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(23) 8. Pullman, Ph. CourriBre, and H. Berthod, J. M e d . Chem., 17, 439 (1974). (24) (a) B. Pullman, A. Pullman, H. Berthod, and N. Gresh, Theor. Chim. Acta, 40, 93 (1975); (b) B. Pullman, H. Berthod, and N. Gresh, FEBS Left., 53, 199 (1975). (25) M. J. Mantione and J. P. Daudey, Chem. Phys. Lett., 6, 93 (1970). (26) M. J. Huron and P. Claverie, Chem. Phys. Left., 9, 194 (1971). (27) V. G. Dashevsky and G. N. Sarkisov, Mol. Phys., 27, 1271 (1974). (28) N. S. Bayless and L. Hulme, Austr. J. Chem., 6, 257 (1953). (29) S. Basu, Adv. Q. Chem., 1, 145 (1964). (30) T. Abe, Bull. Chem. SOC.Jpn., 41, 1260 (1968). (31) A. Schweig, Chem. Phys. Lett., 3, 542 (1969). (32) (a) 0. Sinanoglu in “The World of Quantum Chemistry”, R. Daudel and B. Pullman, Ed., Reidel, Dordrecht, Holland, 1974, p 265; (b) Chem. Phys. Left., 1,340 (1967); (c) Adv. Chem. Phys., 12, 283 (1968); (d) Theor. Chim. Acta, 33, 279 (1974); (e) in “Molecular Associations in Biology”, Academic Press, New York, N.Y., 1968, p 472. (33) D. L. Beveridge, M. P. Kelly, and R. J. Radna, J . Am. Chem. Soc., 96, 3769 (1974). (34) A. J. Hopfinger in ”Molecular and Quantum Pharmacology”, E. D. Bergmann and B. Pullman, Ed., Reidel, Dordrecht, Holland, 1974, p 131. (35) J. Hylton, R. E. Christoffersen, and G. Hall, Chem. Phys. Lett., 24, 501 (1974). (36) (a) M. J. Huron and P. Claverie, J. Phys. Chem., 76, 2123 (1972); (b) 78, 1853 (1974); (c) 78, 1862 (1974). (37) F. Birnstock, H. J. Hoffman, and H. J. Kohler, Theor. Chim. Acfa, 42, 311 (1977). (38) C. C. J. Roothaan, Rev. Mod. Phys., 23, 69 (1951). (39) W. J. Hehre, R. F. Stewart, and J. A. Pople, J. Chem. Phys., 51, 2657 (1969). (40) R. Ditchfield, W. J. Hehre, and J. A. Pople, J. Chem. Phys., 54, 724 (1971). (41) A. I. Kitaygorodsky, Tetrahedron, 14, 230 (1961). (42) (a) P. Claverie in “Localization and Delocalization in Quantum Chemistry”, Vol. 11, 0. Chahret, R. Daudel, S. Diner, and J. P. Malrieu, Ed., Reidel, Dordrecht, 1976, p 123. (b) P. Claverie, “Elaboration of Approximate Formulas for the Interactions between Large Molecules. Applications in Organic Chemistry”, in ”Intermolecular Interactions: from Diatomics to Biopolymers”, B. Pullman, Ed., Wiley, New York, N.Y., in press; (c) A. T. Amos and R. J. Crispin, Mol. Phys., 31, 147 (1976); (d) in “Theoretical Chemistry, Advances and Perspectives”, Vol. 2, H. Eyring and D. Henderson, Ed., Academic Press, New York, N.Y., 1976, p 1. (43) F. Danon and K. S. Pitzer, J. Chem. Phys., 36, 425 (1962). (44) P. Claverie, “Elaboration of Approximate Formulas for the Interactions between Large Molecules. Application in Organic Chemistry”, section V-B and Appendix F in “Intermolecular Interactions: from Diatomics to Biopolymers”, B. Pullman, Ed., Wiley, New York, N.Y., in press.
