Thermodynamic properties of tetraalkylammonium ...

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Canadian Journa of Chemistry

Journal canadien de chimie

Pz~blislzedby

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THE NATIONAL RESEARCH COUNCIL O F CANADA

L E CONSEIL NATIONAL DE RECHERCHES DU CANAD4

Volume 54

Number 14

July 15, 1976

Volume 54

numCro 14

15 juillet 1976

Thermodynamic properlties of tetraalkylammonium halides: vcslnmes, heat capacities, and expansibilities in %%,a, D,O and urea-water mixtures from 278 to 328 M 1 G ~ R A LPERROK. D NICOLE DESROSIERS, A N D JACQCES E. DESNOYERS~ Departmetzt of C l i e r n i s f ~Utricersiii , cfe Sherbroohe, Slrerbrooke, P.Q., Crrr7~1rln J I K 2Rl Received February 4, 1976

G ~ R A LPERRON, D NICOLEDESROSIERS, and JACQUES E. DESNOYERS. Can. J. Chem. 54, 2163 (1976). The densities and heat capacities per unit volume of the symmetrical tetraall 3 nz LJ) vary significantly and even change signs in the region 45 to 50 O C . There are at least tho possible explanations for the apparent contradiction between the erithalpies and heat capacities of transfer. For hydrophilic ions, the Frank and Wen model suggests that the heat capacity of transfer should be a linear function of the difference in the molar heat capacities of the two solvents. It is also possible that a relation similar to [S] holds for hydrophobic solutes. Free Energies. Entllab~les,arzcl Entro~~ies of Since CpD - C,'" decreases only slowly in the Trarz~fer The temperature dependence of the standard temperature range 0 to 100 " C ,it is not surprising free energies. enthalpies, and entropies of trans- that ACOBu,NB,(W+D)is relatively constant. On fer can readily be calculated from the literature the other hand, it is quite possible that HD values at 25 OC and the ACg(W + D) and varies significantly with temperature and ACEO(W+ 3 nz U ) at various temperatures. even changes sign at some temperature, thus These functions were calculated for Me4NBr accounting for the change in sign of (Fig. 9) and Bu4NBr (Fig. 10). All the transfer (W D). An alternative explanation could be functions of Me4NBr decrease in magnitude as that Bu4NBr is becoming a structure breaker at the temperature is increased as we would expect high temperatures (66). If this structure-breaking from a decrease in the structure-breaking effect. effect is increasing with temperature, then the The situation is more complicated with Bu4NBr. sign of the heat capacities of transfer can be explained; the heat capacity being a measure of While ACOB,,xB,(W+ D) and AC0Bu4NB,(W 3 m U) change only slightly in magnitude with the change in the energy distribution with temperature, AH0Bl,4\Br(W+D)and AHOBu,KBr- temperature, a positive c 2 - C,(in) could

