Enthalpies of Urea-Water Solutions
1391
Thermodynamic Quantities for the Transfer of Urea from Water to Aqueous Electrolyte Solutions Martha Y. Schrler, Peter J. Turner, and Eugene E. Schrier" Department of Chemistry, State University of New York at Binghamton, Binghamton, New York 13901 (Received January 29, 1975) Publication costs assisted by the State University of New York
Enthalpies of solution of urea ( c ) were determined a t 25O in water and in solutions of electrolytes, mainly alkali and alkaline earth halides, over a range of salt molalities at urea molalities below 0.03 m. Slopes of the plots of enthalpy of transfer of urea vs. salt molality were obtained at infinite dilution of both solutes. Gibbs free energies of transfer of urea from water to solutions containing one of the salts, LiC1, NaC1, KCl, or CsCl, were obtained by potentiometric and Ssopiestic measurements. Entropies of transfer were calculated. Enthalpy-entropy compensation is observed with a compensation temperature of T* = 280°K. An interpretation of the trends in the data is given based on the Friedman cosphere model. No conclusions can be drawn from data a t low molalities regarding the effectiveness of amide group-salt interaction in the mechanism of isothermal protein denaturation.
Introduction Urea-water solutions have been the subject of intense experimentall and theoretical investigation.2 The polarity of urea, its hydrogen-bonding ability, and its consequent high solubility in water provides intrinsic interest. The protein-denaturing ability of urea in aqueous solutions has also attracted attention. Thermodynamic quantities for the transfer of solutes from water to urea-water mixtures have been determined recently in the hope of elucidating the structural features of the mixed solvent which cause it to differ from water. Studies involving nonpolar solutes3 have suggested that the addition of urea to water destabilizes water structure leading to the disruption of hydrophobic bonds in proteins. Recently, Ben-Naim and Yaacobi4 have reached a different conclusion regarding hydrophobic interactions based on their studies of hydrocarbon solubilities in aqueous urea solutions. They suggested that urea strengthens the hydrophobic interaction rather than weakens it. Quantities for the transfer of electrolytes from water to water-urea mixtures have been obtained by a number of investigators, among them Wen and his collaborator^,^ Stern and Kulluk,6 Desrosiers et al.,7 and Cassel and Wood.8 The results of these workers again suggested the reduction of structure in water-urea mixtures as compared to pure water. The work to be described here was initiated with a somewhat different point of view than that of previous studies. We required a compound to use in modeling the interaction between an ion and an amide linkage in aqueous solution. This interaction appears to be significant in the mechanism of the denaturation of proteins by salts.g Recently, enthalpies of transfer of formamide from water to salt solutions were obtainedlO over a limited range of salt molalities. The greater stability of urea makes it better suited than formamide to be a model for the amide linkage in aqueous salt solutions. The program which we have established for this work calls initially for the development of thermodynamic information which can be accurately extrapolated to limiting molalities of both salt and urea. In this way, parameters are obtained for the interaction of a urea molecule with a particular pair of ions in the aqueous medium. The availability
of these parameters allows us to assess the various factors which influence this interaction. The next phase of the work, more pertinent to protein denaturation, will be directed to solute-solute interaction at higher solute concentrations.