Lucas et al. (45) (a) J. Caillet and P. Claverie, Biopolymers, 13, 601 (1974); (b) Acta Crystallogr., Sect. A , 31, 448 (1975); (c) J. Caillet, P. Claverie, and B. Pullman, Acfa Crysfallogr., Sect. B , 32, 2740 (1976). (46) J. P. Chandler, minimization program STEPIT, Quantum Chemistry Program Exchange, Program No. 66, Indiana University, Bloomington, Ind. 47401. (47) F. M. Westheimer and J. G. Kitkwood, J. Chem. phys., 6, 513 (1938); J. G. Kirkwood, ”Theory of Solutions”, Gordon and Breach, New York, N.Y., 1968, p 208. (48) L. Onsager, J. Am. Chem. Soc., 58, 1486 (1936). (49) B. Linder, Adv. Chem. Phys., 12, 225 (1967). (50) J. A. Abbott and H. C. Bolton, Trans. Fararky Soc., 48, 422 (1952). (51) A. Wada, J. Chem. Phys., 22, 198 (1954). (52) F. Buckley and A. A. Maryott, J. Res. Nafl. Bur. Stand., 53, 229 (1954). (53) H. Eyring, D. Henderson, B. J. Stover, ana E. M. Eyring, “Statistical Mechanics and Dynamics,” Wiley, New York, N.Y., 1964, Chapter 14, Section 5, p 401. (54) H. H. Uhlig, J. Phys. Chem., 41, 1215 (1937). (55) D. D. Eley, Trans. Faraday Soc., 35, 1281 (1939). (56) H. L. Clever, J. Phys. Chem., 62, 375 (1958). (57) J. H. Saylor and R. Battino, J. Phys. Chem., 62, 1334 (1958). (58) (a) R. C. Tolman, J. Chem. Phys., 16, 758 (1948); 17, 118 (1949); (b) J. G. Kirkwood and F. P. Buff, J. Chem. Phys., 17, 338 (1948); 18, 991 (1950); (c) F. P. Buff, ibid., 19, 1591 (1951); 23,419 (1955); (d) S. Kondo, ibid., 25, 662 (1956). (59) D. S. Choi, M. S. Jhon, and M. Eyring, J . Chern. Phys., 53, 2608 (1970). (60) R. B. Herman, J. Phys. Chem., 75, 363 (1971); 76, 2754 (1972). (61) (a) R. A. pierotti, J. phys. Chem. 67, 1840 (1963); (b) 89,281 (1965). (62) H. S. Frank and M. W. Evans, J. Chern. Phys., 13, 507 (1945). (63) G. Nemethy and H. A. Scheraga, J. Chem. Phys., 36,3382, 3401 (1962). (64) N. Bjerrum, Trans. Faraday Soc., 24, 445 (1927). (65) D. L. Beveridge and G. W. Schnuelle, J. Phys. Chem., 79, 2562 (1975). (66) D. L. Beveridge and G. W. Schnuelle, J. Phys. Chem., 78, 2064 (1974). (67) G. W. Schnuelle and D. L. Beveridge, J. Phys. Chem., 79, 2566 (1975). (68) J. D. Payzant, A. J. Cunningham, and P. Kebarle, Can. J. Chem., 51, 3242 (1973). (69) D. H. Aue, H. M. Webb, and M. T. Bowers, J . Am. Chem. Soc., 94, 4726 (1972). (70) P. Souchay in “Fondements th6oriques des red-ierches sur les actions intermolbculaires”, 278 collcque national du CNRS, &litions du Centre National de la Recherche Scientifique, Paris, 1966, p 221. (71) The total charges given do not represent only the electronic populations; besides these,they also take into account suitable “effective” charges which simulate atomic dipole moments.
