306
ANALYTICAL CHEMISTRY
a melting point of 199’ C., and the recrystallization failed to raise this value significantly. The ether-insoluble residue was identified as potassium sulfate by means of qualitative tests and ite optical-cr stallographic properties. Table I%gives the results. POLAROGRAPHICANALYSIS
A polarographic analysis was run on the 2,4-dinitrophenylhydrazones of each of the following aldehydes: acetaldehyde, propanal, 2-methylpropanal, butanal, 2,2,3-trimethylbutanal, 4-methy1pentanalJ n-hexanal, and furfural. The derivatives were prepared by Brady’s method. The modification of the acetaldehyde derivative used melted a t 168’ C. Both forms of furfural 2,4-dinitrophenylhydrazonewere used. The acetone-water buffer of Cray and Westrip ( 1 7 ) with a p H of 9.45 was used as a solvent. All measurements were made on an Aminco polarometric analyzer No. 5-2500. The curves were analogous to those reproduced for the furfural derivatives (Figure 5). I n Table V, El’* designates the half-wave potential and id the diffusion current in microamperes. Concentration is given in moles per liter. The “total wave” was measured from the residual current to the upper limiting current. LITERATURE CITED (1) Allen, C. F. H., and Richmond, J. H., J. Org. Chem., 2, 222 (1937). (2) Bell, F. G., Biochem. J., 35,312 (1941). (3) Braddock, L. I., and Willard, M. L., J . Am. Chem. Soc., 73,5866 (1951). (4) Brady, 0. L., J.Chem. Soc., 1931,756. (5) Brady, 0. L., and Elsmie, G. V., Analyst, 51,77 (1926). (6) BrandstLtter, M., Mikrochemie ver. Mikrochim. Acta, 32, 33 (1944). (7) Braude, E. A., and Jones, E. R. H., J . Chem. Soc., 1945,498. (8) Brederick, H., Ber., 65, 1833 (1932). (9) Brederick, H., and Fritsche, E., Ibid., 70,802 (1937). (10) Brunner, F., and Farmer, H. F.. J . Chem. Soc., 1937,1039. (11) Bryant, W. M. D., J . Am. Chem. Soc., 54, 3758 (1932); 55, 3201 (1933); 58,2335 (1936). (12) Ibid.. 60.2814 11938). (13j Chattaway, F. D., and Clemo, G. R., J . Chem. Soc., 123-2T, 3041 (1923). 1
,
(14) Clark, G. L., Kaye, W. I., and Parks, T. D., IND. ENG.CHEM., ANAL.ED.,18,310 (1946). (15) Conard, C. R., and Dolliver, hl. A, “Organic Syntheses 11,” p. 167, Kew York, John Wiley & Sons, 1943. (16) Cowley, M. A,, and Schuette, H. A., J . Am. Ckem. Soc., 5 5 , 3463 (1933). (17) Cray, L. I., and Westrip, J., Trans. F U T ~Soc., U ~21, 326 (1925). (18) Dirscherl, W., and Nahm, H., Ber., 73,448 (1940). (19) Djerassi, C . P . , J . Am. Chem. Soc., 71, 1003 (1949). (20) Djerassi, C. D., and Ryan, E., Ibid., 71,1000 (1949). (21) Fernandea, O., and Costello, M., Anales soc. espafi. fis. y qulm., 33,81 (1933). (22) Ferrante, J., and Bloom, A , Am. J . Pharm., 105, 381 (1933). (23) . . Gordon. B. E.. W-oDat. F.. Burnham. H. D.. and Jones, L. C., ANAL.CHEM., 23,-1754 (1951). (24) Granick, S., Science, 80,272 (1934). (25) Ingold, C. K., Pritchard, G. J., and Smith, H. G., J . Chem. Soc., 1934,79. (26) Lange, J. J. de, and Houtman, J. P. W., Rec. trav. chim., 65, 891 (1946). (27) Linnell, W. H., and Roushdi, J., Quart. J . Pharm. Pharmacol., 12.252 ,~ (1939). (28) Lockhart, E. E., Merritt, M. C., and Mead, C. D., J . Am. Chem. Soc., 73,858 (1951). (29) Lynn, G., J . Phgs. Chem.,31, 1381 (1927). (30) Malkin, T., and Tranter, T. C., J . Chem. Soc., 1951, 1178. (31) Matthiesson. G.. and Hapedorn. H., &fikrochemiever. Mikrochim. Acta, 2 9 , 5 5 (1941); (32) Mitchell, J., ANAL.CHEM., 21,448 (1949). (33) Morton, A. A , , “Laboratory Technique in Organic Chemistry,” p. 32, Kew York, McGraw-Hi11Book Co., 1938. (34) Roberts, J. D., and Green, C., J . Am. Chem. Soc., 68,214 (1946); IND. ENG.CHEM., ANaL. ED., 18, 335 (1946). (35) Shriner, R. L., and Fuson, R. C., ”Systematic Identification of Organic Compounds,” 3rd ed., p. 171, New York, John Wiley & Sons, 1948. (36) Simon, E., Biochem. Z . , 247, 171 (1932). (37) Strain, H. H., J . Am. Chem. Soc., 57,758 (1935). (38) Samant, H. H., and Planinsek, H. L., Ibid., 72, 4042 (1950). ~~~
