Transfer Energies and Solute-Solvent Effects in the Dissociation of Protonated ''Tris" in NMethylpropionamideWater Solvents Roger G. Bates, James S. Falcone, Jr., and Anthony Y. W.
Ho
Deparfment of Chemistty, University of Florida, Gainesville, Fla. 326 1 1
A recent Investigation showed that the pK of protonated tris( hydroxymethy1)aminomethane (Tris) in N-methylpropionamlde (NMP), a solvent of high dielectric constant, is greater than that in water. In order to understand the nature of the solute-solvent interactions involved, solubility and emf measurements have been used to determine the individual transfer energies of Tris, Tris-hydrochloride, and hydrochloric acid between water and NMP-water mixtures and to derive the pK at 25 OC from them. The enthalpy of dissociation has also been determined by calorimetry. The pK decreases when NMP is first added, passes through a minimum at a mole fraction of NMP near 0.4, and then rises sharply. This same pattern is followed by the pK in methanol-water mixtures, despite the fact that NMP raises the dielectric constant of the solvent while methanol lowers it. Comparison of the transfer energies for the individual species shows that stabilization of the proton is the primary cause of the enhancement in acidic strength of Tris-Hf when NMP is added to a predominantly aqueous solvent.
The properties of both weak and strong electrolytes in water have been studied in detail, but the behavior of electrolytes in nonaqueous solvents is still not well understood. Most of the common nonaqueous liquid media have dielectric constants less than that of water. Although electrostatic forces have long been thought to play a major role in electrolyte behavior, very little is known about ionic processes in solvents with dielectric constant greater than that of water. One such medium is N-methylpropionamide (NMP), the dielectric constant of which is 176 a t 25 "C. In recent studies, the thermodynamics of the dissociation of acids of two charge types, namely acetic acid ( I ) and protonated Tris(hydroxymethy1)aminomethane (Tris) ( 2 ) , in NMP has been studied. Despite the high dielectric constant, both of these acids were weaker in NMP than in water. In the case of Tris-Hf, the dissociation is an isoelectric process and should be insensitive to changes of dielectric constant. In earlier work ( 3 ) ,it was nonetheless found that the pK of Tris-H+ falls when methanol is added to the aqueous solvent and passes through a minimum when the solvent contains about 60% (w/w) of methanol. It was therefore of interest to learn whether NMP (which, in contrast with methanol, increases the dielectric constant of the solvent mixture) causes similar or contrasting changes in the dissociation constant. An attempt has now been made to elucidate the patterns of solute-solvent interaction by studying the change in the pK of Tris-Hf as the composition of the solvent is shifted gradually from pure water to pure NMP. The properties of ( 1 ) E. S. Etz, R. A. Robinson, and R. G. Bates, J. Solution Chem., 1, 507 (1972). (2) E. S. Etz, R. A. Robinson, and R. G. Bates, J. Solution Chem., 2, 405 (1973). (3) P. Schindier, R. A. Robinson, and R. G. Bates, J. Res. Naf.Bur. Stand., Sect. A, 72, 141 (1968).
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NMP-water solvents at 25 "C are summarized in Table I. The densities and dielectric constants were derived from the work of Hoover (4, 5 ) . The Debye-Huckel constants were calculated from the density, dielectric constant, and thermodynamic temperature by the usual formulas (6). The dissociation process can be formulated simply as
Tris -H+(s) k Tris(s)
+
H+(s)
(1) It is evident that the change of pK with the composition of the solvent(s) is the resultant of the relative stabilizations of three species and cannot of itself indicate the dominant interaction bringing about a shift in the equilibrium constant for process 1. For this reason, the pK in NMP-water solvents was derived from the standard changes of Gibbs energy on the transfer of Tris-hydrochloride, Tris, and hydrochloric acid from water to NMP-water solvents, applying a procedure suggested by the work of Kolthoff, Lingane, and Larson (7). If the Gibbs energy change in the transfer of a species i from water (w) to a nonaqueous or mixed solvent (s),name1Y
i(w) = i ( s )
(2)
is designated AGto (i), it is evident that AG,"(diss)
= (RT In 10)(P,K
AG,"(Tris)
+
Act"(")
-
-
P,K) = A G t o ( T r i s -H+) (3)
Since the transfer energies of the ions are additive, AG,"(diss) = h G , " ( T r i s )
i
AG,"(HCl)
-
AG,"(Tris-HCl) (4)
We have determined the transfer energies of the Tris base and its hydrochloride from measurements of the solubility of these substances in water and in NMP-water solvents. The transfer energy for HCl was calculated from emf measurements of cells without liquid junction. By combination of the results according to Equations 3 and 4, the pK (that is, p,K) in NMP-water mixtures has been obtained. As in methanol-water solvents, it decreased when NMP was first added to the aqueous solvent and passed through a minimum at a mole fraction of NMP of about 0.4. The enthalpy of dissociation was also determined by calorimetry and entropies for the dissociation process were calculated. From a consideration of the thermodynamics of the transfer process, it appears likely that the minimum in pK results from a stronger stabilization of the proton in NMP-water mixtures than in water alone, combined with a preference of Tris and Tris-Hf for water over NMP. (4) T. B. Hoover, J. Phys. Chem., 73, 57 (1969). (5)T. B. Hoover, Pure Appl. Chem., to be published. (6) R . G. Bates, "Determination of pH," 2nd ed., John Wiley and Sons, New York, N.Y.,1973, pp 248, 249. (7) I. M. Kolthoff. J. J. Lingane, and W. D. Larson, J. Amer. Chem. Soc., 60, 2512 (1938).
