s . . nitrate, Co(NO 3 ) 2 .6H 2 O sulphate, CoSO 4 .7H 2 O sulphide CoS .. d .. . . s . . . . s . . ...(0)84 (91) 340 ... 43 ..26 39 83.. insoluble in water or dilute acids insoluble ; soluble in acids ..71 79 108.. ..82 150 275.. . . 14 23 75 .. insoluble in water or acids very soluble ; melts 35.5 . 2.0 4.7.... 27.5.. 300 (0) 3.5 (25) 11. 4 .. (70) 64 2 .... 5.0 .... 4.8 .... 1.2 .... 0.6 .... 0.9 Copper: carbonate CuCO 3 s . . chloride, CuCl 2 .2H 2 O nitrate, Cu(NO 3 ) 2 .6H 2 O sulphate, CuSO 4 .5H 2 O sulphide CuS .s,d. ] melts ( 38 . s, e . . . s . . Hydrogen : arsenic acid H 3 AsO 4 -iH 2 O .. d .. boric acid, H 3 BO 3 iodic acid, HIO 3 oxalic acid, H 2 C 2 O 4 .2H 2 O . . s . . . d, s. . . s . . APPENDIX 181 ~ , Formula of crystallized salt. Behavior when exposed to atmosphere. SOLUBILITY IN WATER. Grams anhydrous salt per 100 grams water in a sat- urated solution at 25 100 Mols per liter of solution at laboratory temperature. Hydrogen : phosphoric acid, H3PO4 Iron : chloride (ous), FeCl 2 .4H 2 O (ic) FeCl 3 6H 2 O . ... .. d .. ..d .. .. d .. very soluble ; melts 37 ..(15) 67 .... (80) 100.. very soluble ; melts 31 insoluble ; soluble in acids (18) 82 2 carbonate (ous), FeCO 3 .... . . s . . nitrate (ous), Fe(NO 3 ) 2 .6H 2 O . . (ic), Fe(NO 3 ) 3 .9H 2 O .. sulphate (ous), FeSO 4 .7H 2 O . . (ic), Fe 2 (SO 4 ) 3 . Ox . .. d .. e, Ox .. d .. very soluble ; melts 47 (0) 16.. (30) 33. ..(90) 43 very soluble insoluble ; soluble in acids 50 200.. ..0.5 1.0 4.8 .. ..' 0.0001 07 11 33 .... 1.5 ....0.02 ..0.000003 . . 05 . . s . . Lead: acetate, Pb(C 2 H 3 O 2 ) 2 .3H 2 O . . bromide, PbBr 2 carbonate, PbCO 3 chloride, PbCU . . s . . . . s . . . . s . . . . s . . hydroxide, Pb(OH) 2 001 . . .00004 iodide, PbI 2 nitrate, Pb(NO 3 ) 2 sulphate, PbSO 4 . . s . . . . s . . . . s . . . . 0.04 . . . 0.08 0.44 . ...38 59 132.. 0004 . . . 0.002 .... 1.4 . . 0.00013 sulphide, PbS . . . s . . insoluble ; soluble in concen- trated strong acids .. 1.5 1.3 0.7 .. (13) 5 5 ....0.17 08 Lithium : carbonate, Li 2 CO 3 bicarbonate, LiHCO 3 . . s . . . . s . . chloride, LiCl hydroxide, LiOH.H 2 O nitrate, LiNO 3 3H 2 O .. d .. CO 2 d ...67 82 128.. ..12.7 ... 12.9 ....17.5 . (0)54 (30) 138 (70) 176 ...35 34 30 .. 02 .... 13.3 .... 5.0 .... 7.3 .... 2.8 0.01 sulphate, Li 2 SO 4 Magnesium : carbonate, MgCO 3 .3H 2 O . . s . . . . s . . chloride, MgCl 2 .6H 2 O hydroxide Mg(OH) 2 .. d .. CO 2 ..d .. . . e . . . . s . . ...53 57 73 .. 0.001 ..(OP) 67 (40) 85 .. ...27 39 74 .. insoluble ; soluble in acids ...63 77 115.. .... 5.1 ...0.0002 .... 4.0 .... 2.8 5 nitrate, Mg(NO 3 ) 2 .6H 2 O sulphate, MgSO 4 .7H 2 O Manganese : carbonate, MnCO 3 chloride, MnCl 2 .4H 2 O .e, d. 182 APPENDIX c , Formula of crystallized salt. Behavior when exposed to atmosphere. SOLUBILITY IN WATER. Grams anhydrous salt per 100 grams water in a sat- urated solution at 25 100 Mols per liter of solution at laboratory temperature. Manganese : nitrate Mn(NO 3 ) 2 6H 2 O .. d .. (0)102 (25) 166 (35.5)331 . . 53 65 32 .. insoluble ; soluble in dilute acids 0.0002 5 4 . . 0.00001 .... 0.2 . . .00001 sulphate, MnSO 4 .5H 2 O . s, e . . Ox . sulphide, MnS Mercury : chloride (ous) HgCl . (ic) HgCl 2 . . s . . ..3.7 7.4 61 .. 005 nitrate (ous), HgNO 3 .H 2 O (ic),Hg(N0 3 ) 2 4H 2 0... . s, e . .. d .. } very soluble in a little water > much water ppts. basic salt ) very soluble in HNO 3 0.006 ...0.0001 sulphide (ic) HgS s insoluble ; insoluble in con- centrated acids insoluble ; soluble in acids ..54 67 88 .. (0)80 .(20) 96 (95) 233 (0) 27 . . (30) 43 ... (99) 77 insoluble in water or dilute acids (6)188 (14) 230 (G2)492 ..3.1 8.0 50 .. ..54 68 104.. .. 89 .... 113 156.. (0) 22 . . (25) 36 . . .(60) 60 ..3.1 8.2 56 .. ..28 36 57 .. ..59 64 79 .. .. .5 16 89 .. 4 6 2 25 Nickel : carbonate NiCO<> s chloride, NiCl 2 .6H 2 O nitrate NiNO 3 6H 2 O . s, d . .s,d. . e . . sulphate, NiSO4.7H 2 O . sulphide, NiS s Potassium : acetate KC 2 H 3 O 2 d bromate, KBrO 3 bromide, KBr carbonate, K 2 CO 3 .1H 2 O (bi) carbonate, KHCO 3 . . s . . . . s . . ..d .. . . s . . ....0.38 .... 4.6 .... 5.9 .... 2.8 ....0.52 .... 3.9 2.7 .... 0.4 chlorate, KC1O 3 chloride, KC1 chromate, K 2 CrO4 (bi) rhromate K 2 Cr 2 Oy . . s . . . . s . . . . s . . s fluoride KF 2H 2 O d (18) 92 . . . 12.4 hydroxide, KOH.2H 2 O iodate, KIO 3 iodide, KI manganate K 2 MnO4 .. d .. . . s . . . . s . . . d .. .. 97 .... 119 178.. ..4.7 9.9 32 .. ..128.... 148 208.. very soluble ..13 37 246.. . 38 18 ....0.35 .... 6.0 .... 2.6 1.6 nitrate, KNO 3 oxalate K 2 C 2 O4 H 2 O . . s . . g perchlorate KC1O4 1.5 0.11 permanganate, KMnO4 . . s . . (0) 2.8.. (25) 8.0 ..(65) 25 ....0.33 APPENDIX c i. Formula of crystallized salt. Behavior when exposed to atmosphere. SOLUBILITY IN WATER. Grams anhydrous salt per 100 grams water in a sat- urated solution at 25 100 Mols per liter of solution at laboratory temperature. Potassium : sulphate K2SO4 7 12 24 062 (bi) sulphate, KHSO 4 sulphide, K 2 S.5H 2 O sulphite, KoSO 3 .2H9O (o) 36.. (20) 51 .(100) 122 very soluble very soluble (0)0.7. (25) 1.1.. (80) 2.5 00001 .... 