I
RECEIVED for review December 5 , 1981. Accepted October 9, 1952. Presented in part before the Division of Analytical and Micro Chemistry at t h e 116th Meeting of the AMERICAN CHEMICAL SOCIETY, Atlantic City, N. J., and the X I I t h International Congress of Pure and Applied Chemistry, Section 2, Analytical Chemistry, New York, N. Y., September 9 t o 13, 1981.
X-Ray Investigation of Several Contaminated Barium Sulfate Precipitates C. A. STREULI, H. A. SCHERAGA, ANDM. L. NICHOLS Cornell University, Ithaca, N. Y.
I
N RECENT years much progress has been made in understanding the mechanism by which coprecipitation occurs during the formation of barium sulfate precipitates. The adsorptive properties of freshly precipitated barium sulfate for foreign ions have been investigated by Weiser and Sherrick ( I S ) and by Nichols and Smith (8). Cohen’s method (6, 1 1 ) for the determination of precise lattice constants from x-ray diffraction patterns has enabled Walden and coworkers (9]1 I , I 2 ) to demonstrate the importance of solid solution in accounting for the contamination of barium sulfate precipitates by nitrate, permanganate, water, and several univalent cations. The x-ray diffraction technique has been applied here to investigate the nature of contamination of barium sulfate precipitates by ferric, chloride, and nitrite ions, all of which have been reported ( I O ) to be strongly coprecipitated during the formation of barium sulfate. PREPARATION OF PRECIPITATES
Barium sulfate precipitates were prepared in a manner which excluded all contamination except water and the contaminating
ion under study. Reagent grade chemicals were used throughout. “Pure” barium sulfate was precipitated b adding 50 ml. of approximately 0.2 M carbonate-free barium lydroxide dropwise to 200 ml. of a boiling, mechanically stirred, approximately 0.1 M sulfuric acid solution. The preci itates were digested in the hot solution for 4 hours and washed y! decantation with 2 liters of hot water. They were then filtered through sintered-glass funnels, rewashed, and oven-dried for 2 hours a t 130” C. This procedure was also employed to prepare all the contaminated precipitates except for the changes noted below. Ferric ion contamination was introduced by adding ferric sulfate to the sulfuric acid before precipitation; the sulfate ion concentration was maintained constant by dilution with sulfuric acid while the ferric ion concentration was varied from 0.025 to 1.0 M . These precipitates were washed with dilute acid to prevent subsequent precipitation of ferric hydroxide on the surface of the barium sulfate. The pH of the mixed solutions mas less than 2. Chloride-contaminated precipitates were prepared by adding hydrochloric acid in varying amounts to each of the precipitating solutions. Chloride content, 0.5 to 3.0 M , was adjusted so as to be equal in both solutions before mixing. The instability of nitrous acid made it necessary to premix
V O L U M E 25, NO. 2, F E B R U A R Y 1 9 5 3
Ferric, nitrite, and chloride ions have been reported to be strongly coprecipitated with barium sulfate. This contamination could arise in several ways, such as by adsorption, solid solution, or compound formation. A study w-as, therefore, carried out to determine the nature of the coprecipitation process in these systems, making use of the method developed by Cohen for the determination of precise lattice constants from x-ray powder diffraction patterns, applied earlier by Walden and coworkers in similar investigations. The barium sulfate unit cell volume has been observed to vary with the per