ANALYTICAL CHEMISTRY, VOL. 46, NO. 13, NOVEMBER 1974
Table I. Properties of N-Methylpropionamide-Water Solvents a t 25 "C Wt % N M P
0 20 40 60 80 100 a
Debye-Hdckel Constants'
___~
Mole fraction, NMP
Mean mol wt
Dielectric constant
Density/ g ml-1
A
€3
0 0.0492 0.1212 0.2368 0.4527
18.02 21.41 26.39 34.38 49.30 87.12
78.33 80 .o 80.8 83.0 90.3 176
0.9971 0 ,9934 0,9960 0.9873 0.9626 0.9310
0.5108 0.4935 0.4869 0,4657 0.4051 0.1465
0.3286 0.3245 0.3234 0.3176 0.3006 0.2118
1
Scale of molality. ~~
~~
Table 11. S t a n d a r d Gibbs Energy Change AGto for the Transfer of Tris(hydroxymethy1)aminomethane (E) f r o m Water t o NMP-Water Solvents Derived f r o m Solubility Measurements at 25 "C Solubility of B
~ _ _ __ Wt
cc. N M P
0 20 40 60 80 100 a
mol kg-1
5 ?80b 4 168 2 667 1 372 0 482 0 O6Oc
Mole fraction scale
_
~
mole fraction
0 0 0 0 0 0
09431 08194 06574 04504 02322 00520
/H
1 1 1 1 1 1
110 089 070 047 024 005
Table 111. S t a n d a r d Gibbs Energy Change AGt" for
t h e Transfer of Tris(hydroxymethy1) a m i n o m e t h a n e Hydrochloride (BHCI) f r o m Water t o NMPWater Solvents Derived f r o m Solubility Measurements at 25 "C
A c t " (B) cal mol ' l
Solubility of BHC1
0
94 235 472 878 1776
' Reference (3) ' Reference (2).
Wt yc N M P
mol kg-1
mole fraction
f,
0 20 40 60 80 100
6.750b 5.034 3.463 1.966 0.913 0.352b
0,10842 0.09731 0.08372 0.06330 0,04308 0.02975
0.625 0.636 0.643 0.663 0.722 0.930
Mole fraction scale.
EXPERIMENTAL Tris(hydroxymethy1)aminomethane (Tris), NBS Standard Reference Material, was dried a t room temperature under vacuum in a desiccator. A commercial lot of Tris-hydrochloride was used without further purification. Constant-boiling hydrochloric acid, twice distilled, was prepared. N-Methylpropionamide, obtained from a commercial source, was fractionally distilled a t 65 "C under a pressure of about 0.5 mm Hg in a Todd still, the initial and final quarters of the distillate being rejected. Gas chromatography revealed traces of water, estimated at about 0.03% (w/w), but disclosed no propionic acid or methylamine. Solubility Measurements. An excess of solid was added to solvent contained in bottles of 125-ml capacity. The bottles were rotated end-over-end in a constant-temperature bath for a period of 72 hours or more. Aliquots of the saturated solutions were withdrawn, weighed, and analyzed a t daily intervals. The solutions were assumed to be in equilibrium with the solid phase if no change in concentration was found over a period of 24 hours. Tris was determined by titration with a standard solution of hydrochloric acid with Methyl Red as indicator. The saturated solutions were diluted with water before titration. Solutions of Trishydrochloride were diluted with water, acidified with dilute nitric acid, a slight excess of silver nitrate solution was added, and the weight of silver chloride determined by the usual gravimetric procedure. Emf Measurements. The cells were of the all-glass type in general use in this laboratory (8),incorporating a three-stage saturator for the hydrogen gas. The platinum electrodes were coated with platinum black by the procedure described elsewhere (9). The silver-silver chloride electrodes were of the thermal-electrolytic type. Commercial hydrogen gas was purified by passage over a platinum De-Oxo catalyst. The emf was measured by a Hewlett-Packard digital voltmeter which was checked frequently against two calibrated saturated Weston cells maintained a t a controlled temperature. Enthalpy Measurements. The solutions used for the calorimetric measurements were prepared by adding weighed amounts of water to analyzed stock solutions of Tris in NMP and HCl in NMP. The latter had been prepared some months earlier but had been preserved in a carefully sealed brown bottle. Its stability was confirmed by the acceptable agreement between the enthalpy of (8)R. Gary, R . G. Bates, and R. A. Robinson, J. Pbys. Cbern., 68, 1186 (1964).