Effects of Sodium and Lithium Halides and of Hydrochloric and Hydrobromic Acids on the Coacervation of Aqueous Solutions of Tetraalkylammonium Halides A. Mugnier de Trobriand, M. Lucas,* DGR, B.P. No. 6, 92260 Fontenay-aux-Roses, France
J. Steigman,’ and L. L-Y. Hwang2 Department of Chemistry, Polytechnic Institute of New York, Brooklyn, New York 11201 (Received May 9, 1977) Publication costs assisted by Commissariat 5 Energie Atomique
The coacervation of aqueous solutions of Pn4NBr, Hex4NBr,Hex4NC1,and HeptlNCl was studied at 35 “C in the absence and presence of LiC1, LiBr, NaC1, NaBr, HC1, and HBr. The “invasion” of the R4NX-richaqueous solution by the acids is very marked. A comparison is made between the anion-exchange resin Dowex I X-10 and the R4NX-rich aqueous phase. Several theories which had been advanced to explain the “acid effect” in resins and the corresponding phenomena in quaternary ammonium salt solutions are discussed. Introduction Aqueous solutions of tetraalkylammonium halides of sufficient chain length separate into two aqueous phases a t room temperature upon the addition of inorganic electrolytes. Since one aqueous phase contains most of the 0022-3654/78/2082-0418$01 .OO/O
inorganic salt, with a small fraction of the quaternary ammonium salt, and the other is rich in the organic salt, these coacervates have been used for the extraction of anionic complexes of ions such as U022+,and have been found to be very similar in behavior and selectivity t o 0 1978 American
Chemical Society
Coacervation of Aqueous Solutions of Tetraalkylammonium Halides
anion-exchange resin^.^ Thus further study of these systems may add to the understanding of various properties of the resins. The present paper deals largely with some thermodynamic properties of coacervates based on the ternary systems R4NX, MX, and HzO, in which X is either chloride or bromide, and M represents lithium, sodium, or hydronium ions. R4N represents various quaternary ammonium structures which are described below. In addition, comparisons are made between some properties of the resin Dowex I X-10 and those of the ternary systems. Finally, the “acid effect” in the resin and in quaternary ammonium salt solutions is examined, and a number of explanations are discussed. Experimental Section Bu4NBr (tetra-n-butylammonium bromide) and Et4NC1 (tetraethylammonium chloride) of polarographic grade (Southwestern Analytical Chemicals) were used without further purification, Pn4NBr (tetra-n-pentylammonium bromide) and Hex4NBr (tetra-n-hexylammonium bromide), obtained from Eastman Kodak Co., were precipitated with hexane from benzene solution, and dried in vacuo. Hept4NC1 (tetra-n-heptylammonium chloride) was used without purification after titration with silver nitrate showed that it was pure. Hydrogen bromide gas, a Matheson product, was used shortly after arrival. All other chemicals were reagent grade. The infrared spectra of the various R4NX-HX compounds were recorded on a Perkin-Elmer Model 521 grating double beam spectrometer. The solid products were suspended in nujol, a mull was quickly spread over a KBr crystal, quickly covered with a second crystal, and wrapped around the-crystals’ edges with parafilm. Samples were scanned against air. Bu4NBr, etc. crystals were also run this way. Gaseous HBr was run in a gas cell against air. The R4NX-HX compounds were prepared by passing HBr gas through a P205drying tube into a flask containing the R4NX crystals. The flask was then connected with two drying tubes containing solid KOH and Pz05,respectively; removal of excess HBr was carried out for 4 days. Products made with hydrogen chloride gas were similarly treated, except that saturation with the gas was carried out three times. The solubility of HBr in Bu4NBr solutions was determined by bubbling the gas into solutions with differents concentrations of the organic salt for at least 5 h with constant stirring. Aliquots (0.2 mL) were titrated with standardized NaOH solution. The HBr solutions were exposed to more gas for an additional hour and reanalyzed. Densities were determined by weighing 2-mL aliquots quickly to f1.0 mg. The molality of HBr in the solutions was then calculated. The reproducibility of analysis was better than 1%. The coacervates were stirred at 35 f 0.2 OC overnight, and after decantation each phase was analyzed for halide content and (where applicable) by titration with alkali. Sodium content was determined by flame photometry. LiCl and LiBr in solution were analyzed argehtometrically after the quaternary ammonium halide has been removed by repeated benzene extraction. Results and Discussion The molal concentrations of each electrolyte in the ternary systems containing a common anion X- (that is, MX, RdNX, HzO) which at a given temperature at atmospheric pressure separate into two aqueous phases are shown in Figure 1 for the bromides, and in Figure 2 for
The Journal of Physical Chemistry, Vol. 82, No. 4, 1978 419
R
m
NBr ( l o w e r phase)
.
,&IC_.
.
’
’
---0
’
’
1
’
-- - - --_
m HBr, LiBr,NaBr
in Lower p h a s o
.
.. I
lo
.
.
‘
.
’
.
.
.
.
l
K,,,,, LNa i EBrr
I ’
0
0
m HBr,LiBr,NaBr
10
i n lower phase
Figure 1. (a) Plots of R4NBr molality in the lower phase against inorganic electrolyte molality in the same phase: open circles, NaBr; filled circles, HBr; crosses; LiBr; dotted line, HBr, Hex,NBr. (b) Plots of R4NBr molality and MBr molality in the upper phase against inorganic electrolyte molalii in lower phase.
