Complex Solubilities of AgX in 3-Methyl-2-oxazolidone For CS+the diffeirence SH’- SM’is equal to -17.2 eu.27 The differential entropy change for cesium ion exchange on gel 0.5:48 is fairly constant, averaging about -26 eu up to 8 c s = 0.6 (Figure 4). Therefore AS,, = -9 eu. The corresponding value for sadium ion exchange was -17 eu.12 This lower entropy is expected since Cs+ having less efficiency in binding water would bring less water into the exchanger. Nancollas and Til& also found that AS,, values show significant increases with increasing cation size when exchanged into an erisentially amorphous zirconium phosphateO9 Comparison of our results with those collected in Table I is instructive. To get some idea of the crystallinities of the exchangers we compared uptake of Cs+ from 0.1 M CsCl as reported in ref 6-10 with those of our exchangers.11 Our results for equilibration of 1.0 g of exchanger with 100 ml of 0.1 M CsCl were 1.45 mequivfg for 0.5:48,0.62 mequivfg for 0.8:48, 0.30 naequiv,’g for 2.5:48, 0.13 mequivlg for 3.5:48, and 0.03 meqjuivfg for 4.548. Somewhat greater uptakes were observed in the calorimetric study because of the greater volume to solid ratio. From these uptakes we concluded that all the exchangers used to obtain the results listed in Table 1 are close to 0.5:48 in behavior. On this basis the values reported in the table are reasonable. Baetsle and Ruvarac working with trace quantities obtained a AH which is almost identical with our Ai!lx value at zero load. Amphlest, et a!., working over a broader portion of the isotherm report a somewhat lower enthalpy. Nancollas and Tilak’s measured AH compares very well to our value at the same loading, i.e., -3.6 kcal mol-l at 2 mequivfg. Similarly the free energies reported only refer to that limited portion of Ihe isoSherm where cesium ion is preferred by the exchanger and therefore are negative or onlx slightly positive.
1817
References and Notes (1) (a) This work is part of a cooperative research program jointly sponsored by The National Science Foundation and the Swedish Natural Science Research Council. The Swedish portion is under the direction of Professor Sten Ahriand, Lund University, Lund, Sweden. The present work was supported by NSF Grant No. GP-8108.(b) Postdoctoral Fulbright Fellow, Lund University, Lund, Sweden. (2) A. Clearfield and J. A. Stynes, J. lmrg. Nucl. Chem., 26, 117 (1964). (3) J. Albertsson, Acta Chem. Scand., 20, 1689 (1966). (4) S. Ahrland, J. Aibertsson, A. Alnas, S. Hemmingsson, and L. Kullberg, Acta Chem. Scand., 21, 195 (1967). (5) A. Clearfield, A. Oskarsson, and C. Oskarsson, Ion Exchange Membranes, 1, 91 (1972). (6) L. Baetsie, J. lnorg. Nucl. Chem., 25, 271 (1963). (7) C. B. Amphlett, P. Eaton, L. A. McDonald, and A. J. Miller. J. lnorg. Nucl. Chem., 26,297 (1964). (8) J. P. Harkin, G. H. Nancollas. and R. Paterson, J. lnorg. Nucl. Chem.. 26, 305 (1964). (9) G. H. Nancollas and B. V. K. S. R. A. Tilak, J. Inorg. Nucl. Chem., 31, 3643 (1969). (10) A. Ruvarac, Bu//. Boris Kidric lnst. Nucl. Sci., 20,33 (1969). (11) A. Clearfield and A. Oskarsson, Ion Exchange Membranes, in press. (12) A. Clearfield and L. H. Kullberg, J. fhys. Chem., 78, 152 (1974) (13) J. GrenthB, H. Ots, and 0. Ginstrup, Acta Chi” Scand., 24, 1067 (1970). (14) H. S. Harned and B. B. Owen, “Physical Chemistry of Electrolyte Solutions,” Reinhold, New York, N. Y.. 1957, p 707. (15) S. R. Gunn, J. Phys. Chem., 69,2902 (1965). (16) K. S. Pitzer, J. Amer. Chem. SOC.,59, 2365 (1937). (17) A. Clearfield and J. M. Troup, J. Phys. Chem.. 77, 243 (1973). (18) A. Clearfield and G. D. Smith, Inorg. Chem., 8,431 (1969). (19) A. ClearfleldandA. S. Medina, J. fhys. Chem., 75,3750 (1971). (20) A. Clearfield, G. H. Nancollas, and R. H. Blessing, “Ion Exchange and Solvent Extraction,” Vol. 5, J. A. Marinsky and Y. Marcus, Ed., Marcel Dekker, New York, N. Y., 1973, pp 37-38. (21) G. H. Nancolias, “interactions in Electrolyte Solutions,” Elsevier, New York, N. Y., 1966. pp 120-125. (22) G. Eisenman, Biophys. J., 2, 259 (1962). (23) A. Clearfield. W. L. Duax, A. S. Medina, G. D. Smith. and J. R. Thomas, J. Phys. Chem., 73,3424 (1969). (24) A. Clearfield and J. M. Troup, J. fhys. Chem., 74, 314 (1970). (25) A. Clearfield, W. C. Duax, J. M. Garces, and A. S.Medina, J. Inorg. Nucl. Chem., 34,329 (1972). (26) H. S. Sherry “ion Exchange,” Vol. II, J. A. Marinsky, Ed., Marcel Dekker, New York, N. Y., 1968. (27) D. R. Rosseinsky, Chem. Rev., 65,467 (1965).
Complex Solubilities of the Silver Halides in 3-Methyl-2-oxazolidone ark Salomon Power Sources TechnicalArea, U. S. Army Electronics Technology and Devices Laboratory, Fort Monmoufh, New Jersey 07703 (Received April 75, 1974) Public.ationcosts assistedby the U.S. Army Electronics Command, Electronics Technology and Devices Laboratory
The complex solubilities of the silver halides have been determined in the aprotic solvent 3-methyl-2-oxazolidone. The results appear to be typical of the general behavior of aprotic solvents toward the solubility of the silver halides. 3-Methyl-2-oxazolidone is an interesting solvent in that it is practically isodielectric with water, is stable in the presence of metallic lithium, will dissolve appreciable amounts of alkali metal salts, and is completely miscible with water.
Introduction Propylene (carbonate (PC) and dimethyl sulfoxide (DMSO) are popular aprotic solvents for use in electrolyte
solution studies mainly because of their high dielectric constants and ease of purificati0n.l Recently Huffman and Sears2 have described the physical properties of a series of heterocyclic carbamates which exhibit properties similar to The Journal of Physical Chemistry, Vol. 78, No. 18, 1974
181
Mark Salomon
PC. One of these solvents, 3-methyl-2-oxazolidone
0-
?:W~--CH~-O-CO-NN-CH~ I I
appears to be promising for electrolyte solution studies. At 25' it has a dielectric constant of 77.5 D, a density of 1.1702 g/ml, and a viscosity of 2.450 C P .This ~ solvent is completely miscible with water, will dissolve the alkali metal halides and tetraakkyl salts, and appears to be stable in the presence of metallic lithium. In this paper the complex solubilities of the silver halides are reported in 3-methyl-2-oxazolidone (3Me20x) at 25". The equilibria studied are Ag+ -I
A$
+
x-
= AgX
2X- = A&-
(KSJ
(1)
(6)
(2)
Ag* i 3X- = AgX,2(&) (3 1 K,o is the solubility product, and @ n or stability constant.