Bates, "Determination of pH." 2nd ed., John Wiley and Sons, New York, N.Y.. 1973, Chap. IO.
(9) R . G.
AG,'(BHCi)/ cal mol --I,&
0 107 '273 568 923 1061
Reference (2).
ionization obtained calorimetrically and the valuederived from emf measurements (2) when the solvent was substantially pure NMP. T o minimize the effects of acid-catalyzed hydrolysis of the NME , the measurements with HCl solutions in NMP-water solvents were made within 8 hours after the acid solutions were prepared. The batch microcalorimeter described by Wadso (IO)and manufactured by LKB Instruments, Inc., was used. Experimental techniques were similar to those of Falcone, Levine, and Wood (11). The changes in voltage of the thermopiles were amplified and recorded as a function of time, and heats of the chemical reactions were calculated by comparing the areas with those produced by known amounts of added electrical energy. The areas were measured with a planimeter, and the calculated amounts of chemical energy appeared to have an uncertainty less than 0.2%.
RESULTS Solubility Measurements. The standard changes of Gibbs energy A c t o for the transfer processes
Tris (in H,O) = Tris (in NMP-H,O)
(5)
and T r i s -HC1 (in H,O) = Tris -HCL (in NMP-HzO) (6) were determined by measuring the solubilities of Tris base and its hydrochloride in NMP-water solvents and combining the data with the solubilities in water measured in earlier work (2, 3 ) . T o permit valid comparisons for equal numbers of solvent molecules, the mole fraction scale was used. The data are summarized in Tables I1 and 111. The calculation of A G t ' a t 25 "C (in cal mol-I) was made by the equation
where Y = 1 when i is Tris and 2 when i is Tris-HC1; the subscripts w and s refer to water and to NMP-water solvents, respectively. (10) I. Wadso. Acta Cbern. Scand., 22, 927 (1968). (11) J. S. Falcone, Jr., A. S. Levine, and R. H. Wood, J. Pbys. Cbern., 77, 2137 (1973).
ANALYTICAL CHEMISTRY, VOL. 46, NO. 13, NOVEMBER 1974
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Table IV. S t a n d a r d Gibbs Energy Change AGtO for t h e Transfer of HCl f r o m Water to NMP-Water Solvents Derived f r o m E m f Measurements of Cell A a t 25 "C W t ci; N M P
Molality of HCl/mol kg-1
0 20 40 60 80 100
0.01 0.002 0.002 0.01 0.01 0.01
Mole fraction scale.
E'mIV
YL
0 .46412b 0.5512 0.5574 0.4808 0.4754 0 . 40533c
0.904 0.954 0.954 0.912 0.922 0.970
EOZ/V
0,22234 0.2294 0,2357 0,2395 0.2346 0.16724
0.01597 0.0319 0.0489 0.0663 0,0800 0.04184
AGt"(HCI)/ cnl mol-'"
0 - 367 - 759 - 1161 - 1477 - 597
Reference (14). Reference (15).
