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COPRECIPITATIOS WITH BARIUM SULFATE M. L. NICHOLS AND E. C. SMITH Department of Chemistry, Cornell University, Ithaca, New York Receaved March 1 , 1940 INTRODUCTIOX
“There is probably no single operation in analytical chemistry that has received more attention from investigators than that of precipitation of barium sulfate by means of a soluble barium salt for the determination of sulfate or of barium. The literature on this subject is voluminous and the conclusions reached by different investigators as to the proper procedure to employ in order to obtain a precipitate which will be filterable and reasonably pure are highly contradictory.” (17). The reason for this situation is the fact that “Barium sulfate, to a greater degree than most precipitates, has the property of carrying down various soluble salts which may be present, and these cannot be removed by simply washing the precipitate.’’ (27). This was referred to by Foulk (4) in 1918 as “coprecipitation” and as such has been discussed thoroughly by Kolthoff (9). The case of barium sulfate is not singular, as coprecipitation affects all precipitation processes (11) to a greater or less degree. S o r is the concept a new one, as Turner (29) wrote of the impurities in barium sulfate precipitates over a hundred years ago. Since the beginning of the twentieth century various theories have been advanced to explain the contamination of barium sulfate precipitates. Some believe that the contaminating agent is present in solid solution,-a phenomenon also known as isomorphism or mixed-crystal formation. Iron is a common impurity in barium sulfate precipitates, and some workers (26) attributed the contamination to a solid solution of iron salts in barium sulfate. Grimm ( 5 ) found that barium sulfate formed mixed crystals with potassium permanganate and potassium fluoborate when precipitated in the presence of these salts, and quite recently Walden (31) has obtained evidence from x-ray diffraction studies that permanganate and nitrate ions are retained in the form of a solid solution. Schneider and Rieman (26) believe that this is true for nitrite. Occlusion of mother liquor by imperfect crystals or by crystal aggregates has been advanced to explain the presence of soluble material in this precipitate. Balarew (1) has written profusely regarding adsorption on the inner surfaces of precipitates. “When the particles of a precipitate grow together, it is very easy for a little liquid to become enclosed. This occlusion plays an especially important r81e when the particles dart together, as in the flocking out of colloids. With gradual crystal growth, the occlusion is not ordinarily noteworthy.” (11). However, Kulthoff and
412
M. L. NICHOLS h t i D E. C. SMITH
Bushey (10) say that barium sulfate may contain 3.5 per cent of water a t 100OC. which is lost a t 30OoC. This retained mother liquor may contain soluble salts. Double salts are sometimes formed by two salts having a common ion. The double salt may crystallize from solution before either single salt or sometimes a single salt is deposited and the composition of the solution changes to some critical value beyond which the double salt is thrown down (3). Many such compounds (15) have been postulated to explain the contamination of barium sulfate. However, since double-salt formation has been found to occur rather rarely in analytical precipitations, such explanations have but few adherents today. Adsorption (23) on the crystal surfaces has been proposed to explain the retention of all manner of substances by precipitated barium sulfate. Particles of colloidal dimensions with a large surface will exhibit surface adsorption to the highest degree. Adsorption may be positive, when the concentration of the solute is higher a t the interface; or negative, when it is lower a t the interface than in the body of the solution. The adsorption by barium sulfate is usually conceived to be polar rather than non-polar. Polar adsorption is of two kinds: first, the kind involving preferential adsorption by certain parts of the lattice for ions in solution of the opposite charge and the complementary layer of oppositely charged solution ions; and second, the polar adsorption in which the preferential adsorption of charged ions from solution results in the expulsion of an equivalent number of ions of the same charge from the crystal lattice. Kolthoff calls this “exchange adsorption”. The theories mentioned above have all been advanced to explain the presence of impurities in barium sulfate precipitates. Direct evidence on the problem is, of course, not easy to obtain. Most of the work which has been done has involved the precipitation of barium sulfate in a solution containing ions other than barium ion and sulfate ion. If foreign ions are present when barium sulfate is precipitated, the case for adsorption is not clear-cut, as under these conditions such other factors as occlusion, inner adsorption, and variation of the total surface must be considered. Therefore this study, which is similar to that of Weiser and Sherrick (23), was made under conditions which should eliminate these complicating factors. EXPERIMENTAL