3.5 ....0.06 . 0000006 .. d .. .. d .. Silver : acetate, AgC 2 H 3 O 2 . . s . . bromide, AgBr g carbonate, AgoCO^ s 0.003 ...0.0001 .... 0.6 . . 00001 chlorate, AgClOs s 15 chloride, AgCl s 0.0002 chromate, Ag 2 CrC>4 . . s 0.002 . . 000015 fluoride, AgF d (16) 182 .... 13.5 . . 0.00014 iodate, AglOs ... S . 0.005 iodide. Agl 0000003 00000001 nitrate, AgNO 3 oxide, AgoO, dissolves as AgO H perchlorate, AgClO4 . . s . . . . s . . d ..122.... 257 952.. 0.0025 very soluble . . (18) 073 ... (ioo) 15.. .... 8.4 ...0.0002 . . 024 sulphate AffoSOa. s sulphide, AgS s insoluble in water or acids (0) 34.. (25) 53... (40) 65 (5) 1.3 .(30) 3.9.. (10(>) 53 ..73 87 118.. ..70 28 .. . 46 . . 6 Sodium : acetate, NaC 2 H 3 O2.3H 2 O (tetra) borate (borax), Na 2 B 4 O 7 .10H,O bromide, NaBr.2H 2 O carbonate, Na 2 CO 3 .10H 2 O (bi) carbonate, NaHCO 3 chlorate, NaClO 3 chloride, NaCl chromate, Na 2 CrO 4 .10H 2 O (bi) chromate, Na 2 Cr2O7.2H 2 O fluoride NaF . s, e . . . s . . . . s . . Q 0.15 .... 0.9 1 8 . . S . . . . s . . . . s . . . . e . . .. d .. g (0) 6.9.. (25) 10.. .(60) 16 .. 82 105 233.. ..36 36 40 .. (0) 32.. (21) 90 (ioo) 126 ..(0)16<5.... (98)433 .. (2l)4.2 .... 1.1 .... 6.4 .... 5.4 .... 3.3 .... 50 1 l hydroxide NaOH H 2 O d . . 42 114 348 21 iodide, NaI.2H 2 O nitrate, NaNO 3 .e, d. . . s . . g ..159.... 184 302.. ..73 92 178.. . . (15) 3 2 (100) 6 3 .... 8.1 .... 7.4 024 permanganate, NaMnO4-3H 2 O sulphate, Na 2 SO 4 .10H 2 O (bi) sulphate, NaHSO 4 .H 2 O .. .. d .. . . e . . .. d .. very soluble (0) 5.0 (32.75) 50.65 (100) 43 ..(25)29 (ioo) 50 .. .... 1.2 1 84 APPENDIX c |. Formula of crystallized salt. Behavior when exposed to atmosphere. SOLUBILITY IN WATER. Grams anhydrous salt per 100 grams water in a sat- urated solution at 25 100 Mols per liter of solution at laboratory temperature. Sodium : sulphide, Na 2 S.9H 2 O ( Ox )d,e . . e . . s, e . . s . . very soluble (0) 14 . . (20) 27 ... (40) 50 (10) 60.. (25) 76.. (45) 124 001 . . 2 5 sulphite, Na 2 SO 3 .10H 2 O thiosulphate, Na 2 S 2 O3.5H 2 O . . Strontium : carbonate, SrCOs .. 0.00007 chloride, SrCl 2 .6H 2 O hydroxide, Sr(OH) 2 .8H 2 O nitrate, Sr(NO 3 ) 2 .4H 2 O sulphate, SrSO4 'co 2 ' . . e . . . . s . . ..44 56 101.. ..0.4 1.0 32 .. ..40 79 101.. 01 .... 3.0 ....006 .... 2.7 ...0.0006 Tin: chloride (ous), SnCl 2 .2H 2 O Zinc: carbonate, ZnCOs chloride, ZnCl 2 .3H 2 O nitrate, Zn(NO 3 ) 2 .6H 2 O sulphate, ZnSO 4 .7H 2 O sulphide, ZnS . . . f s iOx . . s . . .. d .. ..d .. . . e . . c (0) 84 (15) 270 7 .. 0.0003? k . . . . 9.2 .... 4.7 .... 3.1 0.004 ? ..208.... 432 615.. 95 . ] 27 .. 42 . . .58 81 .. insoluble ; soluble in acids SUPPLEMENT In the preparation of the third edition of Synthetic Inor- ganic Chemistry it has seemed desirable to proceed slowly in order not to be confronted too soon after the completion of the edition with the desirability of making further alter- ations and additions. This supplement to the Second Edition is therefore printed that the Author may have the opportunity to test new ideas in his own classes and that other users of the book may have the benefit of new work which has already seemed to justify itself. 187 DIRECTIONS FOR WORK Preliminary Report. Before beginning work on a prepa- ration the student should have a clear knowledge of the whole procedure and should understand the reactions in- volved as well as the application of chemical principles to these reactions. To that end study carefully the general discussion of the preparation as well as the procedure. Then write in the notebook all reactions, and, starting with the given amount of the principal raw material, calculate what amounts of the other substances are necessary to satisfy the equations. When the amount specified in the directions is different from that calculated, state the reason for the difference. Calculate also on the basis of the equations the amount of the main product as well as of any important intermediate products or by-products. Present this preliminary report to an instructor and obtain his approval before beginning operations. Manipulation. All references from the procedure to the general notes on laboratory manipulation should have been studied before making the preliminary report. Indeed the instructor will probably make sure by a quiz that this has been done before he accepts the preliminary report. Laboratory Record. The working directions, in the sec- tion entitled procedure, are to be kept at hand while carrying out the manipulations. These directions do not need to be copied in the laboratory notebook; but it is essential, nevertheless, to keep a laboratory record in which are entered all important observations and data; such as, for example, 189 I QO SUPPLEMENT appearance of solutions (color, turbidity); appearance of precipitates or crystals (color, size of grains, crystalline form); results of all weighings or measurements; number of recrystallizations; results of tests for purity of materials and products, etc. Questions. The section under this title gives suggestions