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A study of the double sulphates of some rare-earth elements with sodium, potassium, ammonium and thallium

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A study of the double sulphates of some rare-earth elements with sodium, potassium, ammonium and thallium

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Content A STUDY OF THE DOUBLE SULPHATES OF SOME RARE-EARTH ELEMENTS WITH SODIUM, POTASSIUM, AMMONIUM AND THALLIUM & n /3 s A Thesis Presented to the Faculty of the Department of Chemistry University of Southern California In Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy by Ernest Lawrence Bickerdike June 1937 UMI Number: DP21722 All rights reserved INFORMATION TO ALL USERS The quality of this reproduction is dependent upon the quality of the copy submitted. In the unlikely event that the author did not send a complete manuscript and there are missing pages, these will be noted. Also, if material had to be removed, a note will indicate the deletion. Dissertation Publishing UMI DP21722 Published by ProQuest LLC (2014). Copyright in the Dissertation held by the Author. Microform Edition © ProQuest LLC. All rights reserved. This work is protected against unauthorized copying under Title 17, United States Code uest ProQuest LLC. 789 East Eisenhower Parkway P.O. Box 1346 Ann Arbor, Ml 48106- 1346 This dissertation, written by ... under the guidance of A_is_ Faculty Committee on Studies, and approved by all its members, has been presented to and accepted by the Council on Graduate Study and Research, in partial fu l fillm ent of requirements for the degree of DOCTOR OF PHILO SO PHY Dean . . S . Secretary D ate MAT, ■ .1957........ I f Committee on Studies TABLE OF CONTENTS PAGE INTRODUCTION .................................... 1 GENERAL METHODS FOR THE PREPARATION OF DOUBLE SULFATES OF THE RARE-EARTH ELEMENTS........... 6 Desiccation ................................ 8 Electrolysis , ............................ 8 Equilibrium and Solubility Curves ........... 8 SOURCE OF MATERIALS............. . ............. 11 PART I - PRELIMINARY EXPERIMENTAL WORK........... 14 PART II - PREPARATION AND ANALYSIS OF THE DOUBLE SULFATES...................................... 18 Determination of Water of Hydration.... 20 Determination of the Rare-Earth ............. 23 Determination of the Sulfate Ion ....... 25 Determination of Thallium as Thallous Iodide • 26 Double Sulfates of Samarium and Potassium . . . 35 Double Sulfates of Samarium and Sodium . . . . 38 Double Sulfates of Samarium and Thallium ... 42 Double Sulfates of Praseodymium and Potassium . 47 Double Sulfates of Praseodymium and Sodium . . 50 Double Sulfates of Praseodymium and Thallium . 51 Double Sulfates of Praseodymium and Ammonium . 55 Double Sulfates of Ytterbium and Thallium ... 56 PAGE Double Salts of Europium and Potassium . . . . 58 Double Sulfates of Europium and Sodium .... 59 Double Sulfates of Europium and Thallium ... 60 Double Sulfates of Gadolinium and Potassium . 63 Double Sulfates of Gadolinium and Sodium ... 66 Double Sulfates of Gadolinium and Thallium . . 67 PART III - SOLUBILITY DETERMINATION............. 69 Apparatus ................. 71 Experimental ..... ........ ....... 73 Saturation of solution............. 73 Samples.............................. 74 Analysis of samples................... 74 PART IV - SUMMARY AND CONCLUSIONS............... 80 BIBLIOGRAPHY.................................. 82 INTRODUCTION The rare-earths present themselves as a group of elements whose properties are so nearly alike that their separation into smaller groups and then into compounds of the elements themselves is a long and difficult task. Of the many methods for this preliminary separation, the use of sodium sulfate or potassium sulfate is usually the first used in order to separate these rare-earths into small groups from which separation into still smaller groups can be made. It is generally known that there are two main groups of these earths based upon the differences in solubility of the rare- earth sulfates in solutions of sodium or potassium sulfates. Those which are relatively insoluble in the alkali sulfate solution are known as the cerium group, and those relatively soluble in the alkali sulfate solutions are known as the yttrium group. Other than this, little is known in this sul fate separation either of the actual salts formed or of the solubility of these salts in water. From purely empirical investigation and practical experience it appears that potassium sulfate is superior to sodium sulfate as a reagent for the separation of the cerium group members from those of the yttrium group. Certainly in the case of yttrium, - and probably the same holds for other members of the yttrium group, - the double sulfates 2 increase in solubility with the addition of increased amounts of sodium sulfate, but a point is reached when the addition of further amounts of sodium sulfate causes the precipitate of the double sulfate to become less soluble again. It has been stated"** that during this change of solubility a change in the solid phase occurs, simple yttrium sulfate (as an example) existing as the solid phase up to a certain concentration of sodium sulfate, after which a double sulfate of yttrium and sodium is formed, and this becomes progressively more insol uble as the concentration of sodium sulfate is increased. With potassium sulfate these conditions do not seem to hold, and, therefore, this reagent can be added with less regard to the danger of increasing too greatly the concentration of the alkali sulfate. Even under the best conditions the separation into groups on the basis of different behaviour toward alkali sul fate is by no means sharp and quantitative, as the precipi tate of the cerium group will be contaminated by some yttrium group compounds, and the filtrate will be found to contain a certain amount of the cerium group compounds. A preliminary investigation has been carried out to determine what degree of separation might be reasonably ex pected when the precipitation of the cerium group was forced ^James and Holden, J. Amer. Chem. Soc., 35:559, 1913. 3 to certain degrees of completeness. Details of this inves tigation will be found in the later pages. Many references can be found in the literature to a fairly large number of the double sulfates of rare-earths with alkali metals, and yet many of these salts referred to were not prepared as such, but were indicated from a study of the equilibrium curves showing the equilibria existing in solutions containing varying amounts of the rare-earth and alkali sulfates. In those cases where the salts have been prepared in the pure form, it was seldom that more than their general appearance has been recorded. Therefore, the present work had as one of its objectives, the preparation of pure double sulfates of several of the rare-earths with alkali sulfates, and the determination of some of their char acteristics as well as their solubility in water. In the last few years, among the several methods that have appeared for the cleaner and more rapid separation of the rare-earth elements is one described by L. Fernandes. Instead of using sodium or potassium sulfate, he made use of thallous sulfate. He found that there was a rapid and clean 2a. L. Fernandes, Gazz. chim. ital., 55:3-6, 1925; Chem. Abstr.. 19:2175, 1925. b. Ibid.. 54:623-28, 1924; Chem. Abstr.. 19:221, 1925. c. L. Fernandes, L. Holla and V. Cuttica, Ibid.. 54:617-22, 1924; Chem. Abstr.. 19:220, 1925. 4 separation of the earths into groups and even into compounds of the pure elements. This reagent is becoming more avail able as other uses are found for it, particularly in agri culture. A disadvantage in its use, and an important one, is that it is extremely poisonous. With this new reagent available for the separation of the rare-earths, actual knowledge of the compounds it forms, and the comparison of their properties with those of the salts formed from the alkali sulfates and the correspond ing rare-earths, should be of great value in aiding develop ment and application of it. It was with this view in mind, that another of the objectives of this present work, that of studying the double sulfates of some of the rare-earths with thallous sulfate, was set. Several of the rare-earths were not available, but it was thought that by studying certain representative mem bers of the two main groups of the rare-earths, the general characteristics of the other members of those groups would be indicated. For this reason, the present work includes the rare-earths: praseodymium, a representative member of the cerium group; ytterbium, a representative member of the yttrium group; and europium, samarium, and gadolinium, ele ments whose characteristics place them part way between the two extremes as represented by praseodymium on the one hand and ytterbium on the other. For the alkali sulfates, sodium 5 and potassium were used. Included with these were ammonium sulfate and thallous sulfate. Very few of the double salts of samarium or gadolinium have been prepared, and none has been prepared using thallous sulfate. The present paper pre sents the first report of data concerning the preparation and properties of double sulfates of europium with the alkali metals and with thallium. GENERAL METHODS FOR THE PREPARATION OF DOUBLE SULFATES OF THE RARE-EARTH ELEMENTS The main difficulty in the preparation of the double sulfates of the rare-earths and alkalis lies in ob taining a pure salt rather than a mixture of several double sulfates, or of a double sulfate and either or both of the original salts with which one starts. The several methods that have been used by other workers in this field may be classified into four groups: (l) Those making use of temperature, either to concentrate the solutions, or to decrease the solubility of the rare- earth salts; (2) Desiccation, using special drying agents or reduced pressure; (3) Electrolytic; and (4) The use of equilibrium or solubility curves for mixtures of the rare- earth salts and the alkali salts. This last method does not always result in the actual separation of the pure double salt. Methods Reiving upon Effect of Temperature. The simplest way to prepare these salts is to add excess of the alkali sulfate, either as a solid or a solu tion, to a solution of the rare-earth sulfate. This results in the almost immediate separation of the double sulfates of many of the rare-earths, especially those of the cerium group. With this method there is always present the excess 7 alkali salt that must be washed from the double sulfate salts. Particularly is this true when the solid alkali salts are used, and the more washing required, the less dou ble sulfate is recovered. If, instead of adding an excess of the alkali salt, only enough is added to almost saturate the solution, then, on concentrating slightly, the double salt should form. The mixture of the rare-earth solution and the alkali sulfate solution may be allowed to evaporate slowly at room tempera ture until saturation is reached and precipitation takes place, or it may be held at any temperature so that evapora tion proceeds slowly or rapidly by properly controlling the loss of water. Obviously, since increased temperature has the effect of decreasing the solubilities of the sulfates of the rare-earths, care must be taken to prevent their crystal lization before any possible double sulfate might form. The rapidity with which evaporation takes place may affect also the salt formation either by changing the crystal size or by changing the actual composition of the salts obtained. Another variation in the use of temperature control on the salt formation is obtained by holding the solution at, or a little below, the freezing point of water and at the same time passing cold air across the solution. ^a. C. Baskerville and Hazel Holland, J. Amer. Chem. Soc., 26:71, 1904. b. C. Baskerville and E.G. Moss, Ibid., 26:67-71, 1904. 