All precipitations of barium sulfate were made in the special precipitation flask (figure l), with the volume of solution and weight of barium sulfate such as might obtain in analytical procedures. A volume of 0.3464 N sulfuric acid equal t o 0.004 mole was measured from a buret into the inside compartment and an equivalent amount of 0.2181 N barium hydroxide was placed in the outside compartment. Dis-
COPRECIPITATION WITH BARIUM SULF.4TE
413
tilled water was added to each to bring the volume to 100 ml. The top was moistened with water a t the ground-glass joint, placed in position, and the flask inverted by hand so that the two solutions were suddenly mixed. The inversion was performed once a second for 1 min. The air vent was opened, and a solution of 6 millimoles of the appropriate potassium salt was added from the separatory funnel. The funnel was rinsed with water so that the total volume in the flask became 250 ml. The air vent was closed, and the flask again shaken sixty times for 1 min. Then the cover was removed and the mixture filtered as rapidly as possible
L L- -----I
r----
FIG.1. Precipitation flask
through a No. 4 Jena sintered-glass filter. The precipitate was washed twelve times with small portions of hot water. In the case of chloride ion, this number of washings was sufficient to obviate any test with silver nitrate in the filtrate and the same number of washings was maintained for all other ions. Filtration and xashing required 20 t o 30 min. The precipitate was transferred with a spatula to a porcelain crucible and dried a t 110OC. for about 23 hr. Between 0.56 and 0.8 g. of this dried barium sulfate was then analyzed for the adsorbed impurity. Before determining the amount of contamination by each salt, the methods of analysis mere checked with some of the pure potassium salt.
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M. L. NICHOLS AND E. C. SMITH
The method used by Walden and Cohen (31) and by Schneider and Rieman (26) for the determination of nitrate in the barium sulfate has two objections. The barium sulfate is not completely soluble in the sodium carbonate solution (35), and the evolution of much hydrogen during the ammonia distillation is undesirable with weak acid. Weiser and Sherrick employed the Foerster modification of the Kjeldahl method but had exceedingly large blanks. This was corrected by substituting salicylic acid for phenol and using a solution of sodium thiosulfate instead of the solid. The “reagent grade” potassium nitrite was purified by recrystallization from alcohol, and analysis by the method recommended by Kolthoff and Sandell (12) showed that it was about 98.8 per cent pure. In order to minimize oxidation in solution, the solid potassium nitrite was weighed out in an amount (0.5166 g.) to give 6 millimoles and dissolved just before addition to the barium sulfate. The nitrite was determined by treating the barium sulfate with 25 ml. of sulfuric acid containing 1.5 g. of phenol, and 2 g. of solid sodium thiosulfate was added very slowly to the cold solution. After reduction was complete, the mixture was warmed for the first time, and digestion completed a t a high temperature after the addition of a small amount of copper sulfate and 5 ml. of additional sulfuric acid. The ammonia was then steam distilled and determined by the usual procedure. “Reagent grade” potassium cyanide, labelled 95 per cent, was analyzed by the procedure outlined by Kolthoff and Sandell (13) and gave a result of 94.7 per cent. Six millimoles of this analyzed material was dissolved just before adding it to the barium sulfate. The barium sulfate was dissolved in 10 ml. of sulfuric acid and heated until sulfur trioxide was copiously evolved to convert the cyanide to ammonium ion (16). After cooling, the ammonia was steam distilled. The methods used for determining thiocyanate, ferrocyanide, and ferricyanide were similar to that used for potassium cyanide. The method used by Weiser and Sherrick (23) for the determination of chloride was tried, but satisfactory results could not be obtained. The method finally adopted was similar to that of Robertson (24) and Thompson-Oakdale (34) for organic halogen compounds. The oxidizing mixture consisted of 5 ml. of a suspension of 12 g. of silver dichromate (prepared by the method of Autenrieth (21)) and 12 g. of potassium dichromate in 480 ml. of sulfuric acid. The aspiration apparatus was the same as used by Weiser and Sherrick, with the ground-glass joint moistened with sulfuric acid. It was necessary to add 5 ml. of 0.01 N sodium hydroxide to the first potassium iodide absorption tube, apparently to eliminate oxidation to iodine by air. The aspiration was continued for 30 min. a t 100°c. after complete solution.
415
COPRECIPITATION WITH BARIUM SULFATE
The analysis for bromide was similar to that for chloride, except that more alkali was added to the first potassium iodide absorption tube and aspiration was continued for 1 hr. a t 17OoC. after complete solution. TABLE 1 Tests of analytical methods POTASSIUM &UsT
WEIQET O P SAMPLE OF POTA00IUM SALT
VOLUME OF ,OLDTION FOR TITRATION
AVERAQE
BLANK
WEIQET OP POTASSIUM SALT FOUND
RECOVERY
m9.
ml.
ml.
ms.
per cenl
Ferricyanide . . . . .