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TABLE 5. Intrinsic and partial gram ionic volumes mean that an ice-like co-sphere is melting or that (cm3 mol-1) of R4N+ and Br- at 25 "C the structure-breaking effect near a hydrophobic co-sphere is increasing with temperature. Both Ion &(H,o) ~ ( M ~ o H ) "~ ( D M F ) ~ V(in)b cases would be typical of hydrophobic solutes. However, at very high temperatures, structural effects must disappear. Therefore, eventually, all structural hvdration contributions should tend to zero at high enough temperatures. Again this shows the need for accurate data at high temperatures if we wish to understand the nature of aFrom ref. 74. OFrom scaled-particle theory using compressibility data (62) structural interaction in water. The absolute values of AGOB,,xB,(W +3 nz U ) seem to be inconsistent with those of AGOBuqNBr- estimated from conductance data (68), those (W -+ D). The former are significantly negative from compressibilities agree closely with those while the latter are close to zero. Free energies of estimated from To (35). The values of V(in) transfer are difficult to interpret in terms of calculated with the compressibility set of radii structural changes. Therefore, the observed using the equation of Conway et 01. (68) are effects might be real or one of the transfer compared in Table 5 with the ionic volumes of functions taken at 25 "C might be in error. It R4N- in methanol and in dimethylformamide would therefore be useful to remeasure these free given by Kawaizumi and Zana (71) using the energies of transfer before attempting to interpret neth hod of Zana and Yeager (72). The values of To in nonaqueous solvents are these trends. more negative than V(in) for the Br- and the Stanclarcl Volrlmetric Properties lower R4Nt, as expected from some electroWith simple electrolytes like alkali halides striction of these ions. On the other hand, ToN there are various ways of calculating V(in) which V(in) for the higher homologs, again as expected all give similar values (67). Thus, for simplicity, if these large ions are not very solvated in 11011we will adopt the equation of Conway et al. (68) aqueous solvents. Therefore the values of V(in) of R4Nt are probably fairly close to the true ones. The hydration function To - V(in) of the where the second term is a correction for the alkali bromides and R4NBr is shown in Fig. 11. free volume near the ions. The radii used with With the exception of Pen4NBr, these values for alkali halides are those of Waddington (61) but the R4NBr are all less than 5 cm3 mol-l which is other scales would give approximately the the magnitude of the uncertainty on V(in). same V(in). Therefore the best we can say is that To - V(in) The evaluation of V(in) of R4N+ is more of R4NBr is decreasing with cationic size and is difficult. With such large ions, the calculation of probably negative for Pen4NBr. the volume from ionic radii becomes very The standard volumes of transfer of R4NBr sensitive to the choice of the radii and of the are also compared with those of alkali bromides equation. It has been shown (69) that the ionic in Fig. 11. As with heat capacities, A VEo(W radii of alkali halides derived from the scaled- 3 m U ) qualitatively varies with the cationic size particle theory are in good agreement with in the same way as AV$(W + D) and the signs crystallographic data. There are therefore hopes are in general opposite. Allowing for the unthat the radii of R4N+ derived from this approach certainty in the determination of V(in), it seems should also be reasonably reliable. Two sets of that - AVEO(W+ D) varies in a way analogous to radii of R4N- in water have been obtained from V0 - V(in) and AVEO(W+ 3 m U) in an opthis scaled-particle theory. Masterton et al. (70) posite way. As in the case of heat capacities, it have derived a set from a fit of salting coefficients seems that a good part of the interactions giving and Desrosiers and Desnoyers (62) from iso- rise to - V(in) are accentuated in D 2 0 and thermal compressibilities. attenuated in urea-water mixtures. The signs of While the radii calculated from salting co- V0 - V(in) and of the transfer functions are efficients agree reasonably well with those probably the same for MBr and R4NBr. Negative

v0


.

FIG. 12. Dependence of expansibilities 111 water and expansibilities of transfer from H 2 0 to D 2 0 and from FIG. 11. Dependence of volumes of alkali and tetra- H 2 0 to 3 m urea of alkall and tetraalkyla~nmonium alkylammonium bromides o n cationic size a t 25 " C : 0, bromides o n cationic size in water at 25 "C. @, To- V(in) in H 2 0 ; 3, volumes of transfer from H 2 0 to 3 m urea; A, A volumes of transfer from H 2 0 transfer functions are silown in Fig. 3. These to D,O.

functions all vary in a fairly regular way. From these slopes we can derive the standard partial AVEn(W+ D) are also observed with bolaform molal expansibilities En and the corrzsponding electrolytes (49) and with asymmetrical R4NX expansibilities of transfer, AEEO(W+ ID) and (50). The effect therefore seems general. The AE$(W 3 m U). These functions obtained AVEO(W D) of alkali halides are nearly twice indirectly are compared with the directly meaas negative as those of the three first members of sured ones in Fig. 12. In the case of AEEO(W the R4NBr series. If we correct for the presence D) only the indirect values are available. The E0 of Br- assuming that large alkali metal and and AEEO(W4 3 m U) obtained directly and halide ions have about equal contributions to the indirectly are approximately the same. While in transfer functions, it then appears that much of principle direct measurements should yield the transfer volumes of the R4NBr are due to the better values, in actual fact it is not possible to anion contribution; it is therefore only with decide which set of data is the best. Pen4N+ that there seems to be some evidence for The values of E0reflect directly the hydration a solvent isotope effect above the experimental function since V(in) can usually be taken as uncertainty on the transfer functions. Similarly, temperature independent. The posit~ve B0 of with A J'En(W 4 3 m U), a correction for the alkali halides tells us that the ion-solvent interpresence of Br- would make the transfer volumes actions giving rise to the negative pq - V(in) are slightly negative for Et4N+ and Pr4NL and posi- decreasing with temperature. The negative value tive for the higher homologs. In addition, it was of AEhfXn(W+ 3 m U) suggests that a good part observed that A V,-B,oH(W 3 m U) is slightly of VO- V(in) is structural in origin (32). If there negative (42). Therefore, contrary to previous is a shift in equilibrium of a bulky phase of water suggestions (26, 481, it seems that the sign alone to a more close-packed one as the temperature is of A VEn(W D) and of A VEO(W+ 3 m U) increased, then the model of Frank and Wen (1) cannot give much information on hydrophobic as extended by Desrosiers et al. (32) explains a hydration. positive contribution to E0 of the structureThe temperature dependence of Toand of the breaking effect.