8 Desiccation. The process consists simply in holding the solu tions under diminished pressure in a suitable container, usually a disiccator. A drying agent may or may not be added, but its use allows a more rapid evaporation at ordi nary or diminished temperatures. Electrolysis. 4 This method, known also as Puccini’s method, is widely applied to the formation of salts where oxidation and reduction must take place in order that the ions of the salts may have the proper valence. In the application of this method to the preparation of the double sulfates of the rare- earths and the alkali metals, it appears that the electric current is used to direct the movement of the ions rather 5 than employed for their oxidation or reduction. Equilibrium and Solubility Curves. This method has been used extensively by Zambonini and his co-workers^, as well as by Keyes and James7, and 4J.L. Howe and E.A. O’Neal, Ibid.. 20:759-65, 1898. ^See References 3a and 3b. 6a. F. Zambonini and G. Carobbi, Atti. accad. Lin- cei. (v),II, 33, 301-8, 1924; Chem. Abstr.. 19:2309, 1925. b. F. Zambonini and V. Caglioti, Ibid.. 308-13, 1924; Chem. Abstr.. 19:2309, 1925. 7D .B. Keyes and C. James, J. Amer. Chem. Soc.. 36: 634-38, 1914. 9 0 Bissell and James, Keyes, Bissell, and James rotated bot tles containing various mixtures of the rare-earth sulfates and the alkali sulfates for several months in a constant temperature bath in order to establish equilibrium in the mixtures. Samples of both the liquid and the solid phases present in each bottle were analysed, and the results plotted as solubilities in water of the salts used. The curves thus obtained indicated the presence of the various salts in the solutions and the presence of double salts as well. By mak- 9 ing use of F.A.H. Schreinemaker*s method for analysis of residues, the actual composition of the solid phases in the bottles in which double sulfates were indicated could be determined. Zambonini and his co-workers used essentially the same method, but evidently obtained equilibrium after only fifteen days1 rotation in a bath at 25° C. This time pro bably could be shortened considerably. Workers making use of this method did not always separate the double sulfates that were indicated by the equilibrium curves they obtained. Another disadvantage, 8D.W. Bissell and C. James, Ibid., 58:875-75, 1916. 9a. F.A.H. Schreinemaker, Z. nhysik. Chem.. 9:65-71, 1892. b. Cameron and Bell, J. Amer. Chem. Soc.. 32:869, 1923. 10 that is really serious when ordinarily only very small amounts of the rare-earth salts are available, is that relatively large amounts of salts are needed in order to cover the range of con centration that is desired. After a consideration of these methods as used by other workers, it was decided that those methods making use of temperature effects were the most desirable to use for the preparation of pure double sulfate salts, particularly where the amount of rare-earth material available was very small. SOURCE OF MATERIALS The group of cerium earths used in the sulfate separation study had been obtained by fractionation of the earths from monazite sand as double magnesium nitrates until no absorption spectrum lines of the yttrium group were visi ble • The yttrium group earths used in the separation tests were originally obtained from gadolinite by fractiona tion of the bromides until lines of the cerium group elements could no longer be seen. The praseodymium salt was that used by Sarver and Brinton in their study of the solubilities of some rare- earth oxalates. The ytterbium compounds used were those prepared and used by Brinton and James^ during their work on methods for concentrating the erbium earths. The gadolinium and samarium compounds were from 1P those used by Pagel and Brinton*^ in their study of the high- ^Sarver and Brinton, J. Amer. Chem. Soc.. 49:945, 1927. •^Brinton and James, Ibid., 43:1397-1401, 1921. 12Pagel and Brinton, Ibid.. 51:43, 1929. 12 er oxides of some of the rare-earth elements. The europium sulfate was kindly furnished by Dr. Herbert N. McCoy.^ As all of the rare-earth compounds, with the ex ception of europium, were in the form of oxides or oxalates, a common method was used for the preparation of the sulfate solutions to be used in this work. The oxalate was first ignited to the oxide. This was then dissolved in the least possible dilute sulfuric acid. Evaporation of this solu tion to crystallization, followed by recrystallization of the material formed from water, gave the desired rare-earth sulfate salt. A small part of the thallous sulfate used was ob tained in a high state of purity. The major portion of the material used, however, had to be purified. The thallous sulfate that was available was the grade that is used as an agricultural poison. A solution was made of a small amount of this and tested with hydrogen sulfide to determine the presence of any of the metals which precipitate as sulfides from an acid solution. There was none present. The entire amount of the thallous sulfate was then made into a solution and treated with excess of hydrochloric acid to precipitate white insoluble thallous chloride. This was washed many 13H.N. McCoy, Ibid.. 58:1577-80, 1956 13 times by decantation, and then transferred to a large porce lain casserole. As much water as possible was drained from it, and then an excess of concentrated sulfuric acid was add ed. This mixture was heated, first slowly to drive out the water, and then more strongly to decompose the thallous chloride, drive out the hydrogen chloride formed, and to change the thallous chloride to thallous acid sulfate. Strong ignition until dense fumes of sulfur trioxide were formed, in sured complete removal of the hydrogen chloride. The thal lous salt was allowed to cool and solidify, and then ground up and mixed with water. This treatment changes the acid sulfate to the normal sulfate as the thallous sulfate is sol uble in water only to the extent of 4.859 grams per 100 grams 14 of water at 20° C. Filtering and washing gave thallous sulfate. This was further treated by dissolving in hot water and crystallizing the thallous sulfate from this solu tion. The resulting salt was then ready for use. 14Noyes, Ibid.. 33:1657, 1911. PART I PRELIMINARY EXPERIMENTAL WORK It was desired to find the relative amounts of the rare-earths precipitated as the double sulfates from mixtures of these earths when varying amounts of the cerium group were precipitated. This was to be accomplished by precipitating them from their solution with solid potassium sulfate until approximately one-fourth, one-half, and then all of the neo dymium was removed from the solution. The relative amounts of the neodymium removed were estimated by comparison of the absorption spectra of these solutions with those of standard solutions prepared for this purpose. The rare-earth sul fates precipitated as the double sulfates with potassium sul fate, and those remaining in solution, were finally obtained as oxalates, ignited and weighed as mixtures of their oxides. Five-gram samples of the rare-earth mixture were used and the following results were obtained. The sample used was a monazite residue which had been partially enriched with respect to the less basic elements and contained 28.65 per cent of cerium dioxide. (The ordinary monazite group of rare-earths would carry about 50 per cent of cerium dioxide.) The yttrium group represented probably not over 4 to 6 per cent of the mixture, but this together with samarium, europium and gadolinium might account for 15 per cent of the whole. 15 When the precipitation was carried to the point at which 25 per cent of the neodymium had been removed from the solution, the percentage of the total earths that was preci pitated was between 25 and 30 per cent. When precipitation was carried to 50 per cent com pletion of the precipitation of neodymium from the solution, 60 to 67 per cent of the total was precipitated. When all of the neodymium was precipitated, the total precipitation of the earths ran from 94 to 100 per cent. From these very rough preliminary trials on a mix ture of indefinitely known composition, it seems that probably 75 per cent precipitation of the neodymium would be about the least that could be considered in practical work. After determining the relative amounts of the cerium and yttrium groups precipitated by potassium sulfate, the question arose as to whether potassium sulfate or sodium sul fate gave the better separation of these rare-earth groups. In order to determine this, a mixture of rare-earths was pre pared containing 62 per cent cerium earths and 38 per cent yttrium earths. The same procedure was used on these samples for determining the relative amount of cerium earths precipitated as in the previous work. For this work, three-fourths and all of the cerium earths were precipitated, the amount being determined by spectroscopic examination of the solution. 