for study, which involves: (i) laboratory experiments and direct entries in the laboratory notebook; (2) consultation of reference books, of which all that are necessary will be found upon the shelf in the laboratory; (3) original reasoning. The answers to the questions should be written in the laboratory notebook following the entries for the exercise, and this book should be submitted at the same time as the preparation for the approval of an instructor. General Questions. Besides the specific study questions for each preparation there are, accompanying each group of exercises, general questions relating to the whole group; and these are to be worked out by every student. The answers to these questions are to be written on a certain prescribed kind of paper and handed in at the office, neatly bound, within the times which will be posted. Use of Time in Laboratory. In preparation work it is frequently necessary to wait for considerable periods of time for evaporations, crystallizations, etc., to take place. This time may be utilized for work upon the study questions and experiments, but even then it is advisable to have usu- ally more than a single preparation under way. Thus no time need be wasted by the energetic student who plans his work well. Yield of Product. Where possible the methods employed in these preparations resemble those actually used on an industrial scale; where this is, however, impossible on the limited scale of the laboratory, mention is made of the fact, with reasons therefor. On account of the limitations con- SUPPLEMENT IQI nected with work on a laboratory scale, it is of course im- possible to get as high percentage yields as could be obtained on a large commercial scale. The amounts obtained of each preparation are to be weighed and recorded, but the chief stress is to be laid upon the excellence of the product rather than upon its quantity. NOTES ON LABORATORY MANIPU- LATION 12. AUTOMATIC GAS GENERATOR FOR HYDROGEN, HYDROGEN SULPHIDE, OR CARBON DIOXIDE. Supplies: i 300-00. generator bottle (thick walled). i calcium chloride drying tube (10 inches long exclusive of stem; i inch internal diameter), i reservoir of 300 cc. capacity. i 2-hole rubber stopper to fit generator bottle, i i -hole rubber stopper to fit drying tube, i foot rubber tube to connect generator bottle and reservoir, glass wool. This apparatus is based on the principle of the familiar Kipp generator and it is especially suited to cases in which a solution is to be saturated with the gas in question, as, for example, when an ammoniacal solution of common salt is to be saturated with carbon dioxide in the preparation of sodium bicarbonate by the Solvay process. Directions for Setting up and Starting the Generator. As- semble the apparatus as shown in the diagram. The stem of the generator tube E should reach flush with the bottom of the stopper but not below. The delivery tube C should reach nearly to the bottom of the generator bottle D. Place the requisite amount of calcium carbonate (or zinc, or ferrous sulphide) in the generator tube. Then insert a loose plug of glass wool F about i J inches long so that it will stand about 192 SUPPLEMENT 193 midway between the top of the solid material and the stopper in the mouth of the tube, and act as a gas filter (to remove acid spray). Pour the requisite amount of acid into the AUTOMATIC GENERATOR reservoir A ; clamp the reservoir at just the same height as the generator tube and pour in water cautiously until the acid rises and barely touches the solid in the generator tube. IQ4 SUPPLEMENT The generation of gas will now begin and proceed auto- matically. Do not pour any more water into the reservoir until the air is swept from the receiving flask and the mouth of the latter is closed tight (see below). Then add not more than 5 cc. of water. Later after the vigor of the absorption has slackened the reservoir may be raised to a higher level to give a greater pressure and some more water may be added. If the solid charge or the acid becomes exhausted, lower the reservoir to a little below the generator tube; remove the stopper from the generator tube and introduce more of the solid if necessary; if the acid needs renewal, unclamp the reservoir, lower it and invert it to let the spent liquor siphon out of the generator bottle, and refill with acid as in the first charging. Directions for Using the Generator. Place the solution to be saturated in a flask (G in diagram) fitted with a i-hole rubber stopper through which passes a delivery tube reaching to the bottom of the flask. Start the generator in action (see directions above). Loosen the stopper in the flask a little so that gas escapes until all the air originally in the generator and receiving flask is swept out. Then make the stopper tight; the gas will now pass in as rapidly as it can be absorbed by the solution. Shaking the receiving flask will greatly increase the rapidity of absorption, but observe this caution: At the outset if the gas is drawn too rapidly the liquid may rise so