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With hydrophobic solutes, E3is also positive: as expected L~EO,,,,~(W -> D) is also positive fgr the higher members and AE0R4hBr(W + 3 M? U) negative. It therefore seems from the large positive E0 of the hydrophobic solutes and the signs of the corresponding transfer functions that the structural contribution to To - V(in) is also negative.

Apparent Mo2rt-1Excess hncttoizs -Any solute-solute interactions in addition to the long-range Debye-Hiickel ones will be reflected in the parameter B y af [I]. It would of course be referable to use a better theoretical basis to interpret the various interactions between ions, such 2s the Friedman approsch (73). However, while the MNC equation has helped to elucidate the nature of solute-solute interactions with excess free energles and enthalpies, it is at present not too successful with higher free energy derivatives like & and E(74). A model (8) has becn proposed previously to interpret the excess functions of electrolytes in term., of co-sphere over!ap. With hydrophobic electrolvtes there are three ~ossiblestructural interac6ons: hydrophobic-hidrophobic,hydrophobic-hydrophilic, and hydrophilic-hydrophilic. In general, the interactions involving hydrophobic ions are larger than the hydrophilichydrophilic ones, and these latter can be neglected. This modei had been applied to free energies but it now seems that manv of the anomalies observed with activity data of aqueous electrolyte solutions arc not entirely structural in origin since they also occur in N-methyiacetamide (75). With excess enshalpies and volumes it appears that the hydrophobic-hydrophobic and hydrophobic-hydrophilic interactions have the same sign: BriQ+) and Bv(-). It was concluded from this that the co-sphere overlap was destructive for hydrophobic electrolytes; whatever happens at infinite dilution is reduced when two soiutes approach each other. Philip and Desnoyers (26) rrieasured Be of R4NBr at 25 "C and found thcm to be negative, which seemed to be co~sistent with a Le5tructive overlap model: c2 - C,(in> is positive at infinite dilution and shou!d become less positlve as the concentratioil increases, However, Leduc and Desnoyers (76) found that PP, of tatrabutylammonium octanoate was very positive in contrndiction with the destructive oberlap model. Also, ail very bydro-

CI-

Br-

I

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FIG. 13. Bc:parameter of tetraalkylammonium halides In H20 at 25 " C .

phobic nol~electrolytessuch as t-butanol(1 I) and piperidine (77) have positive Be. It therefore seems that the hydrophobic-hydrophobic interaction led to a positive Be, in contradiction to the destructive overlap approach. On the other hand, the interaction between a hydrophobic solute and a hydrophilic ion seems to result in a negative Bc. For example ACO,.BUolr(W W + MX) (78) and ACOBU,rBr(%' + W MX) (791, where W stands for water and MX for an alkali halide, are both negative. The parameters Bc of [I] are shown for the tetraalkylan~moniumchlorides, bromides, and iodides in Fig. 13. This parameter has a value close to zero for Me4NX and becomes increasingly negative when going from Et4NX to Pr4NX. For B u ~ N XBc , is less negative than for Pr4NX and Pen4NBr (not shown) has a large positive Be. It therefore appears that with the lower meinbers of these series it is the hydrophobic-hydrophilic interaction which predominates but with the higher meinbers the hydrophobic-hydrophobic ones take over. The nature of the anion dots not seem to have a inarked effect on Bc as seen from Fig. 13. The temperature dependence of Be for &WBr is shown in Fig. 14. With Me4NBr, Bc is close to zero and varies only slightly with temperature, with Et4NBr it is negative but becomes less so as the temperature increases. With the two most hydrophobic electrolytes there is an inversion between the order of Bc at low and high temperatures. At high temperatures Be decreases in the order Me4NBr to Pen4NBr suggesting a predominance of hydrophobic-hydrophilic interactions. At low temperatures Bc goes through a minimum with Pr4NBr. The same kind of inversion of Bc with temperature was observed

+


FIG. 15. Temperature dependence of the Bv parameter in H20 and of the difference in D20 and H20 and in 3 nl urea and H20 of tetraalkyfammonium bromides.