the barium hydroxide solution with sodium nitrite and add this solution to the sulfuric acid. The nitrite concentration ranged from 0.1 to 2 M . Precipitates were formed in cold solutions and digested only 2 hours. Some decomposition occurred, however, as shown by the presence of nitrogen dioxide. 'These precipitates were oven-dried for 1 hour a t 110' C. DETERMINATION O F CONTAMINANTS
The precipitates were analyzed for the contaminants added, and also subjected to prolonged heating a t 700" C. t o determine the amount of moisture present after oven-drying. The ferric-contaminated precipitates were analyzed by the o-phenanthroline method described by Sandell (9). The samples were dissolved in hot, concentrated sulfuric acid and poured into distilled water to reprecipitate the barium sulfate. The filtrate was analyzed for iron at pH 3.5 with a Beckman Model DU spectrophotometer a t 510 mp. The method was checked by adding known amounts of ferric ion to "pure" barium sulfate prior to the dissolution of the precipitate in concentrated sulfuric acid.
307
cent contamination in the case of all the ions considered. Because no change was observed in the diffraction pattern, these variations in unit cell volume have been attributed to solid solution. Adsorption has also been shown to play a role in the case of contamination by ferric ion. Solid solution can occur in systems such as those considered here and also those studied by Walden and coworkers, where the degree of contamination is relatively small, even though Grimm's rules for solid solution are not satisfied. Further evidence is thus provided of the importance of this phenomenon in coprecipitation.
Sitrites were deterininctl by dissolving the samples in approxiiiiately 0.1 .If tctrasotiium \'ersenate (eth~~leiirdianiine tetraacetate), reducing the nitrite t ( J ammoiii:t with Drvnrdn's alloy and strong alkali, ant1 distilliiig the :ininionin I'ormrd into an excess of standard acid, which \vas bncsk-tiirattd ivith standard alkali (7). Satisfactory results could not be obtnineJ it' thc preci itates were dissolved in a sodium carbonatc-sodium hydrosiie misture. The instability of the nitrite prevented the use of coiicPiitrated sulfuric acid as the solvent; the Versen:itts d u t i o n , however, proved to be an excellent solvent. Knoivii samples of sodium nitrite containing 20 ing. of nitrite gave theoretical recovery within 2% when detcrrnined by thic mothod. The precision in the analysea of the unknowns wad also within 2yC,
0.4 0.5
3
1
The results indicated that for the concentrations of iron encountered in contaminated precipitates no iron is coprecipitated when the sulfuric acid solution of barium sulfate is reprecipitated in water. The precision and accuracy of the iron analysis were within 5yo. 0.01
0.05 0.1 MOLARITY FE I N SOLN
0.5
I.o
Figure 2. Relation of Iron in Precipitates to Molar Concentration of Iron in Precipitating Solution
0.4 0.8 MOLARITY FE IN PRECIPITATING SOLUTION
I.e
Figure 1. Relation of Iron in Precipitates to Molar Concentration of Iron in Precipitating Solution
The chloride determination was performed by dissolving the precipitates in hot, concentrated sulfuric acid in a closed system, and sweeping the liberated hydrogen chloride into dilute silver nitrate solution with air (IS). After acidification with 1 to 1 nitric acid, the precipitate was flocculated and filtered off, and the filtrate was back-titrated with ammonium thiocyanate solution using ferric alum as indicator. Known samples of pure sodium chloride and barium chloride gave reproducible and accurate results within the limits of the determination-Le., 1 to 2% for a 6-mg. sample. The precision in the analyses of the unknowns was also within 1 to 2%.