far in the generating tube E as to produce too violent an action, which will either blow out the stoppers, or cause froth to pass over into the receiver. Therefore be cautious not to sweep out the air too rapidly, and after the stopper is placed firmly in the receiving flask be cautious during the first 75 minutes not to shake the receiver too strongly. SUPPLEMENT 1 95 I. POTASSIUM NITRATE FROM SODIUM NITRATE AND POTASSIUM CHLORIDE Read the discussion on pages 25-26 of the second edition. In the following procedure equi-molal amounts of sodium nitrate and potassium chloride are taken and enough water added to dissolve at the boiling temperature all of the sodium nitrate taken or all of the potassium nitrate which could result from metathesis, but not enough to dissolve either the potassium chloride taken nor the sodium chloride w r hich could be formed by metathesis. Nevertheless after this mixture is boiled a short time all of the solid potassium chloride disappears and the only solid salt left is sodium chloride. The mechanism of the chemical process may be more easily appreciated if it is represented on paper in the fol- lowing fashion: KC1 -KC1 <=K+ CF NaNO 3 ->NaNO 3 <=NO 3 ~ Na+ IT IT KN0 3 NaCl IT NaCl The formulas printed in bold face type stand for the sub- stances in the solid state, those in common type for sub- stances in the- dissolved state. Single arrows indicate that the reaction runs to completion in that direction under the conditions prevailing (the boiling temperature is supposed to be prevailing in the above representation). The double arrows indicate that an equilibrium is reached and no sub- stance shown on either side of the arrows disappears from the sphere of action. If the conditions were to be shown at o solid potassium nitrate would have to be indicated in equilibrium with the dissolved salt. 196 SUPPLEMENT Materials: crude Chili saltpeter NaNOs, 100 grams. crude potassium chloride KC1, 88 grams. Reagent: i% AgNO 3 solution. Apparatus: 35o-cc. casserole. watch glass. 5-inch funnel. perforated filter plate. 8oo-cc. suction bottle and pump. platinum wire. Procedure. Place 100 grams NaNO 3 and 88 grams KC1 in a 350-cc. casserole. Add 125 cc. of water, cover with a watch glass, and place over a low flame. Keeping an eye on the casserole to see that the contents do not boil, prepare a suction filter according to Note 4 (b), on page 7. Then raise the flame under the casserole and watch it until boiling com- mences. Lower the flame and let the mixture boil gently just one minute, keeping the watch glass over the casserole to prevent too much evaporation of water. While it is at the boiling temperature, pour (see Figure i, page 6) the mixture from the casserole onto the suction filter after first starting a gentle suction. Quickly scrape most of the damp salt onto the filter and suck out as much of the liquid as possible. Then return the solid salt, which is mostly NaCl, to the casserole. Pour the solution into a beaker and cool it to 15 or below by setting it in a pan of cold water or snow r . Separate the crystals of KNOs from the cold liquor by means of the suction filter, observing last sentence of Note 3 on page 7, and pour the liquor into the casserole containing the first crop of NaCl crystals. Bring the solution to boiling point and boil gently three minutes without a watch glass over the casserole, thus allowing some of the water to escape by evaporation. Then filter at the boiling temperature exactly as in the first instance. Cool the filtrate and collect a second crop of KNOa crystals, adding them to the first crop and SUPPLEMENT 197 pouring the liquor into a flask labelled "Mother Liquors." Examine the two kinds of crystals, tasting them and using a microscope. Draw pictures in the notebook of the crys- tals as seen in the microscope. Dissolve about o.i gram of the supposed KN0 3 in 2 cc. of water and test for chloride by adding i drop of AgNOs solution. Considerable chloride will be found and the product must be purified by recrys- tallization. Weigh the crystals roughly while they are still moist, add one half their weight of hot water, and warm until solution is complete. Cool to below 15 and separate the crystals from the mother liquor, adding the latter to the reserve flask. Test as above to see if this crop of crystals is free from chloride. If not repeat the recrystallization as many times as is necessary to get a perfectly pure product. A little of this when dissolved should give no turbidity with silver nitrate solution, and when held in the flame on a plati- num wire should color it the violet color characteristic of potassium with none of the yellow sodium color. Spread the preparation on filter paper or an unglazed porcelain plate and allow it to dry by standing exposed to the air; then put up the salt in a test tube or a small bottle and label it neatly. If the final yield of pure product is not satisfactory in amount the collected mother liquors should be boiled down to about 100 cc., and used as the starting point in a repetition of the above procedure. 