16 Samples weighing near 2 grams were used* The results of this part of the work are found in Table I. From these results it seems immaterial whether sodium or potassium sulfate is used, as long as one is care ful not to add an undue excess of the reagent. When all of the neodymium is precipitated it is quite evident that a very considerable amount of the yttrium group is co-precipitated. When 75 per cent of the neodymium is thrown out, the total of apparent cerium earths precipitated is roughly from 65 to 70 per cent (instead of the theoretical 60 per cent for this particular sample) but in this approximation, even, there is much compensation, as shown by the distinct absorption spec trum of neodymium in the filtrate. Far from being a quantitative method of separation, this double sulfate method appears to be little more than a rough fractionation method, - much more rapid, to be sure, than fractionation methods for the separation of the compounds of individual elements; but it is evident that before any great degree of reliability can be developed in this separa tion a great deal of fundamental study of the individual dou ble sulfates must be carried out. HWtO^lOCOC^CO TABLE I Sample pptg. agent ^ 4 n n Na^S04 n n Pptd. # Nd Separation of sample % in ppt. % in filtrate total all 82.70 17.46 100.16 75# 68.51 31.35 99.66 all 79.16 18.59 97.75 75# 68.05 27.95 96.00 all 82.44 15.20 97.66 all 94.99 4.57 99.56 75# 69.99 32.13 102.12 75# 65.23 33.87 99.10 PART II PREPARATION AND ANALYSIS OF THE DOUBLE SULFATES The method of preparation was similar with all salts. Mixtures of the solutions of the rare-earth sulfate and the alkali sulfate were prepared, and either allowed to evaporate slowly at room temperature, or held at some ad vanced temperature, usually 45° C., for varying periods of time and under different rates of evaporations, A general procedure of drying and analysis was used with every salt thus obtained. Every salt was washed at least five times with small amounts of distilled water, either by decantation or on a filter, depending upon the size of the crystal particle. The excess moisture was re moved by draining on filter paper. The remaining water absorbed by the salt was removed by drying in a desiccator for a period of five days. The desiccant used was anhydrous calcium chloride. With every salt prepared, the purity will depend to a large extent upon the thoroughness with which the ad hering mother liquor is removed. Theoretically a large number of washings would insure a purer product than would only a few washings. Washing, ho?trever, also dissolves part of the salt. In practice, then, a choice must be made so 19 that the greatest purity is obtained with least loss of salt by solution in the wash water. In the present work, five washings were chosen as the number that would remove at least a major portion of any adhering mother liquor and still would not dissolve too much of the prepared salt. That this choice is satisfactory is evidenced by the number of analyses of the salts in which there is very close agreement between calculat ed and theoretical percentage composition. The fact that there was appreciable loss of salt even with this small num ber of washings is shown clearly by the lack of sharp crystal outline in many of the illustrations of the salts prepared. The use of anhydrous calcium chloride as the drying agent for the removal of moisture other than water of hydra tion from the salts raises the question of its possible effect upon the composition of the salt during drying. With some common salts, such as our well known efflorescent hydrates of sodium sulfate and sodium thiosulfate, there would be actual loss of water of hydration if kept over calcium chloride. Other workers,^ however, have used this material as a desic- cant in this type of work without having noted any effect upon the salt other than that of drying it. In addition, there are observations that may be made directly upon the l^See Reference 6a. 20 salts themselves before and after being dried over anhydrous calcium chloride, which can be used to tell whether or not some change in composition has taken place. Salts in losing water of hydration tend to change color and to change or lose their crystal form. Frequently a crystalline material be comes powdery or coated with a fine powder that makes the salt opaque where before it was perfectly clear. Where changes such as this have been observed they have been referred to under the discussion of the salt and its prepara tion. There were, however, few such changes observed. The amount of salt obtained was usually quite small, so that it was necessary to make a complete analysis upon a single sample rather than use separate samples for the determination of each of the substances whose amounts had to be known in order to calculate the composition of the salt. Each salt contains a sulfate of a rare-earth, a sul fate of an alkali metal or of thallium, and possibly water of hydration. These substances can be calculated by deter mining the amount of sulfur trioxide, oxide of the rare- earth, thallium, and water of hydration. After the salt had been dried in a desiccator for at least five days, a sample was weighed into a porcelain crucible. The size of the samples varied from 0.05 gram to 0.2 gram, depending upon the amount of the salt available. Determination of Water of Hydration. The weighed sample in the porcelain crucible was 21 then gradually heated to a maximum temperature of 450 to 500° C., and re-heated to constant weight. The loss in weight is water of hydration, as none of the sulfates that might be present in the salts begin to decompose at that temperature. The dehydration temperature of 450 to 500° C., used in this work was chosen after a consideration of the temperatures found necessary by other workers, and of the effect of this temperature upon the dissociation of the sul fates of the rare-earth and the alkali elements of which the double salts are composed. An advanced temperature of this order is necessary to dehydrate completely many of these rare-earth double sul- fates according to Zambonini and Carobbi. And in order to dehydrate certain double sulfates of cerous cerium and sodium, other workers^ found that a temperature of 400° C., was necessary. As for the individual sulfates of the rare- earth elements, at least two require a high temperature for 18 their dehydration: europium sulfate requires 375° C., and l^See Reference 6a. ^Zambonini and Restaino, Atti. accad. Lincei. 14: 69-71, 1931; Chem. Abstr.. 26:1870, 1932. -^Katz and James, J. Amer. Chem. Soc.. 36:779, 1914. 22 19 erbium sulfate requires a temperature of 400° C. pn Wohler and Grunzweig^ have shown that there is inappreciable decomposition of the rare-earth sulfates used in this work even at a temperature as high as 650° C. The melting point values given for both sodium sulfate and potassium sulfate are several hundreds of degrees higher than the dehydration temperature used here, and there is no indication in the data that there is decomposition of them at or below their melting points. This temperature is also apparently safe to use with thallous sulfate for in the gravimetric determination of thallium in which thallous acid sulfate is formed and then heated to low red heat in order to transform the compound to 21 the normal sulfate, the temperature used is almost the same as is being used in the present work. The effectiveness of this temperature is further illustrated by the many analyses of the double salts prepared in the present study in which the calculated and analytically determined amounts of the constituents of the salt check so well. •^Hofmann and Hoschele, Ber.. 47:238, 1914. ^Wohler and Grxinzweig, Ibid.. 45:1727, 1912. *^Gooch, Methods in Chemical Analysis. pp. 219-220. 23 Determination of the Rare-Earth. The dehydrated material was then washed into a 150 ml. beaker using about 25 ml. of distilled water, and allow ed to dissolve as much as possible. In order to decompose the salt more readily, about one gram of sodium hydroxide was added and the solution was heated. After the salt had decomposed, and the rare-earth was precipitated as the hydroxide, an equal volume of water was added. The solution was then filtered and the precipitate was thoroughly washed. All washings were collected along with the first filtrate, and this solution was immediately acidified with hydrochloric acid in order to minimize any action the alkali hydroxide might have on the glassware. This filtrate contains all of the sulfate, and the alkali metal or thallium. Because of errors in the subsequent analysis for sulfate ion that the 22 use of nitric acid would introduce, hydrochloric acid was used for acidifying this filtrate. A white precipitate of thallous chloride may be found in this filtrate, but it will dissolve on heating and will not interfere with further steps in the analysis. The hydroxide of the rare-earth remaining on the 22 Willard and Furman, Elementary Quantitative Anal ysis . pp. 279-80, 1933. 24 filter was dissolved in dilute hydrochloric acid, filtered free from the disintegrated filter paper, and dilute ammonium hydroxide was added until a precipitate of the rare-earth hydroxide began to form. Dilute hydrochloric acid was added dropwise until this precipitate just dissolved, and then a very small excess was added. A double precipitation of the rare-earth hydroxide with ammonium hydroxide is sometimes needed to free this precipitate from entrapped sulfate ion. Experiment showed, however, that the amount of sulfate ion left in the hydroxide precipitate after a single precipita tion by sodium hydroxide amounted to an unweighable quantity. The volume of the solution at this point was about 50 ml. After heating nearly to boiling, a hot saturated solution of oxalic acid was added slowly and in excess in order to pre cipitate the rare-earth as an axalate. When the sample was very small and the amount of the rare-earth slight, it was found better to use a solution of oxalic acid prepared by diluting the saturated solution about three times, and add ing this reagent dropwise. If this method was not used, the precipitated oxalate tended to pass through the filter paper in a later operation. The precipitated oxalate was allowed to stand at least twenty-four hours to insure complete precipitation and to allow for some crystal growth, being then filtered and washed with distilled water to which had been added about 5 25 ml. of saturated oxalic acid for every 100 ml. of water. The washed oxalate was then ignited to constant weight in a platinum crucible. This ignition gives the rare-earth as an oxide, from which the amount of rare-earth sulfate can be calculated. All rare-earth oxalates used gave oxides having the general formula HgOg with the exception of praseo- 23 dymium oxalate which ignites in air to Determination of the Sulfate Ion. The acidified filtrate from the rare-earth separa tion was diluted or evaporated to 150 to 200 ml., the acid ity regulated so that there was an excess of about 0.5 ml. of concentrated hydrochloric acid, the solution heated to boil ing and the sulfate precipitated by adding dropwise and with constant stirring, 10 ml. of an 0.25 M. solution of barium chloride. The solution and precipitate were kept hot for at least an hour and occasionally stirred before being fil tered. The precipitated barium sulfate was filtered, and washed ten times with small portions of hot water. The pre cipitate and paper were then carefully dried before Igniting to constant weight in a platinum crucible. This gives BaSO^ from which the equivalent of S0^ can be calculated. ^Brinton and Pagel, J. Amer. Chem. Soc., 45:1460, 1923. 