Materials. Tetrapropylammonium perchlorate, chloride, bromide, and iodide (TPAP, TPAC1, TPABr, and TPAI, respectively) were obtained from Eastman Organics and purified by the methods described by Mann.3 e2Ux was supplied by Sears and had been purified by successive fractional freezings as described in Sears' pa,per.2 The solvent was placed in an argon filled VAC drybox ( 0 2 and H20 content less than 1 ppm) and dried over type 4A molecular sieves for at least 2 days. All solutions reported here are of a constant ionic strength of 0.1000 M A Mettler I115 balance was used for all weighings inside the drybox. The TPAP was used as the inert electrolyte to maintain the constant ionic strength. A11 Lhe t e t r a ~ r o ~ y l ~ m m o n i usalts m were colorless in 3Me20x9but the titrant, 0.1000 M AgC104, had a tendency to turn slightly yellow upon standing for more than 3 days. When this occurred, fresh 0.1 M Age104 solution was preared. No attempt was made to shield any of the solutions from light. Potentiometr-ic Meosurements, The titration cell used is identical with the one described earlier.4 The reference electrode was AiJi0.1 M AgC104 in 3Me20x and the indicator electrode was a coiled, etched silver wire. The TPAX TPAP solution (20.0 ml) was placed in the cell which was fitted with a Gilmont 12.000-ml capacity buret. The buret was filled with 0.1000 M AgC1U4 solution in 3Me20x. The cell was then removed from the drybox and placed in a water bath regulated at 25.0 f 0 . 0 5 O . Magnetic stirring was maintained throughout the titration. The emf measurements were made with a Doric Model DS-100 integrating ~ ~ c r o v ~ ~ The ~ mreproducibility e t ~ ~ r ~ of this instrument is f O . l mV.
+
esults The silver ion concentrations were obtained from the Nernst relation E = E' i( R T / F ) In [Ag']
(4 )
where E' is a formal potential and contains contributions from the liquid junction and nonideality of the electrolyte solutions. In eq 4, E i s the measured potential and the The Journal of Physicdl Chemistry, Vol. 78, No. 18, 1974
0.25-
1.00
:
0!5
0
1.0
11s
2!0
cAU%
Figure 1. Potentiometric titration curves. E is in V and CAS/+ is the bromide (U),and iodide (A). relative concentration for chloride (0). The s o l i points were not included in the calculations of the P's.
Nernst slope was obeyed to within f0.5 mV; E' was found to equal 58.0 f 0.6 mV? The titration curves for the three halides are shown in Figure 1. In this figure, the measured emf is plotted against the relative concentration term Cx; CAg is the total silver present and CX is the total halide present in mole/liter. The titrations for TPACl and TPABr appear normal but the TPAI titration presented difficulties. In the first place equilibrium took about twice as long (>IO min) in the region around the point for CA&X = 0.5. In addition a precipitate was not observed for some time after this point. A t first it was thought that the failure to obtain a precipitate a t CAg/cI = 0.5 was due to the formation of soluble AgI. However a value for the equilibrium constant 01 for Ag' + I- = AgI (in solution) (5) could not be refined (see below). This problem was encountered in earlier work^^,^ and it is concluded that this phenomenon is due to a very slow approach to equilibrium. Polynuclear complexes such as AgzI+ which are found in some aprotic solvents in the presence of excess Ag+ were not considered here since they are negligibre in excess I-.' Due to the failure to reach equilibrium, several points had to be omitted in the calculation of the equilibrium constants. These points are denoted by solid points in Figure 1.