I
2
E/V
I
I
I
-
at 25 O C was derived from the standard emf of the cell without liquid junction
P t ; H z ( g , 1 atm). HCl(m), AgC1;Ag by the equation AG,"(HCl)/cal mol-' = 23061(,Ec - ,E") (11) The standard emf on the scale of molality was first calculated by the Nernst equation in the form
E"
0
0.2
0.6
1
XNMP
Figure 1. Standard transfer energies (mole fracfion scale) for Tris, Tris-hydrochloride, and HCI from water to NMP-water solvents at 25 OC
The rational activity coefficients f i have not been measured in the mixed solvents. In water, however, yi (scale of molality) has been found to be 1.005 for Tris and 0.503 for Tris-hydrochloride by isopiestic vapor pressure measurements of the saturated solutions (12). Inasmuch as the solubility of Tris is markedly lowered by the addition of NMP to the aqueous solvent, yi for this substance was taken to be 1 in the NMP-water solvents. The mean ionic activity coefficient y* (scale of molality) for Tris-HC1 was calculated by the Debye-Huckel equation
in which m is the molality of the saturated solution and 6 is the ion-size parameter. The Debye-Huckel constants A and B (scale of molality) for ii values in 8, are given in Table I. When ii is set equal to 4.03 A, Equation 8 yields the known value (0.503) of yh in pure water; hence, this same value of ii was used to calculate yi for the mixed solvents as well. Activity coefficients yi on the scale of molality were converted to rational activity coefficients fi by the customary formula (13) .f, = ~ ~+ (O.O01~,ll'~>7) 1 (9) where M' is the mean molecular mass of the solvent and v is the number of moles of ions produced from one mole of the electrolyte. Emf Measurements. The standard change of Gibbs energy for the transfer process
HC1 (in H,O) = HC1 (in N M P - H 2 0 )
(10)
(12) R. A. Robinson and V. E. Bower, J. Chern. Eng. Data, 10, 246 (1965). (13) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," 2nd ed. revised, Butterworths, London, 1970, Chap. 2.
2006
= E
4
0.11831 log
(,>lyi)
(12)
and then converted to the mole fraction scale by addition of 0.11831 log (0.001M'). The results are summarized in Table IV. The standard emf of the cell in water ( 1 4 ) and in NMP (15) have already been determined. The accuracy of measurements in NMP-water mixtures is impaired by the acidcatalyzed hydrolysis of NMP, particularly marked in 20 and 40 wt % NMP. At these compositions, the molality of HCl was lowered to 0.002 mol kg-' in order to retard hydrolysis but was maintained a t 0.01 mol kg-l in 60 and 80 wt % NMP. After equilibration of the cells, the emf increased a t the rate of 0.1 to 0.3 mV per hour, presumably because of hydrolysis of the amide. For this reason, the measurements of emf were made in as short a time as possible after the addition of HCl to the mixed solvent, and the observed values of E were extrapolated back to the time of mixing. The mean molal activity coefficient of hydrochloric acid ( m = 0.01 mol kg-l) is 0.904 in water (14) and 0.977 in NMP ( 1 5 ) . In the mixed solvents it was calculated by Equation 8 with ii = 5 8,; this procedure yields values of 0.904 and 0.970, respectively, for the pure solvents water and NMP. The Gibbs energies for the transfer of Tris, Tris-HC1, and HCl (mole fraction scale) are plotted in Figure 1 as a function of the mole fraction of NMP in the solvent. Calorimetry. The heat of neutralization of Tris base with hydrochloric acid was determined by the reaction of dilute solutions of the two substances in a microcalorimeter. The molality of the Tris solutions fell in the range 0.027 to 0.050 mol kg-l and that of the acid solutions varied from 0.041 to 0.076 mol kg-1. Weights of the solution samples were from 2.2 to 3.5 g, with the moles of acid always in slight excess over the moles of base. The results are summarized in Table V and compared with earlier measurements, by emf and calorimetry, of the enthalpy of dissociation in pure water and "pure" NMP [which in this investigation and the earlier emf study (2) contained 0.02 to 0.04 wt % of water]. The correction of the experimental molar enthalpy to infinite dilution could not be made with certainty in the absence of heats of dilution for the acid, base, and salt in (14) R. G. Bates and V. E. Bower, J. Res. Nat. Bur. Stand., 53, 283 (1954) (15) E. S.Etz, R. G. Bates, and D. Rosenthal, to be published.
ANALYTICAL CHEMISTRY, VOL. 46, NO. 13, NOVEMBER 1974
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Table V. E n t h a l p y of Dissociation of P r o t o n a t e d Tris(hydroxymethy1)aminomethane ( B H + )in the NMP-Water Solvent System Wt 70
Solvent,
NMP
XNMP
AHa/kea1 mol-'
- q/cal
10.92"
54.20
0.197
83.10
0.504
93.55
0.750
0.10151 0.10816 0.15280 0.22119 0.15680 0.16741 0,22005 0.17654
0,09163 0.09595 0.12939 0.09791 0,13286 0.15330 0.12793 0.14198
1
100
AH3/kcal mol-' fcorr.)