Experimental Section Materials. Urea (Fisher) was twice recrystallized from anhydrous methanol and dried under vacuum for 72 hr a t room temperature. Tris(hydroxymethy1)aminomethane obtained from Sigma Chemical Co. as Trizma Base was dried a t 80" under vacuum. Tetramethylammonium bromide (Eastman), recrystallized twice from a 3:l methanol-ethano1 mixture, was dried under vacuum at 60". The remaining salts were reagent grade and were dried overnight at 110" before use. Concentrated stock solutions were prepared of hydrated salts and of other salts which were difficult to weigh directly in air. These stock solutions were passed through a 1.2-k Millipore filter and analyzed for halide using AgN03 in a potentiometric method. All solutions were made up with glass distilled water on a molality basis, i.e., moles of solute per kilogram of water as the solvent. Calorimetry. Enthalpies of solution were measured using an LKB Model 8700-1 precision calorimeter in conjunction with a potentiometric recorder. Electrical calibrations were performed before and after the solution experiments. All calibrations and solution reactions were carried out a t equal intervals on either side of the mean temperature, 25.000 f 0.005O. The corrected resistance change was determined by linear extrapolation of the before and after drift curves to that time which corresponded to 63% of the heat absorption for the solution experiments and to 50% of the heat evolution for the calibrations. Potentiometry. The technique for obtaining the data has been described.llJ2 In general, potential differences were measured between two cells of the form Ad AgC4 MC1,urea;HzOl cation sensitive glass electrode
where the glass electrode is common to each cell. Each cell contained the same salt molality but different urea molalities. Monovalent cation electrodes (Beckman no. 39137 and Corning no. 476220) were used for the LiC1-urea and The Journal of Physical Chemistry, Vol. 79, No. 14, 1975
M. Y. Schrier, P. J. Turner, and E. E. Schrier
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KC1-urea solutions while a Corning sodium electrode (no. 476210) and the Beckman cation electrode were used to
measure potentials for the NaC1-urea system. An Electronics Instruments Limited (EIL) ammonium-potassium electrode was used in CsC1-urea mixtures. Silver chloride reference electrodes were prepared from Beckman silver billet electrodes by electrolysis in 0.1 M HC1 for 3 min with a 15-mA current. Cell potentials were read to f0.02 mV with a Keithley 630 potentiometric electrometer. All runs were carried out at 25.0 f 0.1' under a nitrogen atmosphere. The pH range of optimum electrode response to each salt was determined potentiometrically by titrating a 0.1 M salt solution (pH 11-12) with 0.1 M HC1. Because of the limited p H range in which the electrode responded to metal ion for CsC1-urea solutions and because of low electrode sensitivity to the metal ion in this system and in LiC1-urea, the apparent slow hydrolysis of urea presented more difficulty in these systems than in the others. Frequent adjustment of the solution pH was necessary. Potential readings were made only when the pH values of both urea solutions were the same. For the NaC1-urea and KC1-urea solutions, the potential differences were reproducible to f O . l mV for two or three electrode transfers between the cells. Results for LiC1-urea and CsC1-urea were less reliable (f0.2 mV) for the reasons mentioned above. Separate studies on pure salt solutions confirmed that the electrodes obeyed the Nernst equation within experimental error. Consequently, the theoretical Nernst slope, 118.3 mV, was used in all calculations. Isopiestic Measurements. The technique and apparatus have been described.13 Urea was used throughout as the reference solute. Generally, the solutions were allowed to equilibrate a t 25.0 f 0.1' for 1 week or less. The experimental results consisted of a set of molalities of the reference solute plus a set of molalities of various mixtures of urea and the salt of interest.
TABLE I: Coefficients for the Representation of AHtr/ml by Eq 3 Coefficients of data points
Range of m i , mol kg-'
hZ1('),cal
26 14 10 11 17 7 8 17 8 8 8 6 7 6 4
0.26-3 .O 0.40-3.0 0.25-3.0 0.2 5 -3.0 0.254.0 0.27-2.0 0.10-2.0 0.2 0-3.0 0.10-1.5 0.10-1.6 0.05-1.5 0.2 0-1.4 0 . 0 9 4 ) .8 0.20-1.5 0.10-1 .o
-237 -302 -370 -394 -292 -398 -533 -644 -748 -778 -684 -805 -824 417 -987
No.
Salt LiCI NaCl NaBr NaI KCI
CSCl MgClz CaC1, CaBr, CaI, SrClz BaCl, LaCiB Me4NBr Na2S04
The Journal of Physical Chemistry, VoL 79, No. 14, 1975
kg mol-'
+5 +3 +6 +
38 i: 5 88 i: 3 106 + 6 109 f 8 77 4 122 f 4 103 f 9 157 & 10 187 7 190 i 9 180 + 7 278 f 10 119 + 37 134 + 6 637 + 3 1
10
*
i5
+4 +8 f 11
*
+5
i7 i5 i8 i 22 & $:
5 21
'?..
Results Calorimetry. The accuracy of the calorimeter was determined by measuring the enthalpy of reaction of tris(hydroxymethy1)aminomethane with 0.1000 N HCI. Our enthalpy of -7113 f 3 cal/mol was in good agreement with that of Prosen and Kilday,14 -7115 f 1 cal/mol. The uncertainty of our value is the standard deviation of the mean of eight determinations. The enthalpy of solution of urea in water, AHsolnW, was measured as 3654 f 1 cal/mol for the molality range, 0.010.03 m. The uncertainty is the standard deviation of the mean of 22 determinations. This value is in good agreement with that of Egan and Luff15 who obtained a value of 3656 f 1 cal/mol by extrapolation of their data to infinite dilution. Our value also agrees within experimental error with that determined by Subramanian et a1.,16 3686 f 33 cal/ mol. The enthalpies of solution of urea in a variety of salt solutions, AHsolnws, were determined as a function of salt molality at urea molalities between 0.01 and 0.03 m. These data are given in Table lM.17 In separate experiments with NaCl a t fixed salt molality, it was shown that AHsolnWS was independent of urea molality in the range considered. These AHsolnWS values are, therefore, assumed to refer to infinite dilution with respect to the urea molality. The enthalpy of transfer of 1 mol of urea from water to a salt solution of a given salt molality
J%,(i),
cal kg3"
0
2
4
6
m,,MOLES
8
1
0
KG-'
Figure 1. Enthalpies of transfer vs. salt molality for the transfer of urea to solutions of: A,LiCI; 0, Nal; 0 ,CaCI2.