The stability constants were obtained by a least-squares fit to those points in the unsaturated region of the titration curve. In this calculation, the relative error function, U, is minimized; for i data points, U is defined by i
u
=
c 1
(1 - Y i / f Y J 2
(6)
where
Yi = ( G , 2
- [X-l,)/r&+I
(7 1
and n
fYi
=
C npn[x-Iin 1
(8)
An iterative method was used to generate the 8, values as described earlier.* The solubility products, KSo, were ob-
Complex Solubilities of AgX in 3-Methyl-2-oxazolidone
1819
TABLE I: Molar Equilibrium Constants at 25 __l_-__l___________
-Log I 3 be refined. For the iodide case, the leastsquares method generated only a value for fi2. The results of these calcu ations are given in Table I. Also included in Table I are the 0, m d KsO values corrected to zero ionic strength. For these calculations, the activity coefficient for 0.1000 M solul ions was obtained from the Davies equationg
bye--Wuckel A factor has the value 0.520 The standard deviations for fino and KSo0 have not been included in Table I. They are expected to be about 10% higher than the experimental values due to the uncertainty in the use of eq I f . There were no unexpected results in the solubility behavior of the eilver halides in 3Me20x. K,o values are generally smaller and K,2 and Ks3 values are much larger than they are in water. The fact that Ks2 and Ks3 are greater than unity is typical of the general behavior of AgX in aprotic organic: solvents.loJ1 In fact there appears to be no exception to this observation. Addition of inorganic protic solvents such 3s H2CWJ3 or aprotic solvents such as SO2* to the aprotic organic solvent does result in Ks2 and Ks3 values less than unity. In aprotic organic solvents containing water, the halide ions are stabilized by hydrogen bonding ~ ~ a r t i the c ~ a ~and ~ by ~ strong electrostatic inC1-~ ion)
teractions with the water dipole. In aprotic organic solvents containing SOz, the halide ions are probably stabilized by bonding through the 3p orbitals on the sulfur andlor by charge t r a n ~ f e r . ~ JIn * , the ~ ~ pure aprotic solvent, anion solvation occurs mainly via mutual polarization.16 For this type of interaction it is expected that the stability of the anion should increase as anion size increases; this is the major factor leading to large values for Ks2 and Ks3.I6.l7 The increased stability of AgX2- compared to X- can be demonstrated by considering single ion standard free energies of transfer from water to 3Me20x, AGto(ion). In an earlier paper? it was proposed that AGto(Ag+)is given, approximately, by AGto(Ag') = {AGto(Ag*,AgClZ-)
+
AG,'(Ag', AgBr2-))/2 (12) The terms on the right-hand side of eq 1 2 are experimental and simple substitution gives AGto(Ag+)N -1.0 kcal/mol. This result was used to obtain the AGto(anion) values which are given in Table 11. The values of ACto(anion) are typical of those found in other aprotic organic solvents.1°J7 Further studies in mixed solvents are in progress and will be reported at a later date. Acknowledgment. The author would like to express his gratitude to Professor P. G., Sears who kindly supplied the purified solvent. Supplementary Material Available. Listings of the experimental data (the concentrations and em€'$) will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper only or microfiche (105 X 148 mm, 24X reduction, negatives) containing all of the supplementary material for the papers in this issue may be obtained from the Journals Department, American Chemical Society, 1155 16th St., N.W., Washington, . C. 20036. Remit check or money order for $3.00 for photocopy or $2.00 for microfiche, referring to code number JPC-74-1817. References and Notes (1) A. K. Covington and T. Dickinson, Ed., "Physical Chemistry of
Organic Solvent Systems," Plenum Press, London, 1973. (2) H. L. Huffmanand P. G. Sears, J. Solution Cbem., 1, 187 (1972). (3)C.K. Mann, Advan. Nectroanal. Cbem., 3, 57 (1969). (4)M. Salomon and 8 . K. Stevenson, J. Pbys. Cbem., 77, 3002 (1973). (5) See paragraph at end of paper regarding supplementary material. (6)R. Alexander, E. C. F. KO, Y. C. Mac, and A. J. Barker, J. Amer. Cbem. SOC.,89,3703(1967). (7)D. C. Luehrs and K. Abate, J. Inorg. Nucl. Cbem., 30, 549 (1968). (8)L. G.Sillen, Acta Cbem. Scad., 16, 159 (1962). (9)C. W. Davies, "Ion Association," Butterworths, London, 1962. (IO)J. N. Butler, Advan. Necfrochem., 7, 77 (1970). (11) M. Salomon, ref 1, Chapter 2.2. (12)J. N. Butler, D. R. Cogley, and W. Zurosky, J. E/ectrochem. Soc., 115,
445 (1968). (13)J. C. Synnottand J. N. Butler, J. Pbys. Cbem.,73, 1470(1969). (14)M. Salomon and B. K. Stevenson, manuscript in preparation. (15)E. R. Lippincott and F. E. Welsh, Specfrocbim. Acta, 17, 123 (1974). (16)A. J. Parker, Cbem. Rev., 69,l(1969). (17)
C.M. Criss and M. Salomon, ref 1, Chapter 2.4. The Journal of Physical Chemistry, Vol. 78, No. 18, 1974