0.507 0.519
8.93 9.06
0.25
0.476 0.483
93.9 93.1
Ferrocyanide.. . . .
0.325 0.334
5.47 5.56
0.31
0.317 0.322
97.6 96.4
Iodide . . . . . . . . . . .
0.265 0.283
8.21 8.72
0.0
0.244 0.259
92.1 91.5
Thi*ocyanate... . .
0.733 0.742
7.65 7.60
0.19
0.725 0.720
98.9 97.0
Cyanide. .
.. , . ...
0.487 0.474
7.37 7.08
0.12
0.472 0.453
96.9 95.6
Bromide
..
0.815 0.805
6.32 6.17
0.0
0.808 0.789
99.1 98.0
Permanganate.. .
0.911 0.908
6.45 6.42
0.0
0.907 0.903
99.6 99.5
Chloride . . . . . . . .
0.433 0.439
0.0
0.436 0.432
100.7 98.4
Chlorate. . . . . . . .
0.685 0.693
7.30 7.42
0.664 0.675
96.9 97.4
Sitrite . , . . . . . . . .
0.711 0.719
9.36 8.77
0.84
0.726 0.675
102.1 93.9
0.866 0.873
6.39 6.37
0.55
0.844 0.838
97.5 95.9
,
. ...
Kitrate. . . . . , , , . , ,
1 ~
i
I
The iodide was determined by treating with sulfuric acid and aspirating the iodine for 1 hr. a t 170°C. into 5 per cent sodium hydroxide. Following absorption, the solution was rinsed into a glass-stoppered Erlenmeyer flask and 15 ml. of a mixture (22) containing 30 g. of sodium acetate tri’nydrate. 100 ml. of glacial acetic acid, 90 drops of bromine, and 200 m!.
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M. L. NICHOLS AXD E. C. SMITH
of water was added. The excess bromine was destroyed with 5 drops of formic acid, and then 2 ml. of 10 per cent potassium iodide and 5 ml. of (1:lO) sulfuric acid were added; after 5 min. the liberated iodine was titrated. The solid potassium iodide, when analyzed by this method, gave an average result of 96.1 per cent of the theoretical. The permanganate and chlorate were determined by treating the sample, plus 0.2 g. of barium chloride, with sulfuric acid containing a few drops of hydrochloric acid and aspirating with nitrogen into alkaline potassium iodide for 30 niin. a t 100°C. The results for the analytical methods and the determination of the anions adsorbed by barium sulfate are given in tables 1 and 2. The ions are arranged in the order of adsorption, beginning with the least adsorbed. TABLE 2 Adsorption of anions b y barium sulfate PERCENTAQE CONTAMINATION BY POTASSICTM BAL"
ION
Ferricyanide. . . . . . . . . . . . . . . . . . . . . . . Ferrocyanide. ..................... Iodide, . . . . . . . . . . . . . . . . . . . . . . . . . . Cyanide. . . . . . . . . . . . . . . . . . . . . . . . . . . Bromide. . . . . . . . . . . . . . . . . . . . . . . . . . . Permanganate. . . . . . . . . . . . . . . . . . . . . Kitrite............................ Nitrate.. ..........................