TABLE I: Solubility of Hydrogen Bromide in n-Bu,NBr Solutions a t 25 ’C n-Bu,NBr, m 0.00 3.19 5.28
HBr solubility,
m 22.1 24.1 21.1
n-Bu,NBr, m 6.76 7.84
HBr solubility,
m 32.5 39.6
the chlorides. The molalities of R4NX and MX (or HX) in the R4NX-rich upper phase are plotted as the ordinate against the molality of MX or HX in the lower phase which contains only a small concentration of R4NX. Figure l a gives plots of the concentrations of R4NBr in the lower (inorganic) phase. Since the molality of Hept4NC1 in the lower phase was always less than lo-’ mol kg-’, it is not shown in the figures. Table I gives the solubility of HBr in Bu4NBr solutions. Coacervation of H20-R4NX Systems Although we have observed coacervation with highly water-soluble salts such as Bu4NBr and Bu4NC1 upon addition of inorganic electrolytes, in this paper we have considered mainly quaternary ammonium salts which exhibit coacervation in the absence of such electrolytes. Such salts are Hex4NC1,Hept4NC1,Pn4NBr, and Hex4NBr. It is apparent from the plots that the water content of the “organic” phase, rich in R4NX, is about 3-5 mol of water for each mol of R4NX, that is for 20-28 methyl and methylene groups.
420
The Journal of Physical Chemistry, Vol. 82, No. 4, 1978 m (upper phase)
Lucas et al. *l2
I
15
HCI/HeptkNCI
,
=---I?-
. HCVDowex 1 -0
-
-0
-+
1
HCI / Pn4 NCI
--- -
- -LiCi/Dowexl -0
-
L E I / Hcpt4 NCI
-o--o-NaCl/Dowexl NaCI/PnkNCl
1
I 0
10
5
15
m Inorg, salt or a c i d in resin o r R,,NX phase
Figure 3. Plots of the Harned coefficient in mixed inorganic electrolyte, R,NX (or mixed inorganic electrolyte, resin) against the molality of the inorganic electrolyte in the same solution. 210g g f (HXI
-1
- 0
m HCI LiCl NaCl in 1w.r phase
10
Figure 2. Plots of R4NCIand MCI molality in the upper phase against inorganic electrolyte molality in the lower phase: open circles, Hept,NCI/LiCI; filled circles, Hept,NCI/HCI; crosses, Pn,NCI/NaCI.
The low solubility of Hex4NBr and of Hex4NC1in the less concentrated aqueous phase may possibly be due to the same reasons that have been given for the low solubility of hydrocarbons in water.4 The organic phase in equilibrium may be viewed as a kind of fused salt with an organic cation with strong dispersion forces between alkyl chains and a hydrated anion. Ternary Systems Li(Na, H)X, R,NX, H 2 0 It is apparent from Figures l b and 2 that the increase in salt in the lower (inorganic) phase increases the R4NX molality in the upper (organic) phase, and that “invasion” of that phase by the inorganic salt is only moderate. The R4NX concentration in the inorganic phase decreases with an increase in MX salt concentration. In keeping with the above interpretation, this may be a salting-out effect similar to that which occurs with hydrocarbons, and which has been treated semiq~antitatively.~~~ The corresponding decrease in water activity may induce a smaller solvation of the anions in the organic phase, hence an increase in R4NX molality in the organic salt-rich phase. In the case of hydrogen halides as electrolytes, the picture is not the same. The increase in HX in the inorganic phase corresponds to a decrease in the R4NX molality in the organic phase and to a very large invasion of that phase by HX. It is treated in more detail below. Acid Invasion Since the inorganic phase contains only small concentrations of the quaternary ammonium ions it is possible to treat the acid activity coefficient on the molality scale r*(W) as if it is equal to that of the pure acid aqueous solution at the same acid molality and then to compute the activity coefficient y+(org) of the acid in the organic phase. Furthermore this activity coefficient can be used to test a kind of Harned’s rule7 2 log y,(org) = 2 log 7,’ - a i Z m R , N X (1) in which y * O is the activity coefficient of the pure acid aqueous solution at m equal to the total molality (mHX +
Hex N C i n ‘ -2
4
t 0
5
A P n NBr
10
m R‘NX
Figure 4. Plots of 2 log y*(HX) at infinite dilution of acid in the R4NX aqueous solution against R,NX molality.