30 grams may be regarded as a very satisfactory yield. The sequence of the operations in this preparation can be followed rather more readily in the tabulated procedure shown on the following page. 198 SUPPLEMENT TABULATED PROCEDURE Treat 100 grams NaNO 3 and 88 grams KC1 with 125 cc. water; heat to boiling and boil one minute; filter hot. Do not rinse out dish but keep it for second boiling. On filter: NaCl, dirt, some KNO 3 : transfer back to dish in which first boiling was made (i) Filtrate: cool and filter Crystals: impure KN0 3 (2) Filtrate is saturated with KNO 3 and NaCl. Pour into dish in which origi- nal mixture was boiled and to which impure NaCl (i) was added. Bring to boil, boil 3 minutes, and filter hot. On filter: Filtrate: NaCl cool and filter and dirt Crystals: fairly free i m p ure from KN Q 3 KNO, (4) Filtrate is satu- rated with KNO 3 and NaCl. Save temporarily in flask labelled "Mother Liquors" (5) RECRYSTALLIZATION Unite impure KNO 3 (2) and (4) ; heat with one half their weight of water until dissolved; cool and filter. Crystals: Nearly pure KNO 3 . Recrystallize repeatedly until entirely pure, adding all mother liquors to (5) in the reserve flask. Filtrate contains nearly all of the NaCl from the impure prod- uct, and is saturated with KNO 3 ; add to (5) in the reserve flask. Discard mother liquors (5) if factory. the yield of pure KNO 3 is satis- Questions 1. Define metathesis. 2. When a metathetical reaction is carried out in the wet way, why is the solubility of the substances involved of importance? Explain why, according to this point of view, the reactions AgNO 3 + KC1 = AgCl + KNO 3 and BaCl 2 + Na 2 SO 4 = BaSO 4 + 2NaCl are much more com- plete than the reaction NaNO 3 + KC1 = KN0 3 + NaCl. SUPPLEMENT 199 3. Explain why fewer operations should be required to prepare potassium nitrate from potassium sulphate and barium nitrate than by the foregoing procedure. 4. Explain why all of the solid salt should change to NaCl when the original materials are boiled with insufficient water to dissolve all of either the NaCl or the KC1. 5. In the tabulated procedure what is the advantage of adding the impure NaCl (i) to the second mother liquor instead of discarding it? 3. SODIUM BICARBONATE BY THE AMMONIA (SOLVAY) PROCESS. Read the discussion on pages 31, 32 of the Second Edition. Materials: table salt, 59 grams. i4-normal ammonium hydroxide, 71 cc. cracked marble, 105 grams. i2-normal hydrochloric acid, 175 cc. Apparatus: automatic gas generator (see Note 12 on p. 192 of Supplement). suction filter (see Note 4, a and b on p. 7). 75O-CC. flask equipped with i-hole rubber stop- per, delivery tube reaching to bottom, and 1 8 inches of rubber delivery tube. 3oo-cc. flask with stopper. Procedure. Place the salt, the ammonium hydroxide, and 130 cc. of water in the smaller flask and shake vigor- ously until the salt is dissolved. Pour the solution through a filter into the 750-0:. flask (large plaited filter for speed). Use this flask as the absorption vessel and connect it with the carbon dioxide generator. Charge the generator with the cracked marble and hydrochloric acid and proceed to saturate the solution with carbon dioxide following with care the directions for starting and using the generator (Note 12). After 10 minutes commence shaking the flask 2OO SUPPLEMENT very cautiously and after 5 minutes more, if it is found to be safe, shake vigorously. Let the absorption continue until practically no more gas passes into the absorption flask even with vigorous shaking. If the shaking has been continuous, this point will be reached within one hour. It is equally as well to allow the absorption to proceed by itself over night and to shake next day to complete the saturation with the gas. When the absorption is complete collect the precipi- tated sodium bicarbonate on the suction filter (Notes 3, 4a, and 4b), drain it thoroughly with suction, stop the suction, pour over the surface of the product 15 cc. of cold water, and after this has soaked in apply the suction again. Wash a second time with 15 cc. of cold water exactly as at first. Spread the drained product on a clean unglazed porcelain plate (or on filter paper spread on a folded newspaper) and leave it 24 hours to dry. Test the preparation for chloride by dissolving about o.i gram in a little water, acidulating slightly with nitric acid, and adding a drop of silver nitrate solution. There will be considerable clouding. Questions 1. What is the purpose of washing the product with water? How much sodium bicarbonate is lost in this way (see solubility table)? 2. Why must the solution be acidulated with nitric acid before testing with silver nitrate? 3. Why does shaking greatly increase the rate of absorp- tion? 4. How do you explain the heat produced in the absorption flask? 5. How can you prepare sodium carbonate from sodium bicarbonate? 6. Why cannot potassium bicarbonate be effectively pre- pared from potassium chloride by the ammonia process? SUPPLEMENT 201 (Look up the solubility of potassium bicarbonate.) What process may be used to prepare potassium carbonate from this source? 