was calculated from the data at different salt molalities for each salt studied. Figure 1 shows the enthalpies of transfer of urea from water to solutions of NaI, LiC1, and CaCl2 as a function of salt molality. Plots of this kind suggested that the dependence of AH,, on salt molality was of the form
+
AH,, = hzl(0)ml + h21(l)m13/2 h21(2)m12+
(2)
where the h21(i) are empirical constants and m l is the salt molality.ls At sufficiently low salt molalities higher sensitivity in the extrapolation was obtained by ignoring all but the first two terms and rewriting eq 2 as
Enthalpies of Urea-Water Solutions
1393
TABLE 11: Coefficients and Their Standard Deviations for the Representation of (0 log y z / i ~ r n ~ ) ,by , Eq 4 Solute molality ranges
System LiCl-urea NaCl-urea KCl-urea CsCl-urea
No. of data points 70 38 45 24
Coefficients mol kg-*
rn 2, mol kg-'
gZ,(O), kg mol"'
0.0056-3.0 0.0013-2.6 0.0028-3.1 0.0012-3.2
0.15-5.7 0.24-7.6 0.434.8 0.95-6.8
-0.0652 i 0.0043 -0.0429 + 0.0021 -0.0288 i 0.0013 -0.0359 i 0.0028
l"il1,
The coefficients, hzl(O) and hzl(l), were then obtained from a plot of AHtr/ml vs. m11/2~In practice, a least-squares routine was utilized to obtain the best fit of the data. Table I gives (1) the salt molality ranges for which the fit was obtained, (2) the number of points utilized, (3) the coefficients, and (4) their uncertainties (standard deviations) for all salts studied. Plots of AHtr/ml vs. ml1I2 for NaC1-urea and for CsC1-urea are shown in Figure 2. The lines were calculated from eq 3 and the coefficients of Table I. The coefficient, h ~ l ( ~represents ), the enthalpy of transfer of 1 mol of urea a t infinite dilution in water to a salt solution of unit molality which has ion-ion interactions equivalent to those a t infinite dilution. As such, h~l(O)is a useful measure of the primary enthalpy of interaction of a urea molecule with a given pair of ions in the aqueous environment. Extraneous concentration dependences are absent. One may question whether the extrapolation to give hzl(O) (see Figure 2) is valid, i.e., whether the extension of the line to m l = 0 should be linear rather than curved. We tested this extrapolation procedure using the system MedNBr-urea. Cassel and Wen5 obtained enthalpies for the transfer of 1 mol of Me4NBr a t infinite dilution from water to urea solutions of various molalities. Extrapolation of their data to zero urea molality gives the coefficient, hlz(O), for this system as -392 f 20 cal/mol. Stern and cow o r k e r ~have ~ ~ shown that h12(O) must equal hz,(O) a t infinite dilution. Our value of h21(O) from Table I is -417 f 5 cal/mol which agrees with the value of Cassel and Wen for h12(O) within experimental error. Since the dependence of the enthalpies of transfer on solute molality must be different in the respective systems, it would be an unlikely accident to obtain numerical agreement within experimental error from two incorrect extrapolations. We suggest that this agreement supports the validity of our extrapolation procedure. Potentiometry. Table 2M gives the measured potential differences between the cells at various salt and urea molalities for the systems LiC1-urea, NaC1-urea, KC1-urea, and CsC1-urea.l' These data were analyzed by methods previously describedl1J2 to give values of the quantity (a log Y z / a r n d m z for each data point. Here yz is the molal activity coefficient of urea. For each salt-urea system, correlation was accomplished by means of an equation of the formz0
and a least-squares routine. Table I1 gives for each system the molality ranges in which data were obtained, the num-
gZl(*) 1k g3 / 2 molm312
0.0114 0.0223 0.0146 0.0148
i 0.0045 i 0,0027
f 0.0012 i 0.0017
g21(2), kg2 mol-'
0.00527 i 0,0014 0.00336 i 0.00065 0.00298 i 0.00046 0.00250 0.00084
*
-'0°/
Yd
-200-
-1
0
I
Y a -1
3 g - 300 2 '
c
I
a
-400
t
Figure 2. Enthalpies of transfer divided by the salt molality vs. the square root of the salt molality for the transfer of urea to solutions of: 17,NaCI; U, CsCI.