26 If the salt contained thallium, the filtrate and washings from the barium sulfate separation were carefully- saved and analysed for thallium. Otherwise, they could be discarded, as the sodium or potassium were not determined by direct analysis. Determination of Thallium as Thallous Iodide. The selection of a method of analysis for thallium suitable for this work caused a great deal of trouble and much loss of time. The choice of a method was dependent up on a consideration of the material to be analysed and upon the time required for the analysis. The analysis of one of the double salts prepared in this work was made upon a single sample rather than upon separate samples, one for each constituent to be determined. According to the method of analysis worked out for these salts, the determination of the thallium would be the last made with any sample. This meant then, that the solution to be analysed would contain in addition to thallium, barium and chloride ions left from the separation of the sulfate ion as BaSO^. With these conditions in mind, the methods recom mended for the determination of thallium were examined. The gravimetric methods that have been most commonly used are those in which the thallium is separated and weighed as an acid sulfate, a normal sulfate, a chromate, or as an iodide. The fact that there would be barium ion in the solu 27 tion to be analysed caused immediate rejection of the first three methods suggested because there would be in the preci pitate to be weighed a salt of barium. And because of the time involved in the analysis, the iodide method was also dis carded. Several volumetric methods are available for the determination of thallium. Most of these depend upon the valence change of the thallium from the thallous to the thallic state. After some consideration, three of these methods were selected as being of possible use. They also had the apparent rapidity that was desired. The first of these was a permanganate method, the second made use of potassium ferricyanide as an oxidizing agent, and the third possibility was an iodometric method. 24 Permanganate Method. - This makes use of potas sium permanganate as an oxidizing agent in an acid solution. Thallous ion is oxidized to thallic ion, and the end point is the appearance of the characteristic pink so common to per manganate titrations. In brief, the method consists of titrating an acid solution of thallous ion with a standard solution of potas sium permanganate. The volume of the solution titrated is kept large, and the acid used is hydrochloric. When this method was tried, a consistent value was obtained for the 24 Crookes, W., Select Methods of Chemical Analysis. p. 176. 4th ed. 28 amount of thallous Ion in the test solution. The value was too small, however, and the end point was frequently hard to determine. In agreement with work of Berry,^ it was found that the titration could not be done when sulfuric acid was substituted for the hydrochloric acid. This method was then discarded because of the difficulties and the possible errors entailed in the use of potassiurn permanganate in an hydro chloric acid solution. Ferricyanide Method. - Browning and Pa l m e r26 sug gested a method whereby the thallous ion is quantitatively oxidized to thallic oxide in an alkaline solution by means of potassium ferricyanide. The insoluble oxide is filtered from the solution, and the ferrocyanide formed is determined by titrating with potassium permanganate in an acid solution. In order to become familiar with the reaction re presented by this method, the titration of a ferrocyanide solution with standardized potassium permanganate was taken up. The titration is done in an acid solution. The end point in the titration is supposed to be the appearance of the characteristic pink color of permanganate titrations. 25A.J. Berry, J. Chem.Soc.. 121:394-399, 1922. 26p.E. Browning and H.E. Palmer, Amer. Jour. Sci.. (4) 27:379-330; Reprinted in Gooch, Method of Chemical Anal ysis . p. 220. In actual use, however, it was found extremely difficult to determine the first appearance of this pink color. Another means of determining the end point was found in the use of a small amount of a solution of ferric chloride used either in the solution being titrated or as an outside indicator. Used in either manner, the end point is reached when the blue color of the ferric-ferrocyanide compound disappears and this point corresponds to the complete oxidation of ferrocyanide to ferricyanide. The potassium permanganate solution was standard ized against sodium oxalate. It was then standardized against a solution of known thallium content by the following procedure. To a known volume of the standra'd thallium solution in 100 ml. of water was added a slight excess of potassium ferricyanide and sufficient potassium hydroxide to cause the complete oxidation and precipitation of the thallic oxide. The precipitated oxide was removed by filtering, and the pre cipitate washed thoroughly. The filtrate and washings were acidified with sulfuric acid, the volume of the solution made to exactly 250 ml., and aloquot portions titrated with the standard potassium permanganate solution using ferric chloride as an outside indicator. After many unsuccessful attempts, consistent values for the potassium permanganate in terms of thallium were obtained. 50 When this method was applied to the analysis of some of the double sulfates prepared in this work, the results were so erratic that the method was discarded. At this time, the work had to be discontinued for a period of several months. During this period of inactivi ty, a rather detailed study had been completed dealing with the use of several methods suggested for the determination 27 of small amounts of thallium. One of the results of this study was the finding that the gravimetric iodide method for the determination of small amounts of thallium was an exceed ingly accurate one. Since the conditions of the solution to be worked with were compatible with those required for the use of the iodide method, this procedure was adopted without further attempts to find a volumetric method. The excellent results obtained in the analysis of the double sulfate salts when using this iodide method more than offset the fact that approximately three days were required for the completion of the thallium determination. The combined filtrate and washings from the sulfate separation were evaporated to a volume of 50 ml. When this solution cools crystals of thallous chloride will probably ^Paul Summers, A Study of Methods for the Deter mination of Thallium in Toxicological Analysis. (unpublished Masterfs thesis, University of Southern California, 1956). 31 be noted. The solution was then made just alkaline to phenolphthalein using ammonium hydroxide, and then slightly acid with 6 N. acetic acid. The solution was next heated almost to boiling and while constantly stirred, 5 ml. of 0.1 N. solution.of potassium iodide was added dropwise. Precipitation from a boiling solution is a departure from the general method, but it gives larger crystals of the thallous iodide and these have less tendency to pass through the filter. With only a small amount of thallium present, a precipitate of thallous iodide will probably not be observ ed as long as the solution is hot. On cooling, however, long yellow to orange needles of thallous iodide form. After allowing eighteen hours for the salt to form, the thal lous iodide is filtered off on a weighed Gooch crucible and heated in an electric oven at 105° C. to constant weight. The equivalent amount of the thallous sulfate originally present can be calculated. A list of the salts prepared during this study is found in Table II. The arrangement of the salts is the same as that used in the discussion of the preparation that follows. With the exception of six of the salts, none of them has been prepared previous to this work. The scarcity of information available concerning the double sulfates of the rare-earth metals used and the alkali elements is clearly indicated by the very few salts 32 prepared and discussed in the literature. Along with each salt discussed in this study there is given the necessary experimental data from which its composition was determined. TABLE II SALTS PREPARED 1. Sm2(S04)3.4.5K2S04 2. SmgCSO^g.BKgSO^HgO 3. Sm2(S04)3.2K2S04.5HgO 4 • Sm2(S04)3.K2S04.4Hg0 *5, 2Snig ( SO^) 3. 3NagS04. 6Hg0 6* 3Smg(S04)3.4NagS04.8HgO *7. Sm2(S04)3.NagS04.2Hg0 8, Sm2(S04)3.Tl2S04.8Hg0 9. Sm2(S04)g.TlgS04.7HgO 10. 2Smg(S04)5.2T12S04.7Hg0 11. Sm2(S04)3.Tl2S04.4H20 12. 2Prg(S04)g.8KgS04.SHgO *13. PrgCSO^g.SKgS^.HgO 14. Pr2(S04)g.Na2S04 •3HgO *15. Pr2(S04)g.Na2S04.2Hg0 *16. P^2(S04)s-T:L8S04-4H20 17. Pt2(S04)3.T12S04.5H20 *18. pr2(S04)3.(NH4)2S04.8Hg0 19. 3YM S04>3-T12S04-H2° 20. Eu2(S°4)3.4KgS04 21. Eu2(S°4)3.4.5K2S°4 22. 2EUg(S04)3.3Na2S04.6HgO TABLE II (continued) 25. 4Eug(S04)3.5Na2S04.8Hg0 24. Eu2(S04)3.Tl3S04.7Hg0 25. 2EUg(S04)3.3TlgS04.10Hg0 26. Gdg(S04)g.2KgS04.2Hg0 27. 4Gdg (S04) g. 9KgS04.14EL>0 23. Gdg (S04)g.2N&gS0^•8HgO 29. 4Gdg(S04)3.5NagS04.8H20 30. Gdg(S04)3.TlgS04.8Hg0 * The six starred salts have been previously pre pared by others and their existence is mentioned in the literature. Double Sulfates of Samarium and Potassium. One salt of samarium and potassium has been re- po ported. Cleve stated that the salt SSmg(SO^g.SKgSO^.SH^O was only slightly soluble in water. Smg(SQ4)g.4.5KgSQ4. - This is the only anhydrous salt of these two elements prepared during this work. The first attempt at its preparation was a failure. About 50 ml. of a saturated solution of samarium sulfate was used and solid potassium sulfate was added until the solution was saturated with this salt. A precipitate started to form during the addition of the alkali sulfate. This solution was set aside at room temperature for only a few hours. When filtering and washing of the precipitate was started, a major portion of the salt dissolved. It was noted, how ever, that two apparently different salts had formed: one was a clear, colorless crystalline salt, possibly potassium sulfate; and the other was much finer, white in color, and was probably the double salt desired. On repeating the preparation, solid potassium sulfate was added until the clear crystalline material form ed. The mixture was then set aside at room temperature for a few days less than a month. At the end of this time, 28Cleve, Bull. Soc. Chlm.. (2) 43:166, 1885. 