7. What is an acid salt? How does a solution of an acid salt such as KHSO4 behave toward litmus? Test the be- havior of solutions of NaHCO 3 and Na^COa toward litmus. Explain the cause of this behavior. 8. Would a precipitate of sodium bicarbonate form if carbon dioxide were passed into a solution of sodium chloride alone? Explain the part played by the ammonia in the formation of the product. 9. Explain how a given amount of ammonia may be used over and over again. 3-A. SODIUM CARBONATE FROM SODIUM BICARBONATE Heat the sodium bicarbonate obtained in No. 3 until it is converted into sodium carbonate. Compare the weight obtained with that calculated. 3-B. CAUSTIC SODA FROM SODIUM CARBONATE Apparatus: 8-inch dish, suction filter, burette with normal HC1. i5-cc. pipette. 500-cc. bottle with rubber stopper. Read the discussion and procedure of No. 2, Caustic Potash from Wood Ashes, and adapt that method to the conversion into sodium hydroxide of the sodium carbonate obtained in No. 3~a. Keep the sodium hydroxide solution in the rubber stoppered bottle. Answer Questions i and 2 under No. 2. 202 SUPPLEMENT 5. AMMONIUM BROMIDE The solution contains practically pure ammonium bromide all of which should be recovered. Evaporate to complete dryness and pulverize the caked salt thereby obtained. If preferred the first part of the evaporation may be hastened by heating to gentle boiling with a small flame held constantly in the hand. It is unsafe to put a flame under the dish and leave it, for the solution may " bump " and spatter. Further- more as soon as the salt is dry it will volatilize freely itself. 5-A. MAGNESIUM NITRIDE AND AMMONIUM SALT FROM ATMOSPHERIC NITROGEN Active metals, as those of the alkali and alkaline earth families, when heated, combine readily with nitrogen. In this preparation powdered magnesium in a closely packed mass is allowed to react with air. The oxygen reacts with the upper layers and only nitrogen penetrates to the interior of the mass. Thus a large part of the magnesium should be converted to nitride. On treatment with water the substance hydrolyzes and the ammonia given off can be absorbed in an acid to yield ammonium salt. The mechanism of the hydrolysis of magnesium nitride is probably similar to that of such salts as sodium carbonate and ferric chloride, and it would therefore appear as follows: Mg 3 N 2 ->3M g ++ 6H 2 O ->6OH~ 6H+ I I 3 Mg(OH) 2 2 NH 3 If magnesium nitride can ionize at all the ions which it would yield are obviously Mg ++ and N as shown in the upper horizontal equation. The fact that such an ion as N is entirely unfamiliar does not weaken our belief in SUPPLEMENT 203 the above mechanism, because, as the direction of the arrow shows, the N ion is formed only to be removed according to the right-hand vertical equation. Materials: powdered magnesium, 10 grams. dry sand, 50 grams. Apparatus: iron crucible and cover of 25 cc. capacity. 300-cc. r.b. flask. 5 2 A1N + 3 CO. Now it is well known that aluminum oxide cannot be re- duced to metal by means of carbon but that, at very high temperature the oxygen may be withdrawn and carbon sub- stituted for it, thus yielding the carbide: 2 A1 2 3 + 9 C -> A1 4 C 3 + 6 CO. In an atmosphere of nitrogen, however, the place of that part of the carbon that unites with the aluminum may be taken by the nitrogen, and aluminum nitride is thus obtained. The manipulation of a powerful enough electric furnace on a small scale for a laboratory preparation is scarcely feasi- ble. But if we start with metallic aluminum instead of the oxide we may easily effect its combination with nitrogen. Materials: finely powdered aluminum, 16 grams. The ma- terial sold for use as pigment and often labelled Aluminum Bronze, although it is nearly pure aluminum, is suitable for the purpose. The oil which still adheres from the grinding is of no disadvantage. 208 SUPPLEMENT Materials: lamp black f gram. sodium hydroxide, 30 grams. Use material which comes in sticks or large lumps. It should contain at least 95% of NaOH and not more than a trace of carbonate. Avoid the granulated material for this preparation. Apparatus: iron crucible and cover of 25 cc. capacity. dark-colored goggles (recommended to protect the eyes from the blinding light). looo-cc. round bottom flask. 50-cc. distilling flask. filter bottle (or 500-0:. flask). U-tube J inches wide, yj inches tall. i thistle tube. 1 2-hole rubber stopper. 