mRdNX) in the organic phase. We have computed values of a12at various acid concentrations in the organic phase and found, as shown in Figure 3, that a12is reasonably constant for a large range of HC1 concentrations in mixtures with Hept4NC1. The same has been found for a few values with Hex4NC1as the salt (not shown on the figure.) Then it becomes possible to calculate the HC1 activity coefficient at infinite dilution in R4NC1concentrated solutions assuming that a12can be extrapolated linearly to zero acid concentration. For HBr-R4NBr mixtures it is not possible to use the same procedure since y,(W) is now known for HBr at the molalities required for the computation. However it is possible to give a value for log y*(org) of HBr at infinite dilution in the Pn4NBr-rich phase, since the plots of log y+ in Pn4NBr against HBr molality in the lower “inorganic” phase are approximately linear in the range 1-5 m as shown in Figure 6. The values of 2 log y* for the acids at infinite dilution in the R4NX-rich phase have been plotted in Figure 4 against the molality of R4NX at zero inorganic electrolyte concentrations, which are found from Figures 1and 2. We have also plotted values of ya for acids at infinite dilution in various Et4NBr and Me4NC1solutions, taken from Table I in ref 8. These values can also be used to compute another parameter a’12as a function of R4NX molality, using 2 log yiHX(O HX, m R4NX) = 2 log Y + H X ( m ) R4NX) (2) where Y + ~ ~is( the ~ ) activity coefficient of the pure acid at that molality. These values of a’12are plotted in Figure 5. These plots show that a’12increases in the order Me4N+
Coacervation of Aqueous Solutions of Tetraalkylammonium Halides
The Journal of Physical Chemistry, Vol. 82, No. 4, 1978 421 A ( ~ ~ ~ / ~ ~ ~ ) / A o ~ ~ ~
0 0
5
lorn R L N X
15
Flgure 5. Plots of the Harned coefficient
for hydrogen halides at infinite dilution in R4NX (and resin) solution of different molality: open circles, R4NX, Me4NCI; filled circles, Et4NCI; crosses, Et4NBr; open square, Hex4NCI; filled square, Hept4NCI. 2 1 0 g $ ~ ( H X) i n upper phase
2
1
0
-1
0
5
mHX 10 Lwrr phase
Figure 6. Plots of the acid activity coefficient in the mixed (R4NX, HX) “organic” solution against acid molality in the lower “inorganic” phase (where X- is an anion).
< Et4Nt = Hept4Nt, exhibiting a saturation effect since the increase in chain length from ethyl to hexyl has about the same effect on as the increase from methyl, and since the effect decreases when the alkyl chain length is further increased beyond that of hexyl. In summary, the acid enters the “organic” phase of the coacervates to a very marked extent, unlike the alkali metal salts. The interactions of hydrogen halides and quaternary ammonium salts in noncoacervating (i.e., homogeneous) aqueous solutions have also been studied, and show similar effects. From electrochemical measurements it has been concluded that whereas an alkali metal halide raises the activity coefficient of the corresponding acid,8 the quaternary ammonium halide will lower it.9-11 Furthermore, in concentrated LiCl solutions, the escaping tendency of even very dilute HC1 becomes so great that it can be quantitatively displaced from solution by a current of nitrogen.12 In contrast, the solubility of HBr in Bu4NBr solutions increases very markedly with added organic salt. This is shown in Table I. A number of explanations have been advanced for these interactions. Two are discussed below: a quantum-mechanical hypothesis, based on infrared spectra, and a structural hypothesis, based on x-ray diffraction studies of solutions. Hydrochloric acid solutions show a characteristic continuum in the infrared spectrum. Zundel and his coworkers have shown that the state of the acid proton in water is best represented by the grouping (H20.H+-OH2), in which the proton oscillates between two water molecules, rather than the grouping ( & o + I z H ~ o ) .They ~ ~ explained the absorbances of the infrared continuum in terms of the projected structure, which is associated with a double minimum in the potential energy curve of the particle, with a small maximum between the two minima. This results
0.2
0.4
*,2
Flgure 7. Plots of the relative continuum I R absorbance of 1 m HCI, 5 m MCI solutions against the Harned coefficient a’,* for acid in the same salt solutions.