All the precipitates were heated for 3 hours a t about 700" C. after oven-drying t o determine loss upon heating. The ferriccontaminated precipitates evolved sulfur trioxide as well as water upon heating and changed from the previous white t o the redbrown color of ferric oxide. The chloride- and nitrite-contaminated precipitates after heating showed complete expulsion of those ions upon reanalysis, within the limits of detectability of the method. The weight per cent was then taken as the total per cent volatility less the weight per cent loss of sulfur trioxide, hydrogeq chloride, or nitrogen dioxide, respectively. X-RAY PROCEDURE
Photographs of the precipitates were obtained using a symmetrical back-reflection, focusing camera of 10.007 & 0.003 cm. diameter and Norelco x-ray diffraction unit. The vacuum x-ray tube contained a copper target; the radiation was filtered through a nickel filter to reduce the intensity of short wave length radiation and to eliminate the copper Kp radiation. If a nickel filter is used with copper radiation, adequate photographs of barium sulfate may be obtained, and i t is unnecessary to resort to the longer wave length radiation of calcium used by Walden and coworkers (2, 6, 11, I d ) . Samples were prepared by grinding the oven-dried precipitates in an agate mortar, passing through a 200-mesh screen, and then dusting onto a thin film of Lubriseal spread on a celluloid sheet. The celluloid mount was then
ANALYTICAL CHEMISTRY
308 clamped into the sample holder on the circumference of the camera.
h
Calculations were carried out by the method of Cohen (6) for determining precise lattice constants. The majority of the results are averages of duplicate x-ray determinations of each precipitate. All parameter values were corrected to 25" C. and are expressed in angstrom units (4). The precision for all lattice constants was 0.01% or better.
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RESULTS
The degree of contamination of the precipitates by the foreign ions studied is low, being of the order of 0.01 mole % in all caaes. Nitrites furnished the highest contamination of the three systems studied, ferric and chloride ion contaminations being approximately equal. The increase in chloride contamination with decreasing chloride content of the precipitating solution may be related to an increase in pH of the solutions with the lowering of the hydrochloric acid content. Values in mole per cent are given in Table I.
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a 12 16 MOLE t GI x 103 Figure 4. Relation of Lattice Constants and Unit Cell Volumes to Mole Per Cent Contamination
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per cent contamination are given in Figures 3, 4, and 5 . In general the unit cell volume and some of the lattice parameter8 change regularly with increased contamination. The values for the parameters of the "pure" precipitates have been corrected for the water content of the precipitates, as Walton and Kalden have demonstrated the solid solution properties of water with
is
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Table 1. Determination of Contaminating Ions and Water in Barium Sulfate Precipitates
4
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MOLE yo FE x
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Figure 3. Relation of Lattice Constants and Unit Cell Volumes to Mole Per Cent Contamination Contamination bv w t e r is much higher than that due to foreign ions. Water contaniination ranges from 5.8 to 30.9 mole yo (Table I). The ferric-contaminated precipitates show the greatrst water content, the chloride-contaminated ones the least. The nitrite-contaminated precipitates show decreasing water content with increasing nitrite content. A plot of weight per cent iron in the precipitates against molar concentration of iron in the precipitating solution is shown in Figures 1 and 2. The curves suggest an adsorption phenomenon described by a Freundlich isotherm, P = ay"; similar b e havior was reported by d'Alcontres ( I ) for the chloride-contaminated system Similar plots for chloride- and nitrite-contaminated systems could not be obtained from the data, although there are indications that one might be obtained for the nitrite system if decomposition of that ion could be prevented. Plots of lattice constants and unit cell volumes against mole
System and Sarriple No. "Pure" 3 4 Ferric-contaminated F-1 F-2 F-3 F-4
_F-6_
F-.5
F-7 F-8 F-9 F-10 F-11 Chloride-contaminated CI-1 c1-2 C1-R -.. (21-4 Cl-5 C1-6 C1-7 C1-8 Nitrite-contaminated S-1 5-2 N-3 N-4 N-5 N-6 N-7 N-8 hT-9 li-10
Molarity Contmnt. in Pptng. Soh. 0.0 0.0
Mole % Contmnt. X 103
... ...