2 i -hole rubber stoppers. See diagram under No. 5~a for assembling the apparatus. A larger generating flask A is used. The stem of the thistle tube i reaches to bottom and the stop cock is dispensed with. The safety tube h in absorption flask is unnecessary. Procedure. Mix the 16 grams of aluminum and f gram of lamp black thoroughly by grinding together in a mortar. Place the mixture in the iron crucible, packing it down as compactly as possible by tapping the crucible on the desk top. The joint of the cover should be fairly tight to exclude air during the preliminary heating. If a little moist shredded asbestos is packed around under the edge of the cover before the latter is pressed down onto the crucible, such a tight joint is obtained. Heat the crucible with a Tirrell burner rather cautiously at first until the oil upon the aluminum has ceased to give off inflammable gas. Then set the Tirrell burner under the crucible and start heating the top of cover with a blast lamp. When the cover is bright red remove SUPPLEMENT 2OQ the lamp from underneath; keep heating the cover as strongly as possible until the under part of crucible has cooled so that it is no longer visibly red. Remove the cover quickly and play the tip of blast flame upon one part of the surface of the charge until it ignites and becomes blindingly incan- descent. The reaction is then self-sustaining and will spread gradually to every part of the crucible. After examining the cooled product in the crucible pul- verize it and transfer it to the generating flask. Place 75 cc. of water in the absorption flask and 30 cc. of 6-normal H 2 SO 4 in the absorption tube. Add 50 cc. of water to the generating flask. Make sure that this flask is supported in such a way that a pan of cold water can be raised from under- neath to cool the flask whenever the reaction becomes violent and has to be checked. Dissolve 30 g. of NaOH in 50 cc. of water and add 10 cc. of this solution to the generat- ing flask. Warm this flask until a vigorous reaction begins. Remove the flame and add the rest of the NaOH little by little. Thereupon warm the flask during J hour just enough to keep the contents boiling gently. It is now well to let the flask cool a little and to leave it stoppered over night. It may then be reconnected and boiled for 15 minutes. If this is not feasible the boiling should be continued imme- diately for another | hour. Save the solution remaining in the flask as the starting material for the next preparation. Unite the contents of the absorption flask and absorption tube, and add enough more 6-n H 2 SO4 to just neutralize the ammonia. Be careful not to add a drop too much of the H 2 SO 4 because it cannot be removed by evaporation, it is well to hold a part of the solution in reserve in case the neutralization of the main part is slightly overstepped. Evaporate the solution to obtain solid ammonium sulphate, following the direction for Ammonium Bromide on page 202 of this Supplement, 2IO SUPPLEMENT Questions 1. Powdered aluminum, unmixed with carbon, does not enter into rapid self-sustaining action with the elements of the air. Why might one expect that such a reaction would take place? Then how may one explain its not reacting according to expectation? Finally how can you account for the fact that a small admixture of soot makes an energetic reaction possible? 2. What information did you obtain in your tests of the gas passing unabsorbed through the U-tubes? What probably is the gas and how do you account for its forma- tion? (Look up Aluminum Carbide in reference book.) 3. Compare the hydrolysis of aluminum nitride with that of magnesium nitride in No. 5. Can you offer any plausible explanation why the hydrolysis of the latter takes place so much more easily? How does the addition of the sodium hydroxide help with the aluminum nitride? 4. In what important detail does the process followed in this laboratory preparation differ from the commercial process of Serpek, and why would the former not be feasible on a commercial basis? Discuss the possibility of the proc- ess outlined in 5~A ever becoming a commercial process of ''fixing''' nitrogen. 5. What is the object of placing glass beads in the second U-tube? 6. Apply Question 2 under No. 5~A to aluminum nitride and aluminum carbide as well as to magnesium nitride. 7. Make for this preparation a calculation similar to that suggested by Question 5 under No. 5~A. II-B. ALUM FROM SODIUM ALUMINATE BY-PRODUCT, SODA CRYSTALS, Na 2 CO 3 .IO H 2 O The solution left in the generator flask from the preced- ing preparation contained mainly sodium meta-aluminate, SUPPLEMENT 211 NaAlO 2 . Addition of an acid to this salt sets free the weak meta-aluminic acid H 2 CO 3 + 2 NaA10 2 - Na 2 CO 3 + 2 HA1O 2 2 HA1O 2 + 2 H 2 O - 2 A1(OH) 3 The precipitated aluminum hydroxide serves as the starting point in the preparation of alum. Evaporation of the filtrate yields soda crystals as a by-product. Read the discussion and procedure of No. 10, Alum from Cryolite, and note that we are here starting with the sodium aluminate solution which is the first intermediate product in that experiment. Note further that solutions of Na 3 AlO 3 and NaAlO 2 differ only by the amounts of NaOH they con- tain. Apparatus: large flask or bottle, 2 liters. i-hole stopper to fit mouth of same. 2 pieces cloth filter 18 inches square, wooden stand 12 inches square for filter, agate pail. 2 8-inch evaporating dishes. Procedure. To the residue in the generating flask in No. n-A add 100 cc. of water, stir, and filter the solution through