FIG.14. Temperature dependence of the B, parameter in H20 and of the difference in H 2 0 and in 3 m urea and M20 of tetraalkylamrnonium bromides.

with ionic surfactants (80). It therefore appears that with hydrophobic interactions, the excess hcat capacities suggest a cooperative effect. Whatever happens at infinite dilution happens even more so when two hydrophobic solutes approach each other without associating. When they do associate (contact pairs) then there is a large decrease in heat capacity. The difference in Bc between H20. DzO, and urea-water mixtures is also shown in Fig. 14. Surprisingly enough these ABc are all very small and possibly with the exception of Pen4NBr are, within the experimental uncertainty, approximately equal to zero. This had also been noted by Philip and Desnoyers (26) for Et4NBr and Pr4WBr. They had seen a solvent isotope effect with Bu4NBr because the +c data in D 2 0 was based on erroneous rpv data. With the new +v data, there is hardly any solvent isotope effect on Bc. The signs of A&(W + 3 m U) and ABc(W + D) are both the same for Pen4NBr and the sign of ABc(W i.D) is opposite that of Bc(W). which seems to contradict the trends in the properties at infinite dilution. The possibility that the sign of ABc could be due to an experimental

error cannot be excluded since in view of the low solubility of Pen4NBr (0.2 mol kg-') the uncerrainty on Bc at low and high temperatures can be relatively high. It therefore seems that even though Bc of Pen4NBr is very large, it hardly changes from one aqueous solvel~tto another. An explanation for the consistency of 3, can be offered for the urea-water mixtures. At low R4NBr and urea conceratrations where only pair interactions occur, the R4N+-R4N' interactions would be independent of urea. The effect of urea would show only when triplet interactions of the type E-E-U would become important. However, this explanation does not hold for the solvent isotope effect.

Apparent Mold Excess Volume~ricProperdties The parameters Bv and changes ABV(W+ D) and ABy(W -. 3 r n U) are shown in Fig. 15. All BITare negative and for the hydrophobic solutes they are more negative at low temperatures. This is consistent with the hydrophobic-hydrophobic and hydrophilic-hydrophilic interactions being both negative. These data are not inconsistent with a cooperative overlap model. The change in voltarne To - V(in) is probably slightly negative and becomes more negative as two solutes approach each other. When they associate (con-


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tact pair) there is then a destructive overlap effect and a large increase in volume occurs as in micelle formation (80). Again it is not easy to see definite trends in ABv(W + D) and ABv(W 4 3 m U). It seems, at low temperatures, that ABv are in the direction expected (Bv is more negative in D20 and less so in 3 m urea), but the trends are comparable to the experimental uncertainty. The slopes of +E - AE(dom)1/2versus m for the R4NBr in Fig. 5 are also consistent with a cooperative overlap effect. At infinite dilution $ J ~ O is positive and +E becomes more so as the concentration increases (BE is positive) with the higher members of the series. It therefore seems that when two hydrophobic solutes approach each other there is first of all some stabilization of the co-spheres. This could occur for example at a separation of one or two water molecules as is usually observed with solid clathrates. At shorter distances there is some collapse of the water structure: two solutes together will perturb the water structure less than two isolated solutes. Both phenomena would be favorable, i.e. lead to a negative excess free energy. The only exception to this rule seems to be excess enthalpies. While it seems that the hydrophobic hydration contribution to AHhOis negative, the experimental excess enthalpies are positive whether the solutes are in contact or not. We can offer no explanation for this at the present time. Acknowledgments We would like to thank Dr. C. de Visser for his useful comments on this manuscript and for his collaboration in repeating the heat capacity measurements of MeoWBr in water. We are also grateful to the National Research Council of Canada for financial assistance and for the award of a scholarship to one of us (N.D.). 1. H. S. FRANK and W.-Y. WEN.Discuss. Faraday Soc. 24, 133 (1957). 2. W.-Y. WEN.111 Water and aqueous solutions. Edited by R. A. Horne. Wiley, N.Y. 1972. Chapt. 15. 3. W.-Y. WEN.J. Solution Chem. 2, 253 (1973). 4. T. S. SARMA and J. C. AHLUWALIA. Chem. Soc. Rev. 2, 203 (1973). 5. M. L u c ~ and s A. FEILLOLAY. J. Phys. Chem. 75,2330 (1971). J. Solution Chem. 4, 6. P. R.PHILIP and C. JOLICOEUR. 105 (1975). Water and aqueous solutions. Plenum, 7. A. BEN-NAIM. N.Y. 1974.

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