36 only a finely divided, white salt was seen. This was wash ed thoroughly and dried over anhydrous calcium chloride for five days. The salt is illustrated in Figure 1. The magni fication required was 430x. Few whole crystals were found, and only one is shown in the illustration. This does showr, however, the six-sided, plate-like form shorn by so many of the other double salts prepared during this work. Calculated Found I II Smo0, 25.67$ 35.34$ 35.33$ SOg 6 43.73$ 43.39$ 43.35$ The solubility is given in Table III. Sttig(S04)g.5KgS04.HgO. - Equal volumes of saturated solutions of samarium and potassium sulfate were mixed, and the resulting solution was evaporated to half-volume at 45° C. The result was a finely divided, white crystalline salt. It is shown in Figure 2 magnified lOOx. Only parts of the crystals are seen, but they indicate the six-sided, plate-like form of the crystals. Calculated Found Sm„0, 23.60$ 24.42$ S05 43.32$ 43.34$ HjD 1.22$ 1.23$ 57 Smg(SO^)g• • 5HgO. - This salt was prepared by allowing a mixture of equal volumes of the saturated solu tions of the sulfates of samarium and potassium to stand at room temperature for a period of ten days. At the end of this time, a pale yellow crystalline salt had separated. The particles were small in size, but, as shown in Figure 5, were definitely crystalline and more prismatic than the preceding salts of these two elements just described. In form, they are similar to the hydrated samarium sulfate. In addition to being finely divided, many of the crystals grew in the form of rosette-shaped groups similar to the rosettes shown in Figure 4. and other crystals shown in following figures is due to the washing that they undergo before being dried and analysed. The solubility is shown in Table III* Sm„(S04)3.K2S04.4H20. - This salt represents the material that formed in a solution prepared by mixing equal volumes of saturated solutions of the sulfates of samarium and potassium and evaporating slowly at 45° C. for a week. Much of the rounding off of the edges of these Calculated Found H?8 33.95$ 38.95$ 8.76$ 35.55$ 39.43$ 8.15$ In color, the crystals are a pale yellow. As 58 shown in Figure 4, the salts grew both as single crystals and in small rosette-like groups. The larger crystal shown illustrates the tendency for the salt to form as flat, clear crystals. Calculated Found Sm?0~ 41.76# 41.47# SOg 38.34# 39.45# HgO 8.62# 5.28# The solubility of this salt is given in Table III. Double Sulfates of Samarium and Sodium. The salt, BSmgCSO^g.SNagSO^.SHgO was prepared by 29 Keyes and James, while they were studying the solubility, at 25° C., of samarium sulfate in solutions of sodium sul fate. The curve they obtained indicated only this one salt. They then prepared the salt by mixing saturated solutions of samarium and sodium sulfates. According to their v/ork, 30 this is the only salt existing at 25° C. Restaino, study ing the isotherm of the system SmgCso^^jNagSO^HgO at 25° C., found that the two salts, 4Smg(S0^)g.5NagS04.8Hg0 and *^See Reference 7 30S. Restaino, Atti. accad. Lincei. 20:192-200, 1934; Chem. Abstr.. 29:1028, 1935. Figure 1 Figure 2 Sing (S 3 • 4 • 5iCg3 04 *^^2 ^^2 ^4 * ^2 ^ Figure 3 Sm2(S04)3*2K2S04.5HgG Figure 4 Sm2(SC4)5.K2S04.4H20 40 Smg(S04)3.NagS04#2Hg0 were present in the solid phases of 31 the system. Cleve also reports having prepared this lat ter salt. 2Smg(S0^)g.3NagS0^.6Hg0. - In order to prepare this salt, equal volumes of the saturated solutions of the sulfates of samarium and sodium were mixed, and the result ing solution was evaporated at 45° C. to three-fourths of its original volume. The salt separated as a finely divid ed, white, crystalline material. The salt is illustrated in Figure 5. This required a magnification of 430x. The salt is seen to consist of long prismatic crystals apparent ly belonging to the monoclinic system. Calculated Found SmpOg 40.75$ 41.29$ SOg 42.08$ 42.52$ HgO 6.31$ 5.95$ 3Smg(S04)^.4NagS04.8Hg0. - Saturated solutions of the sulfates of the rare-earth and the alkali were mixed and set aside at room temperature. After three weeks an extreme ly fine, very pale yellow salt had formed. In form, the crystals appear like those of the salt 2:3:6 shown in Figure 5. Figure 6 shows the salt, but only parts of a few crystals Cleve, Bull. Soc. Chim., (2) 43:53, 1885 Figure 5 2Sm2(S G4)3♦3Na2S04.6H2Q Figure 6 3Sm2(S04)3.4Ha2S04.8H2 0 Figure 7 Sm2(S04)3.TX2S°4.SH20 Figure 8 Sm2(S04)3 #T12S04•8H20 4 2 are seen. As with the preceding salt, a magnification of 430x was necessary in order to see the crystal form. This salt also appears to crystallize in the monoclinic system. Calculated Found ShuO- 43.£1$ 42.01$ SOg 41.98# 42.08# HgO 5.81# 5.85# The solubility of this salt is shown in Table III. Sm^SO^g.NagSO^.SHgO. — Starting with a saturat ed solution of samarium sulfate, solid sodium sulfate was added until precipitation started, and the mixture was then set aside at room temperature for a week. By this time a very fine white powder had formed. Attempts to photograph this material were without success, even magnification to 430x failed to resolve the powder into recognizable particles. Calculated Found Smg03 45.48# 42.75# S03 41.75# 42.08# H20 4.70# 5.11# The solubility was determined and is shown in Table III. Double Sulfates of Samarium and Thallium. Smg(S0^)g.TlgS0^.8Hg0. - This salt, pale yellow in color and usually clear, was prepared by several different methods. The crystals obtained ranged in size from less than 43 0.5 mm. up to 10 mm. a* A solution composed of equal volumes of the saturated solutions of the sulfates of the rare-earth and thallium was held at room temperature for a week. The crystals formed were from 1 to 3 mm. in length. porating at 45° C., a solution prepared by mixing saturated solutions of the sulfates of samarium and thallium in the ratio of 4:1. This preparation was started originally at room temperature; however, at the end of a month no salt had formed, so that the solution was then evaporated slowly at the increased temperature. the ratio of the samarium to the thallium sulfate 1:8, the same salt was obtained. Only ten days were required, and the crystals obtained varied in size from less than 1 mm. up to 10 mm. All were of the same form. The analysis of this sample is shown below. Figure 7 at 40x and in Figure 8 at 4x. In form, the crystals are thick and long. b. The same salt was obtained also by eva- c. Using the same procedure as in b, but with The salt prepared by these methods is shown in Calculated Found S03 TlsO HgO SnigOg 28.18$ 28.50% 25.87$ 26.02$ 34.32$ 34.18$ 11.64$ 11.50$ Solubility values are given in Table III 44 Sm2(S04)g.TlgS04*7Hg0. - In order to prepare this salt, 150 ml. of a saturated solution of samarium sulfate was saturated with thallium sulfate. This mixture was then set aside at room temperature for a week. At the end of this time a clear, pale yellow, finely divided, crystalline mater ial had formed. The salt is shown in Figure 9 as a six- sided, flat crystal. In Figure 10, the crystal is still six- sided, hut more prismatic in form. Calculated Found I II Smp0, 28.60$ 28.50$ 28.73$ SO3 0 26.25$ 26.06$ 26.34$ TlgO 34.83$ 34.39$ 34.64$ HgO 10.34$ 9.91$ 9.86$ The solubility of this salt is given in Table III. The mother liquor remaining after this salt had been removed was evaporated slowly at 45° C. and another crop of crystals obtained. These were clear and pale yellow. Anal ysis showed that they were the salt 1:1:8 already described. Calculated Found Smp0„ 28.18$ 28.30$ SO3 6 25.87$ 26.02$ TloO 34.32$ 34.01$ HoO 11.64$ 11.56$ The two following salts were prepared and photo graphed, but the analyses are incomplete, so the existence of the salts must be regarded with less certainty than in the Figure 9 SiBg (8B4) 3. TlgS 04. 7HgO Figure IX 4)3.2TlgS04.7Hg0 Figure 12 Smg (S 04) 3 . TlgSG4 .'45^ 0 Figure 10 Smg(S 04)3.TlgS04.7Hg 0 46 case of some others. SSnig(S0^)g.2TlgS0^.7Hg0. - This salt was prepared by slowly evaporating at 45° C. a solution prepared from equal volumes of saturated solutions of the sulfates of the two elements. The salt is a pale yellow in color. The analysis of this salt was based only upon a determination of the rare-earth and of water. For this reason, the anal ysis may be of doubtful value. The salt is shown in Figure 11. A sub-stage dark field ring, and a magnification of 40x were used. Calculated Found SnuO, 30.16$ 38.96$ HgO ° 5.45$ 5.43$ Smg(S0^)g.Tl2S0^.4Hg0. - Equal volumes of the sul- fates of samarium and thallium were mixed, and the resulting solution was evaporated slowly at 60° C. After a few days, a clear, pale yellow salt had formed. Examination of the salts shown in Figure 12, shows that the srystals are six- sided, thick, flat prisms. Evidently the salt has crystal lized in the monoclinic system. Analysis of the salt, as with the preceding one, was for rare-earth and water only. Calculated Found SmpO, 39.93$ 38.18$ HgO 6.18$ 6.15$ 47 Double Sulfates of Praseodymium and Potassium. The isotherm at 25° C. of the system Pr2(S04)3: <Zp KgSO^iHgO was determined by Restaino. His results indi cate the presence of these six salts: (l) Pr2(S04)s.5KgS04. HgO, (2) Pr2(S04)3.4.5K2S04, (3) Pr2(S04)3.KgS04.2Hg0, (4) Pr2(S04)3.3KgS04.2Hg0, (5) 2Pr2(S04)3.3K2S04.8Hg0, and (6) PTg(S04)3,4K2S04.Ho0. He finds that this series corres- ponds exactly to the series between cerium and potassium. 33 The specific gravity of (4) is reported by Von Scheele to be 3.29, who further reports that the salt Prg(S04)3.3KpS04. HgO is only slightly soluble in water. The salt Rr2(S04)s# KgS04.4Hg0 was prepared in two different ways by Baskerville 34 and Holland. These co-workers were able to prepare it by the electrolytic method and also be evaporating in a desic cator a solution containing the sulfates of the rare-earth and potassium. 2Pr2(S04)3.8K2S04.3H20. - The formation of this salt took place immediately on mixing equal volumes of sat urated solutions of the sulfates of praseodymium and potas sium. This was let stand for only a day before filtering ^See Reference 30 33 Von Scheele, Ztschr. anorg. Chem., 17:310, 1898. ^See Reference 3a 48 and washing. The salt is a light orange in color. That the crystal is of a fair size is shown by Figure 13, in which the magnification is only lOOx. As with several of the other salts prepared, the crystals are flat and six- sided, and appear to be crystallizing in the monoclinic system. The solubility of this salt is given in Table III. PrgCsO^Jg.bKgSC^.HgO. - The preparation of this salt was the same as for the preceding salt. Saturated solutions of the sulfates of praseodymium and potassium were mixed at room temperature. Precipitation of the salt was immediate. The salt, although the method of preparation was the same, was a pale yellow color in contrast to the orange of the other salt. Most of the salt was much finer than the first as is shown in Figure 14 using 430x magnifi cation. That the crystal form is similar to that of the 2:8:3 salt is shown by Figure 15 showing a single crystal enlarged 430x. The analysis suggests contamination of the crystals by a little praseodymium sulfate, which would show results for praseodymium a little above, and for sulfate a little below, the theoretical figure. Calculated Found 25.50$ 84.85$ 45.39$ 2.30$ HgO 43.31$ 2.09$ h d C O C O o I # * * X j oTo? • £ C n * ■ * C O H o C n m C O o i V f / J o • * . * . ! £ , A & . jf • ; / j m / 53* * > 50 Calculated Found 32.45% 24.35% 43.60% 41.65% 1.23% 1.09% The solubility of this salt is given in Table III. Double Sulfates of Praseodymium and Sodium. References to only two salts of praseodymium and sodium are found in the literature. While determining the isotherm of the system Pr2(S04)3:NagS0^:Hg0 at 25° C., Zam- bonini^5 found that the two salts Pr2(S04)3.Na2S04.2Hg0 and 4Prg(S04)g.5NagS0^.8H20 were present in the solid phases of the system. Beyond the fact that he found both salts to lose water with difficulty, nothing is given concerning the properties of either of these salts. immediately on mixing equal volumes of the saturated solu tions of the sulfates of praseodymium and sodium. The salt is greenish-yellow in color, very finely divided, and definite ly crystalline. Figure 16 shows this compound, the magnifi cation being 430x. Pr2(S04)3.NagS04.5Hg0. - This salt precipitated Calculated Found 43.05$ 41.80# 7.05# 41.28# 41.78# 6.67# ^See Reference 17 51 Prg(S0^)3.NagSC>4.2Hg0. - In order to prepare this salt, an excess of a solution of sodium sulfate was added to a solution of praseodymium sulfate. A fine yellow precipitate formed immediately. This was left for a week, in order to promote crystal growth. The resulting salt was in appearance a fine yellow powder. At 430x magnification it was found to be composed of long prismatic crystals as shown in Figure 17. Calculated Found 44.09$ 42.69$ 42.81$ 42.03$ 4.81$ 5.32$ The solubility of this salt is given in Table III. Double Sulfates of Praseodymium and Thallium. The system Prg(S0^)s:TlgS0^;Hg0 at 25° C. was studied by Zambonini and Restai.no.