paper. Place the clear solution in the large flask or bottle and nearly fill the latter with water. Through the tight-fitting stopper lead a gas delivery tube to the bottom of the liquid and connect with the automatic carbon dioxide generator (Note 12, page 192, of Supplement). Loosen the stopper until all air is expelled, then close the flask and allow carbon dioxide to be absorbed until the solution is saturated. Tack one piece of the cloth filter to the wooden frame, allow- ing the middle to hang 4 inches lower than the edges. Lay the other cloth over the first one. Collect the precipitated aluminum hydroxide on this filter and catch the filtrate in a 212 SUPPLEMENT clean agate pail. Wash the precipitate twice with hot water. Transfer the A1(OH) 3 to a dish and treat it with three times the volume of 6-normal H 2 SO4 which it took to neutralize the ammonia formed in II-A. Warm the mixture for 5 minutes. If much undissolved substance is left, add 5 cc. more of 6-normal H 2 SO 4 , warm one minute, and continue with such additions until the solution is nearly clear. It is important to avoid adding more than a trifling excess of acid. For every 100 cc. of the 6-normal acid used in dissolving the A1(OH) 3 add 17.4 grams of solid K 2 SO 4 . Warm until this is dissolved. Filter the solution and bring it to crystallization as directed in No. 10. Obtain soda crystals from the nitrate in the agate pail. n-c. MODIFICATION OF II-B. BY-PRODUCT, GLAUBER SALT, Na2SO 4 .IO H 2 O Follow the same procedure as in No. II-B except in using H 2 SO 4 instead of carbon dioxide in precipitating aluminic acid. An excess of carbonic acid does no harm, but an excess of H 2 SO 4 dissolves an equivalent amount of A1(OH) 3 . Before beginning operations, submit to an instructor your plan for telling just how much H 2 SO 4 to add to completely precipitate the A1(OH) 3 without redissolving any. Obtain crystals of alum and of Glauber salt as the by-product. 23. LEAD NITRATE Note to Procedure. The solution which is set to crystallize should be slightly acid, enough to redden litmus. If in- sufficient nitric acid is used, the excess of PbO dissolves in hot concentrated Pb(NO 3 ) 2 solution forming the basic salt PbOH.NO 3 which separates as a fine granular or flaky pre- cipitate when the solution cools. SUPPLEMENT 213 35. CHROMIC ANHYDRIDE, CrO 3 Materials: sodium dichromate, Na 2 Cr 2 O7.2H 2 O, 100 grams. concentrated sulphuric acid, 400 cc. Apparatus: 8-inch evaporating dish. glass plate to cover the 8-inch dish. glass marble. unglazed porcelain plate. glass-stoppered sample bottle. Procedure. Dissolve the 100 grams of sodium dichromate in 250 cc. of water and filter from any sediment. Add rather slowly with constant stirring about half of the con- centrated sulphuric acid until a slight permanent precipitate of CrO 3 is formed. Let the mixture cool for half an hour or longer, then add slowly, while stirring, the rest of the sul- phuric acid. Let the mixture stand over night covered with a glass plate in order that the crystal meal may become somewhat coarser. In such a crystal meal standing in its saturated solution, the smaller grains dissolve and their material deposits out on the larger crystals. But even now the crystal meal will be rather fine and it will at first run through the filter; if, however, while waiting, the mixture is heated with stirring to 100 and allowed to cool slowly, and this process is repeated once or twice, a more satisfactory product will be obtained. To collect the crys- tals, use a suction filter, but place a small glass marble in the funnel instead of the usual plate and paper. If the red crystals at first run past the sides of the marble, pour the liquid in the bottle repeatedly back onto the filter until finally the filtrate runs clear (see last sentence of Note 3 on page 7). After draining the crystals completely and press- ing the surface with a glass spatula, stop the suction and pour 15 cc. of the concentrated nitric acid so as to wash down the sides of the funnel and cover the surface of the product. 2 14 SUPPLEMENT Stir up the product with this washing fluid for a depth of about \ inch. Suck dry and repeat the operation twice with 10 cc. of nitric acid each time. Finally transfer the product to an unglazed porcelain plate, place the latter on an iron ring and heat it by playing under it the burner held in one hand while with the other hand the crystals are con- tinually stirred. Continue this operation, being very careful not to overheat, until nitric acid vapors cease to be given off. Transfer the product at once to a dry previously weighed glass-stoppered bottle. 364847 UNIVERSITY OF CALIFORNIA UBRARY UNIVERSITY OF CALIFORNIA LIBRARY, BERKELEY THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW expiration of loan period. THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW AN INITIAL FINE OF 25 CENTS WILL BE ASSESSED FOR FAILURE TO RETURN THIS BOOK ON THE DATE DUE. THE PENALTY WILL INCREASE TO SO CENTS ON THE FOUR! DAY AND TO $1.OO ON THE SEVENTF OVERDUE.