in a high polarizability of the projected structure, and a great sensitivity to its immediate environment. Schioberg and Zundel found that lithium chloride produces a decrease in the absorbance of the infrared continuum of 5 m HC1 which is proportional to its concentration, whereas Et4NC1 produces a concentrationdependent increase.14 They explain the difference in the behavior of lithium and quaternary ammonium salts toward HC1 in terms of the difference in ions size, and hence in the electrical field in the vicinity of the ions. The double minimum potential associated with the structure (HzO. H+-HzO)is more symmetrical in the presence of the large Et4N+ ions. Lithium ions, because of their high charge density, will displace the easily polarized proton toward one of the H5-02+ water molecules, effectively producing H30+.Hz0, where the much larger organic cation will produce much less polarization. Figure 7 shows that there is a correlation between the integrated absorbance of the infrared continuum of HC1 in the presence of various salts, and the values of a’12for these salts, the latter taken from ref 15. As a consequence, since the symmetrical H5OZtstructure is more stable than the unsymmetrical H30+-H20structure, quaternary ammonium salts will decrease the activity coefficient of HC1, whereas LiCl and similar inorganic salts will raise it. Another explanation is based on the ability of quaternary ammonium salts to organize cages of hydrogenbonded water molecules around them~elves.~~,~’ This leads to a possible explanation of the interaction of acids and quaternary ammonium salts. Narten and Triolo’s studied the x-ray diffraction pattern of H20.HC1 solutions. They concluded that the basic solvent structure was retained even in saturated HC1 solutions, with some local distortion to accomodate H30+and C1- ions, but was not “broken”; in particular, there was not much evidence for hydration of the chloride ions. This is in marked contrast to the results obtained with LiCl solutions;20the basic water structure was “broken”, and replaced by hydration shells around both the Lit ions and the C1- ions. If a compound such as n-Bu4NC1 can dissolve in water with a minimum disturbance of the water structure,lg and if HC1 added to it will do the same, then the chloride ions of the two electrolytes may be similar in their solvent effects (because of their accompanying cations), they can accomodate each other, and lower each other’s escaping tendency. On the other hand, a mixture of LiCl and HC1 would have, in crude terms, two “different” kinds of chloride ion solvates to deal with, one “hydrated” (in the vicinity of lithium ions) and one “unhydrated” (near H30f ions). While only averaged properties of chlorides in mixtures can be detected, it is reasonable to expect that the escaping tendencies of the two compounds will be raised. Comparison of t h e Aqueous RINX Systems w i t h Anion-Exchange Resins The anion-exchange resin Dowex I X-10 shows some similarities to the R4NX systems in the invasion of the
422
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The Journal of Physical Chemistry, Vol. 82, No. 4, 1978
resin by HC1, and by NaCl aqueous solutions. Nelson and Kraus,22who studied the behavior of Dowex I X-10 toward a great many electrolytes, showed that a Harned rule could be applied to the resin, assuming that it behaved like an aqueous electrolyte solution of a salt of cohcentration mR which is equal to the number of resin sites for 55.51 mol of water. In the case of NaC1, the equation is