ber of data points taken, the values of the coefficients, and their standard deviations. Isopiestic Measurements. Table 3M gives the solute molalities of solutions in isopiestic equilibrium and the values of A/m1mz13 for the systems LiC1-urea and CsC1u1ea.l' The osmotic coefficients of aqueous solutions of urea, LiC1, and CsCl required in the calculation of A/mlmz were obtained from the l i t e r a t ~ r e . ~ l - ~ S The values of Alrnlmz were correlated using a leastsquares routine by means of an equation of the form20
+
+
Alrnlmz = 2.303(g21(~) gzl(1)m11/2 g21(2)m2]
(5)
For the two systems, Table I11 gives the molality ranges in which data were obtained, the number of data points taken, the values of the coefficients, and their standard deviations. Comparison of Potentiometric and Isopiestic Results. In principle, the numerical values of the coefficients of eq 4 and 5 should be the same. Comparison of the coefficients in Tables I1 and 111 indicates that the agreement is good in the case of the urea-LiC1 system and poor for CsC1-urea. Some variation in the parameters is produced by the different ranges over which the data are fit. A better test of the agreement of the data from the two The Journal of Physical Chemistry, Vol. 79, No. 14, 1975
M. Y. Schrier, P. J. Turner, and E. E. Schrier
1394
TABLE 111: Coefficients a n d Their Standard Deviations for the Representation of A/mlm2 by Eq 5 Solute molality ranges Coefficients
No. of data
'97 2,
1,
System
points
mol kg-'
mol kg-'
LiCl-urea CsCl-urea
16 23
0.12-1.45 0.11-2.13
0.24-3.33 0.234.14
TABLE IV: Comparison of the Values of -log
0
0.0576 (0.0481) 0.0523 (0.0437) 0.0470 (0.0394) 0.0418 (0.0350) 0.0365
0.1088 (0.0914) 0.0983 (0.0826) 0.0878 (0.0738) 0.0772 (0.0650) 0.0667
y2/y2O
g2,('), kg mol-' 4 . 0 5 4 1 i 0.0020 4 . 0 5 7 2 i 0.0043
g21(2), kg2 mol-'
*
0.00439 & 0.0065 0.00377 i 0.0013
0.00889 0.0023 0.0322 i 0.0044
Obtained from Potentiometric and Isopiestic Measurementsa9b
0.1560
0.0280 (0.0130) 1 0.1402 0.0247 (0.0113) 2 0.1244 0.0213 (0.0097) 3 0.1086 0.0179 (0.0083) 4 0.0928 0.0146 (0.0070) 5 0.0312 0.0562 0.0770 0.0112 (0.0058) 6 0.0260 0.0456 0.0612 0.0079 (0.0048) 7 0.0045 (0.0039) a Isopiestic results are in parentheses. Units of m l are moles of of water. Values in parentheses are from ref 25,
methods is provided by comparing the values of -log y2/ yz0 calculated from the coefficients by means of the equation
where yzo is the activity coefficient of urea in a solution of molality r n 2 which contains no salt. The comparison between the potentiometric and isopiestic results is shown in Table IV. The values calculated from the isopiestic data are indicated in parentheses. The quality of the agreement between the two sets of values in the LiC1-urea and CsC1-urea systems is now comparable and is satisfactory. It does not approach what can be achieved (see the recent study of Phang and Steelz4 on glycine-NaC1) but (1) the slow hydrolysis of urea, (2) the possible attack of the silver electrodes by urea, (3) the low level of the response of the electrodes to Li+ and Cs+, and (4) the generally small deviations from ideality in these systems tend to limit the quality of the data. The isopiestic data of Bower and Robinsonz5 are compared with our potentiometric data for NaC1-urea in Table IV. The agreement is poorer than in the other cases. In addition to the points just given, another possible explanation is the lack of emphasis in their isopiestic study on the low solute molality region. In further discussion, we shall employ the potentiometric values of gZ1(O) for all systems investigated here. The Journal of Physical Chemistry, Vol. 79, No. 14, 1975
gZ1('), kg3l2mol-3/2
0.0437 0.0515 (0.0192) (0.0208) 0.0370 0.04 14 (0.0167) (0.0184) 0.0303 0.03 13 (0.0145) (0.0162) 0.0236 0.0212 (0.0124) (0.0143) 0.0169 0.0111 (0.0106) (0.0127) 0.0101 0.001 1 (0.0090) (0.0112) -0.0090 0.0034 (0.0077) (0.0101) 4.0033 4,0191 (0.0065) (0,0091 salt per kilogram of water;