0.019 0.041 0.281 0.308 0.208 0.420 0.677 0.462 0.918 0.950 1.61
ANION PER 100 MOLE0 OF BESO'
grama
gram .quivalen ta
2.87 5.56 50.3 43.1 19.4 66.1 120 54.8 147 121 234
0.041 0.105 0.396 0.742 0.747 0.827 1.01 1.55 1.76 2.63 3.78
EFFECT O F .IGING
Four millimoles of barium sulfate was precipitated in the usual manner, washed into a 400-ml. beaker, and digested on a hot plate below the boiling point for 10 hr. The suspension was cooled to room temperature, transferred back to the precipitation vessel, and diluted to the original volume. Then 0.006 mole of potassium chlorate was added in 50 ml. of Eater, the mixture shaken for a minute, and the precipitate filtered, washed, and dried for 33 hr. a t 110°C. The filterability was not improved, but the analysis of three samples indicated that no potassium chlorate had been adsorbed. DISCUSSION
There are three methods of investigating the impurities carried down by barium sulfate precipitates. The easiest is to weigh the precipitates
COPRECIPITATION WITH BARIUM SULFATE
417
and compare the weights with controls in which the soluble material is absent. It is simple and useful for finding the magnitude of the error in analytical procedures. It suffers from the defect that it is indirect and what the impurity really is cannot be ascertained with certainty, if more than a single foreign substance is present in the solution from which precipitation is made, Furthermore, a compensation of errors may completely mask serious contamination. The second method is to analyze the solution for soluble material, after removing the precipitate, and to attribute any decrease in concentration of a salt to removal by the precipitate. It is possible for both water and a soluble salt to be removed by the precipitate in such a ratio that the concentration of the salt in solution is unchanged. Also, since the amount of salt remaining in solution is large, a small relative error in determining its amount may be a serious error in estimating the quantity removed by the precipitate. The last and most difficult method, but the one which will give the most accurate results, is to analyze the precipitate. “The formation of precipitates from molecularly or ionically dispersed solutions takes place in two stages: (a) the formation of crystal nuclei and (b) the growth and agglomeration of these particles to sizes sufficiently large to settle out under the action of gravity. I n the formation of any precipitate there is therefore a stage a t which the particles are of colloidal dimensions.” (2). Most substances, even though of colloidal dimensions, have a crystalline structure. I t is not possible t o follow the changes from complete dispersion to a particle size visible under the microscope, but it is generally believed that the growth is crystalline below the range of microscopic visibility. There are several factors which may delay or prevent the change into particles which settle by the action of gravity. One of these, according to von Weimarn (30), is the velocity of nuclei formation, V’,which equals K ( Q - L)/L. The growth of initial particles is represented by the Xernst-Soyes equation (32) that
D V = - * O *(Q - L) 9 That there are other factors which also affect the process cannot be denied. Iiolthoff and Sandell (14) agree nith Haber (6) that one factor which cannot be neglected is the orientation velocity in the growing crystal. However, a strongly polar substance like barium sulfate has a high orientation velocity and this is not the rate-determining step. An important consideration affecting Ton Keiniarn’s conclusions regarding particle size is the ratio S,/S. Hulett ( 7 ) shoned that fine wbdivision of barium sulfate increased the solithility, and Kolthoff (9) has calculated
418
M. L. NICHOLS AND E. C. SMITH
this ratio for various substances. This increased solubility of small crystals is the explanation usually given for the beneficial effects of digestion. It appears possible, however, that in most cases before digestion is begun, the particles are already larger than they are in the zone where increased solubility plays a &le. Lindsly (17), Trimble (28), and O d h (19) have contended that digestion improves filterability because it induces dendritic aggregation and not because the particles grow larger. This latter seems doubtful, however, when we consider the effect of digestion upon adsorption. With potassium chlorate and contact immediately after precipitation, the barium sulfate carried down 1.76 equivalents of chlorate ion per 100 moles, while after 10 hr. digestion no detectable amount of chlorate ion was found to be adsorbed. Substantially the same results were found by Hulett and Duschak (15). Particles in the colloidal zone are frequently stabilized by acquiring an electrostatic charge, and a precipitate shows a preferential adsorption for its own ions. Barium sulfate, as usually obtained, is positively charged by the adsorption of barium ions, although Kichols and Thies (18) have prepared the negatively charged colloid. Barium sulfate formed with sulfate in excess is easier to filter than when barium is in excess, especially if some hydrochloric acid is present. However, since sulfate is more commonly determined than barium, the excess of barium is unfortunately the more common case. Yon-polar substances may be expected to follow the Gibbs’ rule in regard to adsorption, but polar substances do not necessarily follow this behavior. The adsorption of cations by barium sulfate seems to follow the order of sulfate solubilities (20). According to the data &en by Weiser (33) for the molar solubilities, the barium salts of the various anions show the order Fe(CN)b
< NOT < Cloy < C1- < MnOl < NO; < Br- < CN- < I- < SCN-
Barium ferricyanide is listed “very soluble.” found in this study is
NO;
The order of adsorption
> NO; > ClOT > C1- > MnO; > Br- > CN- > SCN- > I- > Fe(CN)b > Fe(CN)?