(3) In this equation y + O is the activity coefficient of NaCl in water at the total molality ( m R+ mN,cJ(resin phase). Figure 3 shows that cyl2 values are similar for resin and R4N solutions in the case of NaCl and of HC1, in spite of the fact that these two electrolytes behave very differently toward either resin or quaternary ammonium salt. Thus, the resin invasion by HCl (the acid effect,*l which is very large compared to invasion by NaCl solutions, can probably be explained by the same interaction as that found in R4NC1 solutions. It has been suggested that in the resin phase HC1 associates with chloride ion to form HC1z-.21 If it is granted that there is a strong similarity between anion-exchange resins and quaternary ammonium salts in aqueous solution, two sets of observations cast doubt on the hypothesis. One is the disappearance of certain infrared bands in mixtures of R4NX-HX or R4NX-HY on exposure to water vapor. The infrared spectra of HC12-, HBrz-, and HClBrions in solidsz4and in solvents25have been reported. HC1, shows two strong bands at 1180 and 1565 cm-l, HBrf has two at 1170 and 1670 crn-l, and HClBr-1 has two at 1100 and 1650 cm-l. These have been assigned to bending and stretching modes, respectively. In the work reported here anhydrous mixtures of Bu4NBr and HBr, Bu4NBr and HC1, and Et4NC1and HC1 were prepared. Infrared spectra of the solid Bu4NBr-HBr showed three broad bands near 900, 1550, and 2000 cm-l, which were not evident in the spectrum of solid Bu4NBr or of gaseous HBr. There were two smaller bands at 1150 and 1670 cm-l. Bu4NBr-HC1 mixtures were liquid and showed three broad bands at 1150,1560, and 2050 cm-l. The Et4NC1-HC1 solid mixture absorbed at 1170 cm-'. All of these bands were observable only with anhydrous products. If the infrared cells were opened so that the mixtures were exposed to air for several minutes, these bands either were diminished or disap-
peared, and the usual water bands appeared. From these experiments we have concluded that the dihalide salts cannot exist in the presence of water and certainly cannot exist in aqueous solution no matter how concentrated. The second set of observations stems from the current research. Mixtures of Pn4NC104 and HC104 form coacervates in which invasion of the R4NX-rich upper layer by HC104 is very marked, quantitatively very much like the invasion by HBr or HC1 of the quaternary ammonium halide-rich layers, which were described earlier. Since the existence of the biperchlorate anion C104-H+C104-is very doubtful, and in fact has not been reported, by analogy the bihalide ions are also not involved. Such a conclusion has already been reached by Diamond and Whitney26for perchloric acid in resins.
References and Notes Present address: Division of Nuclear Medicine, Department of Radiology, Downstate Medical Center, 450 Clarkson Ave., Brooklyn, N.Y. 11203. Taken from the Ph.D. Thesis of L. L-Y. Hwang, Polytechnic Institute of New York, June 1976. M. Lucas, J. Inorg. Nuci. Chem., 32, 3692 (1970); 33, 1883 (1971). R. A. Pierotti, J. Phys. Chem., 69, 281 (1965). M. Lucas, Bull. SOC. Chim. Fr., 2994 (1969). S. I. Shoor and K. E. Gubbins, J . Phys. Chem., 73, 489 (1969). H. S. Harned and 6. B. Owen, "The Physical Chemistry of Electrolyte Solutions", Reinhoid, New York, N.Y., 1958, Chapter 14. M. Lucas and J. Steigman, J. Phys. Chem., 74, 2699 (1970). Reference 7, p 575. K. Schwabe, Sitzungsber. Saechs. Akad. Wiss. Leipzig, Math. Naturwiss. Ki., 109, no. 6 (1972). E. Scarano, G. Gray, and M. Forina, Anal. Chem., 43, 206 (1971). K. Schwabe, Z . Phys. Chem. (Leiprig), 247, 113 (1971). P. Schuster, G. Zundel, and C. Sandorfy, Ed., "The Hydrogen Bond, Recent Developments In Theory and Experiments", North Holland Publishing Co., Amsterdam, 1976, Chapter 15. D. Schioberg and G. Zundel, J. Can. Chem., 54, 2143 (1976). Reference 7, p 467, ref 3. J. Steigman and J. Dobrow, J. Phys. Chem., 72, 3424 (1968). (a) M. M. Marciacq-Rousselot, A. de Trobriand, and M. Lucas, J. Phys. Chem., 76, 1455 (1972); (b) M. Lucas, A. de Trobriand, and M Ceccaldi, ibid., 79, 913 (1975). R. Trioio and A. H. Narten, J. Chem. Phys., 63, 3624 (1975). A. H. Narten and S. Lindenbaum, J. Chem. Phys., 51, 1108 (1969). A. H. Narten, F. Vaslow, and H. A. Levy, J. Chem. Phys., 58, 5017 (1973). S. Lindenbaum and G. E. Boyd, J . Phys. Chem., 66, 1983 (1962). F. Nelson and K. A. Kraus, J . Am. Chem. Soc., 80, 4154 (1958). B Chu and R. M. Diamond, J. Phys. Chem., 66, 1383 (1962). T. C. Waddington, J . Chem. SOC., 1708 (1958). M. E. Peach and T. C. Waddington in "Nonaqueous Solvent Systems", T. C. Waddington, Ed., Academic Press, New York, N.Y., 1965, pp 83-115. R. M. Diamond and D. C. Whitney in "Ion Exchange", Vol. I, J. D. Marinsky, Ed., Marcel Dekker, New York, N.Y., 1966, p 318.