(uncorr.)
0
0
a
HCI, mmol
Tris, mmol
11. 3 8 b 11. 3 3 c 10.57 f 0.02
10.63 10.59 11..30 10.91 12.72 12.43 15.51 15.76
0,9740 1.0164 1,4624 1,0684 1.6897 1,9057 1,9842 2.2376
11.03 i 0.22 12.52 i 0.14 1 5 . 5 6 j=0 . 1 4 1 5 . 63d
Sturtevant, calorimetry (16). Bates and Hetzer, emf ( 1 7 ) . Datta, Grzybowski, and Weston, emf (18). d Etz, Robinson, and Bates, emf ( 2 ) .
Table VI. S t a n d a r d Thermodynamic Q u a n t i t i e s and pK for the Dissociation Process Tris-H+ % Tris H + in NMP-Water Solvents at 25 "C (Scale of Molality)
+
AGC
Wt Yc N M P
0 20 40 60 80 100 a
kcal mol-'
11 10 10 10 10 12
Reference (17)
02 74 45 14 09 05
AH"/ kcal mol-'
11 11 10 10 10 15
>SO
cal K-1 mol-'
72
11 0 9 0 9
54 85 63
2 5 12 0
35 00
1 3
pK
8 7 7 7 7 8
075a 87 66 44 40 831h
Reference ( 2 ) .
NMP-water solvents. Nevertheless, correction terms were estimated and applied, even though the correction was smaller than the uncertainty of the measurements. The heat of dilution for Tris base was assumed to be negligible, and the heats of dilution of the electrolytes were estimated from the Debye-Huckel slopes calculated with the aid of the densities and dielectric constants of NMP-water mixtures given in Table I. The total correction amounted to 70 cal mol-l or less at all solvent compositions. The hydrolysis of NMP is catalyzed by acid, and it was necessary to avoid long delays between the dilution of the stock solution of HCl in NMP with water and the measurement of the heat of reaction. In one experiment, the measurements a t XNMP = 0.504 were repeated with the same base and acid solutions one week after the solutions were prepared and the data reported in Table V had been obtained. The mean AH (uncorrected) was 9925 cal mol-l. Thus it appears that the measured heat of neutralization is reduced as hydrolysis of the amide proceeds and that the rate of decrease is about 170 cal mol-I per day. The results given in the Table were obtained within 6 to 8 hours after preparation of the acid solutions in NMP-water solvents, and therefore an uncertainty of some 60 cal mol-' can be expected from this source.
DISCUSSION It is now possible to determine AGto (diss) by Equation 4 and, inasmuch as pK in water is known, to derive pK in NMP-water solvents from it by Equation 3. The results are given in Table VI, and p K on the mole fraction scale is plotted as a function of the mole fraction of NMP in Figure 2. (16)J. M. Sturtevant, J. Amer. Chern. SOC.,77, 1495 (1955). (17) R. G.Bates and H. B. Hetzer, J. Phys. Chem., 65, 667 (1961). (18) S. P. Datta, A. K. Grzybowski, and 6. A. Weston, J. Chem. Soc.,London, 792 (1963).