Mole 5; Water 17.3 16.0
1.0 1.0 0.5 0 5 0 25 0,20 0.10 0.10 0.0: 0.05 0,023
13.4 14.3 11.6 11.6 10.0 9.1 6.0 6.6 5.2 4.8 3.8
3.0 3 0 2.0 2.0
9.8 9.8 5.9 7.8 5.9 7.8 13.0 lZ.6
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22.5 14.1 13.6 16.6 15.0 15.0 12.5 12.9 3.9 6.0
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309
V O L U M E 25, NO. 2, F E B R U A R Y 1 9 5 3 barium sulfate precipitates and consequent lattice expansion in the aa and co directions (12). DISCUSSION
Pure barium sulfate has orthorhombic symmetry and space group VAG,and contains four barium sulfate molecules per unit cell. The positions of the atoms p-ithin the cell are given by both Basche and 3,lark ( 3 ) and IVyckoff (14). iiccording t o Wyckoff, barium coordination is not clear-cut and the rrlation of the ions to one another does not closely resemble that of simpler crystals. The sulfur coordination is, of course, four with a sulfur to oxygen bond length of 1.65 A. In light of the known structure of barium sulfate and the inapplicability of Grimm’s rules, as demonstrated by Walden and Cohen (11) in cases where the amount of contamination is low, the possibility of the introduction of contaminants by solid solution is not incompatible with the barium sulfate structure. Solid solution will be evidenced by distortion of the lattice and, therefore, slight changes in parameter values, but does not necessitate a change in crystal arrangement. Thus, in the systems studied, barium ions may be replared by feiric ions in some form, or sulfate ions by chloride or nitrite ion,, or a whole c d l may be replaced by water as suggested by Walton and Walden ( 2 2 ) . If solid solution occurs, some accounting must be made for charge balancing, since the contaminating ions are not divalent and electrical neutrality must be maintained within the crystal. Averell and Walden (2)have suggested that in the case of solid solution of permanganate in barium sulfate, electrical neutrality is achieved if a barium ion is replaced by an oxonium ion for every divalent sulfate replaced by a univalent permanganate ion. This same idea may be applied to the systems studied here. Nitrite and chloride ions may be balanced by the presence of oxonium ions in the crystal to equalize the charges. Sodium ions may also be included in the lattice in the nitrite case, as Walton and Walden ( 1 2 ) have shown that they enter into solid solution with barium d f a t e In the case of the ferric-contami-
nated precipitates the inclusion of a ferric ion and an oxonium ion to replace two barium ions will achieve electrical neutralitv. .4nother possible explanation for the inclusion of foreign and differently charged ions in the precipitates is the presence of “holes” in the imperfect crystal. .4rcording to Clark (6),silver chloride may include up to 10% cadmium chloride in solid solution. The excess of charge producwl by the replacement of univalent silver by divalent cadmium is overcome by the absence of positive ions at adjacent positions in the lattice. Similarly, the presence of two ferric and one barium ion to replace €our barium ions in zi unit cell would leave one empty position and balance the charges of the four sulfate ions present. I n chlwideand nitrite-contaminated precipitates the replacement of two sulfate ions by either two chloride 01 t a o nitrite ions could be compensated by the absence of one barium ion from a unit cell, leaving three barium ions to balancr the two sulfate and two univalent ions present.