Except for KO*-, SCN-, and Fe(CN)6-4, the adsorption 01anions supports the rule that decreased solubility of a compound of the anion with the positive ion of the precipitate causes increased adsorption of that anion. The Schulze-Hardy rule does not appear to be verified, or eke high adsorption of ferrocyanide and ferricyanide should have been obtained. From the authors’ experience, the high contamination of these anions
COPRECIPITATION WITH BARIUM SULFATE
419
found by Weiser and Sherrick (23) may have been due to local excess of barium causing actual precipitation of these barium salts. The order of adsorption obtained by Weiser and Sherrick was Fe(CN)b
> NO; > NO, > ClO; > MnO; > Fe(CN)pa > C1- > Br- > CN- > SCN- > I-
which agrees with the order found if chloride ion and permanganate ion are reversed and if ferrocyanide and ferricyanide are omitted. The data obtained support the idea that adsorption plays an important part in the coprecipitation of impurities. The fact that finely divided barium sulfate takes up potassium salts of various anions means that similar adsorption can occur on the primary particles during analytical precipitations. As ordinarily precipitated in an acid solution, barium sulfate is peculiarly suitable for anion adsorption. The short contact period minimizes the solid solution theory, since time is necessary for fusion into the crystal interior. The order KCl> KBr > K I generally found has been verified by these data. If the polarizability of the anion tends to increase adsorption, this order should be reversed. Potassium salts are generally adsorbed more than the corresponding sodium and ammonium salts (8). Using the same technique as here employed, one of the authors in 1930 obtained the following values of gram-equivalents per 100 moles of barium sulfate: 1.62 for chloride ion from potassium chloride, 0.87 for chloride ion from sodium chloride, and 0.85 for chloride ion from ammonium chloride. The adsorption values which Weiser and Sherrick obtained for the strongly adsorbed nitrate, nitrite, chlorate and permanganate ions are much higher than our values, although a lower molar concentration was used. This might be expected, since the contaminating salt was present when the barium sulfate was precipitated and because barium salts are probably more strongly adsorbed than potassium salts. However, our values for cyanide ion, thiocyanate ion, and iodide ion are higher and the values for chloride ion and bromide ion are practically the same. The data obtained give strong evidence that adsorption on the surface of the primary particles of barium sulfate is a major factor in the coprecipitation of impurities. Aggregation of these primary particles with their adsorbed impurities will cause a random distribution of the impurity throughout the mass of the precipitate. Future x-ray investigation will substantiate or destroy the claims that certain anions are retained in solid solution. SUMMARY
1. Barium sulfate has been precipitated in a finely divided form by mixing stoichiometric quantities of sulfuric acid and barium hydroxide.
420
M. L. NICHOLS AND E. C. SMITH
The quantities of eleven potassium salts adsorbed by the finely divided barium sulfate upon short contact have been studied. 2. Methods for determining the anions in the contaminated barium sulfate have been tested with known quantities of material. 3. The gram-equivalents of anions taken up by the barium sulfate from potassium salts show the following decreasing order:
SO;, NO;, ClO;, C1-, MnO;, Br-, CX-, SCN-, I-, Fe(CX)b, Fe(CN)i3 4. That digestion is effective in reducing contamination has been demonstrated. 5. Strong evidence has been found that adsorption on the surface of the primary particles of barium sulfate is a major factor in the coprecipitation of impurities. REFERENCES (1) BALAREW: Z. anal. Chem. 116, 104 (1938). (2) BURTON:T h e Physical Properties of Colloidal Solutions, p. 37. Longmans, Green and Company, London (1938). AND MASON:Handbook of Chemical Microscopy, Vol. I, p. 360. John (3) CHAMOT Wiley and Sons, Inc., New York (1930). Notes on Quantitative Chemical Analysis, p. 44. McGraw-Hill Book (4) FOULK: Company, Inc., New York (1918). See BROWNING: Am. J. Sci. 46, 399 (1893). (5) GRIMM:Z. Elektrochem. 30, 467 (1924). HUTTIGA N D MENZEL:2. anal. Chem. 68, 343 (1926). (6) HABER:Ber. 66B, 1717 (1922). (7) HULETT:2.physik. Chem. 37,385 (1901). (8) KATOA N D NODA:J. Chem. SOC.98, 11, 895 (1910). (9) KOLTHOFF: J. Phys. Chem. 36, 860 (1932). (10) KOLTHOFF AND BUSHEY:Bull. soc. chim. Belg. 47, 689 (1938). (11) KOLTHOFFA N D FURMAN: V o l u m e t k Analysis, Vol. 1, p. 152. John Wiley and Sons, Inc., New York (1928). (12) KOLTHOFF A N D SANDELL: Quantitative Analysis, p. 574. The MacMillan Company, New York (1937). (13) Refermce 12, p. 545. (11) Reference 12, p. 94. (15) KUSTERA X D THIEL: 2 . anorg. Chem. 22, 424 (1900). JANNASCH A X D RICHARDS: J . prakt. Chem. 39, 321 (1889). RICHARDS:Z. anorg. Chem. 23, 383 (1900). SILBERBERGER: Monatsh. 26,220 (1904). HULETTA N D DUSCHAK: Z. anorg. Chem. 40, 196 (1904). FOLIX: J. Biol. Chem. 1, 85 (1905). SMITH:J . Am. Chem. SOC.39, 1152 (1917). KARAOGLAXOW:z. anal. Chem. 67, 77 (1918). KOLTHOFF ASD VOGELESZASG: Pharm. Weekblad 66, 122 (1919). BALAREW: 2. anorg. allgem. Chem. 123, 69 (1922). (16) LATIMERA X D HI~.DEBRASD : Reference Book of I n o ~ g a n i cChemistry, p. 227. The 3IaclIillsn Company, Xew York (1931).