®® They found four dif ferent double sulfates of these two elements to be present in the solid phases of this system. Beyond giving the com position of the salts, no other data concerning them is re corded. The salts they report are: Prs(S04)3.5TlgS04, Pr2(S0 )3.4.5T12S04, Pr2(S04)3.3Tl2S04.Hg0 and Pr2(SC>4)3. T12S04.4H20. PrpO-r sol 3 Hg8 30 F. Zambonini and S. Restaino, Atti. accad. Lin- cei, 13:650-654, 1931; Chem. Abstr.. 26:1204, 1932. Figure 17 Prg(S04)3.NagSO4.2H20 Figure 18 Pi*q (S04)3* g•TlgSO^•4Hgu Figure 19 Pr2(S04)3.Tl2S04.4H20 Figure 20 Pr2(S04)g.(KH4)2S04#8E20 53 Prg(S04)3*Tl2S04.4HgO. - With a single exception, every attempt to prepare double salts of praseodymium and thallium resulted in the formation of this particular salt. In every case, too, the crystals formed were of a good size and well formed. Figure 18 shows some of the crystals mag nified 40x, while in Figure 19, the magnification is only 4x. In both, the form of the crystal is seen to be thick and short. a. On mixing equal volumes of saturated solu tions of praseodymium and of thallium sulfate and setting aside at room temperature, crystals were obtained at the end of five days. These were about 2 mm. in length, and of a clear orange color. Analysis of this salt was for the rare- earth and water only. b. The same method as in a repeated required almost a month before a clear, pale greenish-yellow salt was obtained. c. The same salt was prepared by mixing equal volumes of saturated solutions of the sulfates of praseo dymium and thallium, and holding at 45° C. until crystalliza tion took place. The salt was similar in color and appear ance to that obtained from b above, except that the crystals were much smaller. d. Using equal volumes of the saturated solu tions of the sulfates again, and holding at room temperature 54 for three weeks, two different salts were obtained. After only a few days, a crop of very small crystals was obtained. These were left in the solution and they gradually grew in size until they were about 2 mm. in length. Along with these clear crystals, however, there were some apparently non-cyrstalline particles. After washing and drying, these were found to be opaque and a pale yellow in color, very much in contrast to the clear, pale yellow-green crystals ac companying them. These two salts were separated and anal yzed. The clear crystals were the 1:1:4 salt; and in spite of their apparently dehydrated appearance, the opaque yellow material was found to have the composition represented by the formula: Pr2(SO^J^.Tl^SO^.SH^O. ^Average of seven analyses. The solubility of the salt is given in Table III. Pr2(SO^g.TlgSQ^.bHgO. - The preparation of this salt has already been explained in the preceding discussion. Attempts to photograph this salt were without success. The material appears to be amorphous. Calculated Found* 28.77$ 28.38% 27.92% 28.11% 37•04$ 6.28% 35.71$ 6.19$ TI2O H20 55 Calculated Found 27.87% 27.33% 28.32% 27.49$ T120 HgO 36.46# 7.73% 34.83% 7.83% The solubility of this salt is given in Table III. Double Sulfates of Praseodymium and Ammonium. Zambonini and Restaino3^ report that after study ing the equilibrium curves of the system Prg(S04)3:(NH4)gS04: HgO, only two salts were indicated. These had the formulas two salts are among the few for which crystallographic data latter salt is only slightly soluble in water, and that it has a specific gravity of 2.53. In addition, he indicated that he had prepared the salt Prg(S04)3.3(NH4)2S04.8Hg0, but there is no further data given. equal volumes of saturated solutions of the praseodymium and ammonium sulfates wexe mixed and the resulting solution was allowed to evaporate slowly at room temperature. At the end of two months, the crystals had formed and grown to a good 3?F. Zambonini and 8. Restaino, Ibid.. 11:774-779, 1930; Chem. Abstr.. 24:5656, 1930. Pr2(S04)3.(HH4)gS04.8H20 and Pr2(S04)3.5(NH4)2S04. These 38 has been determined. Von Scheele indicated that this PrgCsO^g. (NH^gSO^SHgO. - In preparing this salt, 3^See Reference 33 The analysis of this salt required a slight change in the general method used in the analysis of all of the other salts prepared. Analyses were made only for the rare- earth and sulfate present. The water of hydration had to be determined by difference since heat sufficient to dehydrate the salt also decomposed it. it is a clear light green. The salt apparently crystallizes in the monoclinic system, and the individual crystals are flat with little thickness. Because of the small amount of salt prepared, no solubility determination was attempted. Double Sulfates of Ytterbium and Thallium. 3Ybg(SO^Jg.TlgSO^.xHgO. - Only an extremely small amount of the rare-earth oxide was available so that just the one salt was prepared. This was made by adding sufficient powdered thallium sulfate to almost saturate a solution of the ytterbium sulfate. The resulting solution was set aside at room temperature to slowly evaporate. The salt had form ed by the end of two weeks. The crystals of the salt are clear and colorless. In size, they varied from less than a millimeter in length to almost three millimeters. As Figure The salt is illustrated in Figure 20. In color, Calculated 38.98$ 37.84$ Found 38.52$ 38.22$ Figure 21 3Yb2(S04)3.T12S04.xHg 0 Figure 22 Eug(S04)3•4KgS04 21 shows, they are prismatic in form. There is some doubt as to the degree of hydration of this salt for analyses varied widely for water. An insufficient amount of salt was avail able for final determination of the wr ater content and for a solubility determination. Double Sulfates of Europium and Potassium. Eug(S0^)g.4KgS0^. - The unusual fact concerning this salt and following one is that the salts are both anhydrous and yet the same methods of preparation were used that resulted in hydrated salts when sodium or thallium were used with europium. urated solution of europium sulfate with solid potassium sulfate and holding the resulting solution at 45° C. After a day, a white, finely divided material had settled out. Apparently a powder, it shows when enlarged a definitely crystalline form. Figure 22, at lOOx magnification and using the sub-stage ring, shows only one complete six-sided crys tal, a few broken pieces, and the blurred masses commonly noted. Calculated Found I II Tl2° 46.66$ 46.03$ 45.41$ 31.60$ 30.99$ 31.04$ 16.77$ 17.03$ This salt was prepared by almost saturating a sat- 59 Calculated Found I II 27.50$ 27.47$ 26.92$ 43.47$ 43.42$ 43.19$ 1.06$ 0.94$ The solubility of this salt is given in Table III. EugCSO^J^.^.bKgSO^. - The mixture of the two sul fates was prepared as for the preceding salt. The resulting solution was maintained at room temperature. Three days were necessary for the salt to form. Again, finely divided and white in color, the salt was apparently a powder. Fig ure 25 taken at 40x magnification with the sub-stage ring in place, shows parts of what are evidently clear, six-sided crystals. Careful examination of the material failed to yield any complete crystals. Double Sulfates of Europium and Sodium. SEugCSO^^.SNagSO^.BHgO, - Sufficient solid sulfate was added to saturate a saturated solution of europium sul fate, and this solution was then kept at room temperature. After three days a precipitate had formed. As shown in Fig ure 24, taken at lOOx magnification, the salt apparently has no distinct crystalline form. Even at 450x magnifica tion, the material still appeared as a fine powder. Calculated Found 25.58$ 25.57$ 43.63$ 42.82$ 42.82$ 0.48$ Hb0 60 Calculated Found 40.85$ 41.93$ 6.29$ 41.15$ 42.39$ 6.79$ 4Eug(S04)g.5NagS0^.8Hp0. - Starting with a solution prepared in the same way as for the preceding salt, the temperature was maintained at 45° C. At the end of only a day, a fine white precipitate had formed. For a long time it seemed that this was merely a fine powder, and attempts to note its form at high magnification resulted only in dark, opaque masses. In an attempt to reduce the amount of pow der on the slide, a finger was used to remove the powder. Figure 25 shows the result of this operation. This picture clearly shows that the salt is crystalline, and that the crystals are long and prismatic. Beyond this, it is not possible to go, since the magnification used to obtain this photograph represents the limit of enlargement that was ob tainable • Double Sulfates of Europium and Thallium. Eug(S°4)5.ILgS04.7Hg0. - This salt was prepared in two different ways. A saturated solution of europium sul- Calculated Found 43.69$ 42.23$ 4.47$ 43.71$ 42.37$ 4.83$ Figure 25 4Eu2(S04}3.5KagSQ4.8Hg0 Figure 26 Eug(S04)g.T12S04.7HgO s \ \ - a Figure 27 Eu2(S°4)3.T12S04.7H20 ■ Figure 28 2Eug(S04)3.3T12S04.lOHgO 62 fate was added to just saturate it. The resulting solution was left at room temperature for five days. At this time, a white, finely divided crystalline salt had formed. This material was removed, washed, and dried. The mother liquor and washings were united and slowly evaporated at 45° C. After two days, another crop of crystals was obtained. These last crystals, however, were larger in size than the first ones obtained. Analysis showed that both crops of crystals had the same composition. Figures 26 and 27 show these salts. Figure 26 was taken at 40x magnification, tising a sub-stage dark field ring. The illustration is of one of the