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MOLE % WATER
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Figure 6. Relation of Lattice Constants and Unit Cell Volumes to Mole Per Cent Contamination of Water in Ferric-Contaminated Precipitates
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16 20 x io3 Figure 5. Relation of Lattice Constants and Unit Cell Volumes to Mole Per Cent Contamination 4
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The vwiation of lattice parameters with increasing contamination of the precipitates demonstrates in all three of the systems studied that solid solution of foreign substances occurs. T h e increase in the parameters cannot be ascribed to solid solution of water alon?, for in no case does the behavior of the increase correspond to that for solid solution of water in otherwise pure barium sulfate precipitates (12). In fact, in the case of nitritecontaminated precipitates the lattice expands with increasing nitrite contamination despite the decrease in water content of the precipitates. There may be solid solution of water molecules in the precipitates investigated, but there must also be solid solution of the foreign ions to account for the manner in which the lattice cxpmsion occurs. The water content of the precipitate? may be in some way related to or regulated by the contaminating ions, as shown by the large variation in water content among the three systems studied. The adsorption isotherm shown in Figures 1 and 2 for the ferric-contaminated system suggests that ferric ions in some form may be initially adsorbed to the surface of the growing precipitate; after this the precipitate continues to grow over the entrapped ions, producing a lattice of barium sulfate containing the contaminant in solid solution. i l n inspection of Figure 3 shows that although there is no significant change in the bc parameter there is an expansion in the a0 and co directions, resulting in an expansion of the unit cell volume with increased iron contamination. Since the ferric
310
ANALYTICAL CHEMISTRY
,
ion has a radius of 0.60 A. ( 5 ) while the barium ion has a radius of 1.35 A. (6), replacement of barium by ferric ions could not account for the lattice expansion which occurs. That this expansion is not due to hydration alone is shown in Figure 6, in which a plot of ao and CG and the unit cell volumes vs. mole per cent water in the ferric-contaminated precipitates is given. The data of Walton and Walden ( 1 2 ) show an increase in these parameters from 0 to 10% water and then a negligible increase for mole percentages of water greater than 10. The data of Figure 6, on the other hand, show relatively no change up to 20 mole % water and then a large increase in expansion a t mole percentages greater than 25. Therefore, the phenomenon present in the ferric-contaminated system cannot be due to solid solution of water alone. Because of the expansion of the lattice with increasing iron concentration, the greater amount of water present in these precipitates compared to “pure” barium sulfate, the difference in lattice expansions from that due to water alone, the obvious color changes of the precipitates upon heating, and the wellknown tendency of ferric ions to form complex ions, it appears that the ferric ions enter the precipitate as some sort of hydrated, complex ion. The amount of water present in the precipitates is more than adequate to account for such complexes. These ions would probably be somewhat larger than barium ion and would expand the lattice, as is observed. The degree of hydration and actual size of these complex ions, however, are unknon-n. The lattice constants for the chloride-contaminated precipitates in Figure 4 show a slight expansion along the ao, a contraction along the bo, and no change in the co axis. The expansion in a. is less than that which would be expected on the basis of the nrater content of these precipitates, while solid solution of water alone should have no effect on the bo axis (12). The unit cell volume shows a slight contraction. The chloride ion has a smaller radius than the sulfate ion (1.81 A., 5, compared to 3.00 A., 13). Solid solution of the chloride ion would give rise to a contraction of the lattice. This contraction could be overcome along the a0 and co axes, on the average, by solid solution of water, since the two effects are in opposition to one another, but there should still be a contraction in the bo direction, as this one is unaffected by water. As shov-n in Figure 4, there is less expansion in the a0 and c g direction than Tyhen there is only solid solution of water and the expected contraction occurs in the bc direction. The data are then in accord vith solid solution of both water and chloride ions. The data for the nitrite system, indicating expansion of the