0
0.2
0.4 X
0.6
0.8
7
NMP
Figure 2. Change of pK(molality scale) for protonated Tris with solvent composition in the NMP-water solvent system
With the aid of AHo from Table V, one can also obtain the standard entropy for the dissociation process. The thermodynamic functions AG," (diss), Uto(diss), and ASto (diss) for the dissociation process (mole fraction scale) may be a useful indication of the relative stabilizations of the acid-base species as solvent water is replaced by solvent NMP. These are the thermodynamic constants for the transfer process
Tris(w)
-t
H'(w) + Tris-H'(s) =
Tris(s) + H'(s) iT r i s - H ' ( w ) (13) The values plotted in Figure 3 were obtained by conversion of the results given in Table VI to the mole fraction scale. It is evident that Tris-H+ increases in acidic strength when NMP is added to a solvent that is initially pure water. The pK value passes through a minimum when the mole fraction of NMP in the solvent is about 0.4 and subsequently rises, reaching a value in pure NMP that is some 0.8 unit higher than that in pure water. The effect of added NMP on p K is indeed very similar to that produced by addition of methanol, where pK falls to a minimum of about 7.8 and rises again sharply when the mole fraction of water in the solvent becomes low. Furthermore, the course of the curve of pK us. x 2 resembles closely that for the pK of m-
ANALYTICAL CHEMISTRY, VOL. 46, NO. 13, NOVEMBER 1974
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Table VII. Comparison of t h e Transfer Energies of the Tris-HT Cation and t h e Proton from Water to NMP-Water and Methanol-Water Solvents I G t a ( T r i s H-, - I G t - ( H A )
0
0.2
0.6
1
XNM*
Figure 3. Thermodynamic functions AG" (diss), Aho (diss), and TAS," (diss)for the transfer dissociation of protonated Tris (Equation 13)in NMP-water solvents (mole fraction scale)
nitroanilinium ion in mixtures of water and dimethylsulfoxide (DMSO) (19, 20), which, like NMP, is an aprotic acceptor solvent. These similarities of behavior are a t first thought surprising, in view of the fact that addition of methanol or DMSO to water produces a series of solvents of progressively decreasing dielectric constant, while addition of NMP forms a series of mixtures of increasing dielectric constant. The plots of the individual transfer energies in Figure 1 make it clear that both Tris and Tris-hydrochloride are more strongly stabilized by water than by NMP. On the other hand, hydrochloric acid appears to be stabilized by NMP so strongly that addition of NMP to the aqueous solvent promotes dissociation of the Tris-H+ when the mole fraction of NMP is 0.4 or lower. By contrast, the transfer energy for HCl in water-methanol solvents is always positive (preferential stabilization by water) and rises sharply in methanol-rich mixtures. Here the decrease of pK in predominantly aqueous solvent mixtures can be attributed to the fact that AGto for the Tris-hydrochloride exceeds that for the Tris base in a region where A c t o for HCl is small. In mixtures containing large amounts of methanol, the stabilization of HCl by water is dramatically lessened, whereupon pK reaches a minimum and rises once again. It is tempting to conclude immediately from these results that NMP is a more basic solvent than methanol, and even more basic than water. Nevertheless, the energy change in the transfer of an acidic species from one solvent to another is not an unequivocal index of relative basicity, (19)R. K . Wolford, J. Phys. Chem., 68,3392 (1964). (20)R. G. Bates, L. Johnson, and R. A. Robinson, Chem. Anal. (Warsaw), 17,479 (1972).
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Mole fraction, N M P or MfOH
NMP-H 0, cal mol-'
MeOH-H 0, cal mol-'
0.2 0.4 0.6 0.8
1550 2280 2450 2200
450 950 1120 960
for several reasons. If the first solvent is water, addition of the second solvent alters the structure of the water and the patterns of solute-solvent interaction in a manner not well understood. Furthermore, the process for which SG," was determined here includes the transfer of an anionic (chloride) species as well as the proton, and it well may be that halide ions are more strongly stabilized by one solvent component than by the other. Thus, the decrease in Acto for Tris-HC1 at high concentrations of NMP may mean that chloride ion "prefers" NMP to water. A third factor that must be taken into account is the increased stabilization that is to be expected from simple Born electrostatic considerations as the dielectric constant of the medium increases. Inasmuch as standard transfer energies are additive, it is possible to eliminate the effect of preferential solvation of the anion. Thus, the data given in Tables I11 and IV permit one to evaluate the difference between the transfer energies of the Tris-H+ cation and the proton. A comparison of the results with those in methanol-water mixtures ( 3 ) is shown in Table VII. The differences in transfer energy found for the two solvent systems follow similar paths with changing composition of the media, but those in NMP-water are greater than those in methanol-water by 1.1 to 1.3 kcal mol-'. I t appears that the dielectric constant has little influence here, as there is a difference of about 100 between the dielectric constants of the two solvent systems a t x = 0.8 but only about 15 a t r = 0.2. These results thus provide further evidence of the primary importance of preferential solvation in determining the behavior of ionic processes in solution.
ACKNOWLEDGMENT The authors are grateful to R. A. Robinson for helpful suggestions. RECEIVEDfor review April 2, 1974. Accepted June 13, 1974. This work was supported in part by the National Science Foundation under Grant GP14538. Presented at the I. M. Kolthoff 80th Anniversary Symposium, Division of Analytical Chemistry, 167th National Meeting, ACS, Los Angeles, Calif., April 1974.
ANALYTICAL CHEMISTRY, VOL. 46, NO. 13, NOVEMBER 1974