larger crystals obtained. Figure 27, using lOOx magnification and the dark field illum ination, shows some of the smaller crystals. These are clear and appear similar in form to some of the other salts prepared. As can be seen, the form is that of a prism, and the crystallization is apparently in the rnonoclinic system. Calculated Found I II EupO, 28.78$ 28.74$ 29.03$ S0§ 0 26.18$ 26.21$ 26.34$ TlpO 34.73$ 34.48$ 35.17$ HgO 10.31$ 10.10$ 10.40$ The solubility is given in Table III. 2Eug(S04)3.3TlgS04.10Hg0. - The preparation of this salt followed one of the methods used to prepare the preced- 63 ing salt. A saturated solution of europium sulfate was just saturated with solid thallium sulfate, and the resulting solu tion was slowly evaporated at 45° C. for a period of five days. By this time, a clear, colorless salt had formed. Individual crystals range in size from 0.5 mm. to 3 mm. in length. Figure 28 shows this salt. A magnification of 40x was used, a Wrattan Filter B, and a sub-stage ring. In form, the salt appears to be similar to that of the 1:1:7 salt. (Figure 27) Calculated Found 24.33$ 24.12% 25•03$ 24.64$ 44.26$ 43.75$ 6.26$ 6.27$ Double Sulfates of Gadolinium and Potassium. The only data available concerning salts of gado linium and potassium, gives the meager information that the 39 salt, Gdg(SO^Jg.KgSO^^HgO is sparingly soluble in water. Gd2(S04)g.2KgS°4.2Hg0. - A week at room tempera ture was required for the salt to separate from a solution made by saturating a saturated solution of gadolinium sulfate with potassium sulfate. The salt was white and very finely divided. Under the microscope, fragments of crystals were so|°3 TloO Hg6 s9Benedicks, Ztschr. anorg.. Chem.. 22:409, 1900 64 observed along with only a few complete crystals. Figure 29 shows one of these at lOOx magnification and with the sub-stage ring in place. It appears as a long prismatic crystal belonging to the monoclinic system. Calculated Found I II GdpO^ 36.73# 36.42# 37.01# SOg ° 40.55# 40.25# 40.84# KgO 3.65# 3.73# 3.93# 4Gdc>(SO^)g.9KgSO^_.14i3gO. - Starting with the same type of solution as for the preceding salt, five days at 45° C. were needed for the formation of this salt. The salt is clear and colorless. No complete crystals were observed. Magnified, it appears as plate-like crystals with six sides and probably belonging to the monoclinic system. Figure 30, at lOOx and using the sub-stage ring, shows portions of crystal particles. Calculated Found GdP0, 34.28$ 34.29$ S0| ° 39.73$ 33.51$ HgO 5.96$ 5.89$ The formula of this salt appears somewhat unusual, especially w^ith respect to the high degree of hydration. It appears to be a definitely crystalline material as is shown in the illustration, and it has been inserted for the sake of completeness and not as a salt whose existence has been definitely shown. " P M P"nT*fi Gdg (SO^)3.2KgS04.2H20 rr I* s <■•••■ < L r t-'rWam* Figure 31 4Gdg(S04)3.5NagS04.8Hg0 Figure 30 4Gag(S04)3.9KgS04.14Hg0 Figure 32 Gdg(S°4)3.TigS°4.8Hg° 66 Double Sulfates of Gadolinium and Sodium. There is only one double sulfate of gadolinium 40 and sodium reported in the literature. Bissell and James studied the solubility of gadolinium sulfate in a solution of sodium sulfate. The curve they obtained indicated only one double salt with the formula Gdg(S04)g.NapS04.2Hg0. No further information is given concerning this salt and there is no indication that these workers actually prepared this salt. Gd2(S04)3.2NagS04.3Hg0, - The salt was the result of setting aside at room temperature for a period of two weeks, a solution prepared by saturating a saturated solution of gadolinium sulfate with sodium sulfate. The material formed was white in color and powdery in form. High magni fication did not resolve the salt into any definitely crystal line form. Calculated Found Gd«0s 38.54$ 39.54# SO3 42.54# 42.40# H20 5.74# 5.37# The solubility of this salt is given in Table III. 4Gdg(S04)5.5NagS04.8Hg0. - Instead of holding the solution at room temperature as in the preceding preparation, ^^See Reference 8 67 a similar solution was kept at 45° C. for a day. The result ing material was a fine white powder. It was only after rub bing most of the salt off of a slide with a finger that the crystalline character of this compound was noted. Figure 31 shows the salt magnified 430x. It is evident that the salt is definitely crystalline and that the crystals are prismatic in form. The salt is similar in appearance to the analogous salt of europium and sodium. (Figure 25) Calculated Found GdpCU 44.42$ 43.86$ SO* ° 41 • 68$ 41.295? 4.41$ 4.58$ Double Sulfates of Gadolinium and Thallium. GdgCSO^g.TlgSO^.SHgO. - This salt was prepared by adding sufficient thallium sulfate to almost saturate a sat urated solution of gadolinium sulfate. This mixture was set aside at room temperature. It required three weeks for the formation of sufficient salt with which to work. The salt formed was a mixture of white and of colorless clear crystals. Figure 32, taken at 40x magnification and using the sub-stage ring, shows the two kinds of crystals obtained. Both forms however, are prismatic and long. They appear to belong to the monoclinic system. 68 Calculated Found 28.97# 29.07# 25.58# 24.45# 33.94# 32.09# 11.51# 11.88# Using the same type of mixture as used for the preceding salt, but evaporating slowly at 45° C., the salt formed was Tl^SO^. G d o O ^ so§ 3 T120 HgO PART III SOLUBILITY DETERMINATION The concluding part of this problem was to be the determination of the solubility in water of as many of the salts prepared as was possible. There are at present no data in the literature concerning this particular property of the double sulfate salts of the rare-earth elements and the alkali elements. It was believed, therefore, that the solubility data to be determined would add to the present available knowledge of these salts. Several questions are awaiting answers: 1. There is apparently a definite order of solu bility of the rare-earth double sulfates in alkali sulfate solutions. Does this same order hold with their solubilities in water? 3. There is probably a change in the solubility of the salts varying with the alkali element combined with the rare-earth in the double sulfate salt. Does this change appear if water is used as the solvent? 3. Thallous sulfate has been suggested as a reagent for separating the rare-earth elements because of the greater solubility differences between the double sulfates formed. Does this order of difference also appear when water is used as the solvent? 70 4. There is no data available concerning the sol ubility of the double sulfate salts varying with the percent age of rare-earth element or alkali element present. Is there such a relationship to be found? The results of the solubility determinations made during this work are given in Table III. This data repre sents work done during two widely separated periods of time. This accounts for the two different temperatures found in the table. During the time that the first solubility work was being done, the room and water temperatures were low enough so that a temperature of 25° C. could be maintained easily in the thermostatically-controlled water bath. Dur ing the second period of time, however, the general room temperature was only slightly below 25° C. and it was found to be impractical to attempt to maintain the previous tempera ture used. Instead, the temperature of the constant tempera ture bath was maintained at 27.5° C., and it is at this tem perature that the majority of the solubility determinations were made. The choice of a period of seven days of rotation at constant temperature was based upon several factors. Some of the workers in the field of rare-earth compounds had rot ated their samples several months at a time in order to assure saturation or equilibrium.4^ - Others had used much shorter 4^See Reference 7 71 periods of rotation.^2 It was evident from the nature of the compounds with which the work was being done, that to attain equilibrium between the salts and their saturated solutions would require a longer time than if the salts were of lower-valence elements. From an examination of the sol ubility values obtained for some of the salts at 25° C., it is evident that the period of rotation, 3, 7, or 10 days caused no appreciable change in the solubility values pro vided the solution still contained excess salt. A period of seven days was thus chosen as being of sufficient length to assure saturation of the solution at 27.5° C. The results shown in Table III for those salts which were rotated for longer than seven days show that the shorter period had suf ficed for the attainment of equilibrium between the salt and the saturated solution. Apparatus. A constant temperature bath was available. It was made of wood and lined with sheet zinc. The inside dimen sions were 2x2x2 feet. A DeKhotinsky standard thermo-regu lator with a precision of 0.01° C in conjunction with a Cenco polarized type sensitive relay made up the temperature regu lating apparatus. A single 200 watt immersion type knife heater proved sufficient to maintain the temperature of the bath at the point desired. The bath was stirred by means ^See Reference 6a 72 of a Cenco centrifugal stirrer rated to circulate almost 100 gallons of water per minute. There was no apparatus on hand for holding the bottles and solution while being rotated in the bath. The apparatus finally constructed was based somewhat upon an ap- Ar z paratus used and described by Sarver and Brinton. In the absence of actual data concerning the solubility of the double sulfate salts in water, it was thought that 100 ml. of solu tion would be sufficient with v/hich to work. Accordingly, an apparatus was constructed to hold twenty-four 125 ml. glass stoppered narrow-mouthed reagent bottles. These bottles were supported on two separate frames. Each frame was equipped to support twelve bottles. A frame consisted of six radiat ing spokes, with each spoke carrying two bottles so arranged that as the frame rotated, the bottles were turned end over end. For many of the solubility determinations, the amount of salt available was very small and the 125 ml. bot tles were too large to use. Because of this, small, bake- lite screw-capped, glass vials of 20 ml. and 35 ml. capacity were used. These vials, not fitting the frames, were wired on instead. As with the bottles, a coating of paraffin was applied over the cap to further guard against leakage. The shaft to which these two frames were attached 4^Sarver and Brinton, J. Amer. Chem. Soc.. 49:945, 1927. 