unit cell in all three directions, are shown in Figure 5 . Since the precipitates show decreasing water content with increasing nitrite content, the lattice p-ould contract in the and cg directions if only solid solution of water were involved, Evidently this decrease is overcome, on the average, by an expansion due to solid solution of nitrite ions. Coprecipitation of sodium ions may have also occurred, but Walton and Walden (13) have shown that solid solution of this ion produces no variation in lattice constants except for a slight decrease in the Q direction with increasing sodium concentration. Because of the instability of the nitrite ion, the analyses for this ion may be low in comparison to the amount of nitrite originally present in the precipitates. Walton and Walden ( 1 2 ) have shown that the rigidity of the barium sulfate lattice prevents change after formation when water originally included in solid solution is expelled by heating. It is reasonable to assume that initial decomposition of the nitrite ion within the lattice would not alter the cell dimensions. Thus, possible decomposition of nitrite ion would only affect the nitrite analyses and not the calculated lattice constants. This would shift the curves of Figure 5 along the abscissa and not invalidate the conclusion based upon increased lattice constants with increasing contamination. LITERATURE CITED
d’Alcontres, G. S., Gazz. chiin. ital., 79, 609 (1949). Averell, P. R., and Walden, G. H., J . Am. Chem. SOC.,59, 906 (1937). Basche, W., and Mark, H., 2. Krist., 64, 1 (1926). Bragg, W. L., and Wood, E. A, J . Am. Chem. SOC.,69, 2919 (1947). Clark, G. L., “Applied X-Rays,” Chap. XVI, New York, McGraw-Hill Book Co., 1940. Cohen, .M.U., Rea. Sci. In&., 6, 68 (1935); 2. Krist.,94, 288, 306 (1936). Hillebrand, W.F., “Applied Inorganic Analysis,” p. 639, New York, John Wiley & Sons, 1929. Nichols, M.L., and Smith, E. C., J . Phys. Chem.,45,411 (1941). Sandell, E. B., “Colorimetric Determination of Traces of Metals,” p. 271, New York, Interscience Publishers, 1944. Schneider, F., and Reiman, V.,J . Am. Chem. SOC.,59, 354 (1937). Walden, G. H., and Cohen, >I. E., Ibid., 57, 2591 (1935). Walton, G., and Walden, G. H., Ibid., 68, 1742, 1750 (1946). Weiser, H., and Sherrick, J., J . Phus. Chem., 23, 205 (1919). Wyckoff, R., “Crystal Structures,” Chap. VIII, text pages 6, 9, 24, table page 21, illus. page 5 , New York, Interscience Publishers, 1951. RECEIVED for review July 10, 1952.
dooepted October 20, 1952.
Titrations in Nonaqueous Solutions Efects of Solvents f o r Increasing Sensitivity CHARLES W. PIFER, ERNEST G. WOLLISH, AND MORTON SCHMALL Products Control Laboratory, Hoffmann-LaRoche, Znc., Nutley, N . J.
T
HE acidic or basic strength of the solvent for nonaqueous
titration has been the subject of many investigations, the importance of which has been duly recognized. Effects attributed to the ionization constant of the solvent have played only a minor part in such studies. The purpose of this report is to demonstrate with experimental data the striking changes in the sensitivity of the reaction upon addition of solvents of low dielectric constant, which permit a further sharpening of the end point in many cases. As early as 1889, it was observed by Kabulkoff ( 1 4 ) that substances were far less ionized in alcohol than in water. This author also reported a decrease of ionization with increasing molecular weight of the alcohol. Fohn (7, 8) in 1910 was able to titrate
acids in toluene, benzene, chloroform, and carbon tetrachloride with 0.1 N sodium ethylate or amylate in the corresponding alcohol, using phenolphthalein as indicator. Conant, Hall, and Werner (1,12) pointed out the effect of the acidic or basic strength of the solvent upon the titration and studied the effect of alcohol and water on the sharpness of the end point. Many authors have since used alcohols and other solvents for increasing the solubility of organic compounds in aqueous solution and for sharpening of the indicator end points. Because only a limited number of indicators had been available for titrations in nonaqueous solutions, Davis and associates ( 2 - 6 ) investigated various new indicators for titrations in aprotic solvents.