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LINDSLY: Ind. Eng. Chem., Anal. Ed. 8, 176 (1936). NICHOLS AND THIES:J. Ani. Chem. SOC. 48,302 (1926). O D ~ J.: Chem. SOC.120, 25 (1921). PAXETH AND HORORITZ: Z. physik. Chem. 89, 513 (1915). HAHK:Saturwisseiischaften 14, 1196 (1926); Ber. 69, 2014 (1926). (21) PREGL(translated by R o t h ) : Qctantitaliue Organic MicroanaZ?jsis, 3rd English edition, p. 110. P. Blakiston’s Son and Co., Philadelphia (1937).
(17) (18) (19) (20)
(22) Reference 21, p. 113. 2. anorg. Chem. 23, 3S3 (1900). (23) RICHARDS: KORTE:J. Chem. SOC.81, 1503 (1905). JOHSSTONA X D ADANS: J . .Am. Chem. SOC.33, 829 (1911). WEISERAND SHERRICK: J. Phys. Chem. 23, 205 (1919). GERMANX: J . Ani. Chem. Soc. 43, 1615 (1921). DOERSER AND HOPKINL: J . Am. Chem. SOC.47, 662 (1925). PANETH AND HOROWITZ: Z. physik. Chem. 89, 513 (1915). EBLERAND V A N RHTS: Ber. 64, 2896 (1921). KAhI3IER A N D SILBERlI.4 J. .Im. Chem. So?. 47, 2514 (1925). BALCAR AND STEGEMAN: J. I’hys. Chem. 32, 1411 (1928). HAHN:Ber. 69, 2014 (1926); 2. physik. Chem. 144, 161 (1929). C H I T T E R J I : z. anOrg. allgem. Chem. 121, 128 (1922). DUTOIT AND GROBET:J. chim. phys. 19, 32s (1921). GHOSHASD DHAR:Kolloid-Z. 36, 141 (1924). CH.4KRAWARTI AXD DHAR: KOllOid-Z. 44, 63 (1928). DE BROUCKERE: Bull. acad. sci. roy. Belg. 13, 415 (1927); 16, 1263 (1930). (24) ROBERTSON: J. Chem. SOC.107, 902 (1915). 2 . physik. Chem. 10, 425 (1893). (25) SCHNEIDER: RICHARDS: Z. anorg. Chem. 23, 3S3 (1900). AND RIEMAS:J. Am. Chem. soc. 69, 351 (1937). (26) SCHSEIDER (27; S m m : Quantilatiue Chemical Analysis, p. 64. The >Iachlillan Company, New York (1928). (28) TRIIIIBLE: J. Phys. Chem. 31, 606 (1927). (29) TURNER: Phil. Trans. 119, 295 (1829). (30) vos WEIMARN:Chcm. Rev. 2, 217 (1926). (31) WALDENA N D COHES: J. Ani. Chem. SOC.67, 2591 (1935). AVERELL A K D WALDEX:J. .Im. Chem. SOC.69, 906 (1937). G R I m r , PETERS, ~ S ~ DY O L F F :2. anorg. ahgem. Chem. 236, 57 (1938). (32) WEISER: The Colloidal Salts, p. 12. J. Wiley and Sons, Inc., New T o r k (1927). (33) Reference 32, p. 37. J. Am. Chem. SOC.62, 1893 (1930). (34) WILLARDA X D THOMPSON: (35) WOLESKY:Ind. Eng. Chem., -inaI. Ed. 1, 28 (1929).