73 was made of 3/4 inch material and supported in a horizontal position by means' of two upright pieces of 2x4 inch material placed on opposite sides of the tank. The shaft was support ed and free to rotate in holes drilled in these two uprights. Near one end of this shaft, a large bicycle sprocket was at tached. A bicycle roller chain around this and around a smaller sprocket attached to a shaft directly above and out side of the bath, provided means for rotating the frames using a quarter-horse power motor. The frames were rotated at a speed of twelve revolutions per minute which was suffi cient to keep the salt and solution continually stirred and not so fast that the salt was kept at one end of the bottle through the action of centrifugal force. Each bottle was carefully checked to see that it did not leak. In order to more securely fasten the stoppers in place, a TVf shaped groove was cut in the top of the stop per. The stoppers could then be fastened on by means of wire looped around the neck and over the stopper. Experimental. Saturation of solution. - The salt was added to a bottle, the water added, and the stopper wired, on and well paraffined. The bottle was then rotated in the constant temperature bath for a period of seven days. At the end of this time, it was removed from one of the frames and placed in a settling rack in the bath. Here it stayed for a period 74 of twenty-four hours in order to give the excess salt a chance to settle out before sampling. Samples. - Samples of the saturated solutions were obtained by means of 10 and 20 ml. volumetric pipettes. In order to be certain that the sample obtained contained no sus pended material, the liquid was filtered through a plug of dry cotton in a short piece of tubing attached to the end of the pipette. The filtered solution was examined carefully in a strong light for any evidence of opalescence or of sus pended particles, and then placed in a previously weighed 50 ml. Erlenmeyer flask. The flask and its contents were again weighed, and the difference was taken as the weight of the sample of solution. Analysis of Samples. - The weighed sample of solu tion was then ready for analysis. For the purpose of this work, the total rare-earth element was determined as the oxide, and the total sulfate ion determined as barium sulfate. By means of factors calculated for each of the salts used, the solubility of the salt could then be calculated on the basis of two different substances present. In this way a check was had upon the solubility determinations and possibly upon the original analysis of the salt as well. The weighed solution was transferred to a 250 ml. beaker, and made ammoniacal with dilute ammonium hydroxide. This precipitated the rare-earth element as the hydroxide. 75 This was removed by filtration and thoroughly washed. The method of analysis from this step was the same as that used in the analysis of the double salts after the rare-earth hydroxide had been precipitated with ammonium hydroxide. The hydroxide of the rare-earth was changed to the oxalate, ignited, and weighed as the oxide. The filtrate, plus all of the wash water from the separation of the hydroxide, was used for the determination of the sulfate as barium sulfate. In the case of those salts that contained thallium, no deter mination was made of this element. The results of these solubility determinations indi cate that the order of solubility of the double sulfates in water varies in the same general way as do the solubilities of the corresponding rare-earth sulfates. The double sul fates of praseodymium are definitely less soluble than those of samarium. Because of lack of sufficient data on the double sulfates of europium and gadolinium, it can only be stated that these salts are of about the same general order of solubility as those of samarium. From the molar concentration of the rare-earth oxides as given in Table IV, it is evident that for the salts of praseodymium, those containing thallium have a higher con centration of R0CL than those containing either sodium or d 3 potassium. The results are not conclusive as to the order of the salts containing sodium and potassium with respect to TABLE I I I SOLUBILITIES OF SOME RARE-EARTH DOUBLE SULFATES Salt Temp. °G. Sm2(S04) 5,K2S04.4H20 27.5 Sm2(S04)5.2KgS04.5H20 25.0 27.5 Sm0(SO .) .4.5HoS0. 27.5 2V 4 0 & 4 Sm2(S04)^.HagS04.2HgO 27.5 5Sm2(S04)^.4Na2S04.8H20 27.5 Smg(S04)b.TlgS04.7HgO 27.5 Smg(S04)b.TlgS04.8H20 25.0 27.5 2Smg (SO4) g. 2TlgS04 • 7Hg0 25.0 Prg(S04) .5K2S04.Hg0 27.5 2Pr2(S04)£ > .8K2S04.3HJ e 0 25.0 Pr2(S04)^.Na2S04.2Hg0 27.5 Pr (S04)5.Na2S04.3Hg0 25.0 Time Solubility-gm./lOO gm. HgO Days Based on Based on SO3 det. R2Q3 det 7 1 . 2 0 1 . 2 0 4 2.06 1.82 7a 3.07 2.70 7 2.51 2.96 14 2.53 2.90 7 2 .33 2.44 7 0.62 0.60 7b 0.62 0.60 7 0.67m 0 . 6 6 7 2.26 2.34 14 2.07 2.07 3 1.36 1.36 7 1.34 1.35 1 0 1.39 1.31 7 1.23 1.37 7c 2.06 2.34 14c 2.03 2.36 4 0.87 0.92 7d 1.34 1.29 7 0.56 0.63 7e 1.18 0.87 14e 1.13 1.03 3 0.25 0 . 2 1 5 0.28 0 . 2 2 9 0.25 0 . 2 2 1 2 0.25 0 . 2 1 7 0.36 0.37 1 0 0.36 0.36 3 0.30 0.27 5 0.32 0.25 9 0.31 0.28 1 2 0.32 0.28 TABLE I I I (continued) Salt Temp. Time Solubility-gm./lOO gm °C. Days Based on Based SO,, det. 0 R2°5 ' Pr2(S04)5.Tl2S04.4H20 27.5 7 0.99 1.08 7f 1.24 1.56 14f 1.25 1.58 Pr2(S04)5.Tl2S04.5H20 27.5 7 1.54 1.56 Eu2(S04)5.4K2S04 27.5 7+5g 5.15 27.5 7 2.95 2.97 Gd2(S04)5*2Na2S04‘5H2° 27.5 7+5g 1.54 1.56 a This solution was not saturated. b More salt was added after the first sample was taken, and again rotated for seven days, c More salt was added after the first sample was obtained, and again rotated for seven days. 14c is a contin uation of 7c. d This solution was still unsaturated after having added more salt, and no more salt was available for this purpose, e Same as for c. f Same as for c. g These samples were run for seven days at 27.5° G. During the time that they were placed in the settling rack, the thermostat regulator broke, and the temperature had risen to 40° before it was noted. The solution was re moved, more salt added, and rotated in the bath at 27.5° C. for an additional three days. 78 each other• The results with the samarium salts again indicate that those containing thallium give a higher concentration of RgOg than those containing sodium, that those containing potassium give a higher concentration of the oxide than those containing sodium, but that there is some question concerning the relation of the potassium and thallium salts. The data is insufficient to place europium and gadolinium salts. These results do not agree with the practical separa- 44 tion work of Fernandes who found that thallous sulfate was a better reagent for the separation of the rare-earths into smaller groups because of the lower solubility of the double sulfates with thallous sulfate and because the differences in solubility among the various earths are greater than for the corresponding potassium salts. 44 See Reference 2b TABLE IV MOLS OF RgOg PER 100 GRAMS OF SATURATED SOLUTION Salt Mols RgO; Sm2(S04)3.4.5K2S04 0.00178 Sm2(S04)3.2K2S04.5H20 0.00282 Smg(S04)3.KgS04.4Hg0 0.00144 3Smg(S04)g.4NagS04.SHgO 0.00108 SiDg ( S04) g«NagS04 • 3HgO 0.00081 Smg(S04)g.TlgS04.8Hg0 0.00191 S f f l g(S04)3.TlgS04.7Hg0 0.00170 2Prg(S04)3.8KgS04.3Hg0 0.00016 Prg(S04)3.5KgS04.Hg0 0.00070 Prg(S04)3.NagS04.3Hg0 0.00037 Prg(S04)3.NagS04.2HgO 0.00048 Prg(S04)3.TlgS04.4Hg0 0.00120 ^ V a - ’V W 0.00117 ^ ( 804)3.4^804 0.00243 Eug(S04)3.TlgS04.7Hg0 0.00243 Gdg(S04)3.3NagS04.3Hg0 0.00144 PART IV SUMMARY AND CONCLUSIONS A preliminary examination of the method of separat ing the rare-earths into two groups on the basis of the rela tive solubilities of the double sulfates of the various mem bers with the sulfates of sodium or potassium gave results so lacking in uniformity and concordance that this method, in the present state of our knowledge, could be considered only as a method of rough fractionation. The sharpness with which the members of the cerium group could be separated from those of the yttrium group by this method has probably been greatly overstated in the past. For the double purpose of forming a basis for a scientific approach to this problem of the double sulfate separation, which has heretofore been attacked by purely em pirical methods, and to enrich our knowledge of the important field of the double sulfates of the rare-earth elements, a study has been made of the preparation and properties of dou ble sulfates of several rare-earth elements taken from typical positions in the serial order. The combining sulfates studied included those of sodium, potassium, ammonium, and thallium. The thirty double sulfates prepared and reported upon in this research include twenty-four salts that have never been reported before, and their existence has been established 81 for the first time by this investigation. The methods of preparing these salts have been in vestigated, and certain of their properties, such as solu bility, and crystal system, have been closely studied. Of the few double sulfates that have up to this time been recog nized in the literature, most of them have been studied only by equilibrium diagrams, and they have not been actually pre pared and separated as entities* In the present investiga tion, each salt has been separated in a very near pure state, analysed, and in almost all cases, photographed and classified as to the crystal system. BIBLIOGRAPHY' BOOKS Browning, Philip Emburg, Introduction to the Rarer Elements. New York: John Wiley and Sons, 1919. Coraey, A.M., A Dictionary of Chemical Solubilities. 2nd edition, revised. New York: The Macmillan Company, 1921. Gooch, Frank Austin, Methods in Chemical Analysis. New York: John Wiley and Sons, 1912. Hopkins, B. Smith, Chemistry of the Rarer Elements. Me?/ York: D.C. Heath and Company, 1923. 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Asset Metadata

Creator Bickerdike, Ernest Lawrence (author)

Core Title A study of the double sulphates of some rare-earth elements with sodium, potassium, ammonium and thallium

Degree Doctor of Philosophy

Degree Program Chemistry

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag chemistry, inorganic,OAI-PMH Harvest

Language English

Contributor Digitized by ProQuest (provenance)

Advisor [illegible] (committee chair), [illegible] (committee member), Clements, Thomas (committee member), Roberts, L.D. (committee member)

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c17-17877

Unique identifier UC11348852

Identifier DP21722.pdf (filename),usctheses-c17-17877 (legacy record id)

Legacy Identifier DP21722.pdf

Dmrecord 17877

Document Type Dissertation

Rights Bickerdike, Ernest Lawrence

Type texts

Source University of Southern California (contributing entity), University of Southern California Dissertations and Theses (collection)

Access Conditions The author retains rights to his/her dissertation, thesis or other graduate work according to U.S. copyright law. Electronic access is being provided by the USC Libraries in agreement with the au...

Repository Name University of Southern California Digital Library

Repository Location USC Digital Library, University of Southern California, University Park Campus, Los Angeles, California 90089, USA