0.0261 (0.0357) 0.0236 (0.0320) 0.0210 (0.0282) 0.0185 (0.0244) 0.0160 (0.0206) 0.0135
0.0439 (0.0537) 0.0389 (0.0461) 0.0339 (0.0386) 0.0289 (0.0311) 0.0239 (0.0235) 0.0189
0.0189
0.0110
0.0139
0.0114
0.0085
0.0089
0.0039
units of
m2
0.0564 0.0489 0.0414 0.0339 0.0264
are moles of urea per kilogram
Discussion The set of h21(O) and g21(O) coefficients obtained in this work together with the g21(O) value from the NazS04-urea system13 and the h12(O) and g12(O) values obtainable from the data of Wen and collaborator^^^^ for the Me4NBr-urea and Bu4NBr-urea systems comprise the most complete set of limiting interaction coefficients available for a given nonelectrolyte in various salt solutions. Although a quantitative interpretation of these results is well beyond the present level of theoretical understanding of these interactions, a useful working perspective can be developed. The trends in the h21(O) values in Table I are marked. As with the formamide-salt systems studied previously,1° the exothermicity of the hz,(O) values increases with increasing ion size for a given series of alkali or alkaline earth halides. These is also a roughly 1:2:3 ratio of the numerical values of h21('3)from the alkali halides to the alkaline earth halides to lanthanum chloride. A few values of hzl(0) for formamide-salt systems are available. They are, h21(O) = -275, -305, and -281 cal/mol for NaC1, NaBr, and KC1, respectively. The uncertainties for these numbers are f 1 5 cal/mol. The average 10%difference between these numbers and those for urea given in Table I is in keeping with the 12% difference in dipole moments between urea, 4.20 D, and formamide, 3.71 D.26 Alternatively, this difference could be the effect of the addi-
1395
Enthalpies of Urea-Water Solutions
6oor
TABLE V: Values of the Limiting Thermodynamic Quantities of Transfer of Urea from Water to Various Salt Solutions at 25"
I
400
RT In gZl(O),
Salt LiCl NaCl KC 1 CsCl Me4NBr" Bu4NBr" Na2S04b
kZi(0),
cal kg mol-' cal kg mol"
*
-89 6 s 9 * 3 -39 2 4 9 1 4 -148 -166 4 6 a Data from ref 5 . RT lng,,'ol
*
s.21
(0)
,
gibbs Itg mol-'
-237 f 5 -302 f 3 -292 f 5 -398 4 -392 i 20 466 f 17 -987 j, 21 data from ref 13.
-0.50 f 0.04 -0.81 0.02 -0.85 i. 0.02 -1.17 i 0.03 -4.82 2.12 -3.02
2oo
*
P
t
ot W
d
-200
-
-400
c
0
-N
L
tional -NHz group present in urea. In any case, urea appears to be quite similar to formamide in its interactions with salts, an important consideration in these model studies. In our discussion of the differences between h z ~ (for ~) formamide and urea, we suggested the possible significance of ion-dipole interactions in these systems. In our view, the contributions of terms involving the work of cavity formation, ion-dipole interactions, dispersion interactions, etc. to the numerical magnitude of the parameters are of real importance. Limited correlationsz7which can be made involving hz,(O) as a function of the ionic radii and the ionic polarizabilities add strength to this contention. There seems little doubt, however, that the trends in these parameters and the companion entropy of transfer values are due to perturbations of the structure of water produced by the solutes. Table V collects the values and their uncertainties of h ~ l ( RT ~ ) ,In gZI(O), and sz1(O) calculated from the data of the various investigators. The quantity RT In g21(O) represents the Gibbs free energy of transfer while sz1(O) is the entropy of transfer of 1 mol of urea at infinite dilution in water to a salt solution of unit molality which has ion-ion interactions equivalent to those a t infinite dilution. A plot of hz1(O)vs. s21(O) is shown in Figure 3. The slope of the line, T * ,is 280 f 8°K and the intercept is 104 f 13 cal/mol. The fit standard deviation is 15%. This type of enthalpy-entropy compensation behavior with T* N 280'K appears to be characteristic2s~29 of water structure related processes. Friedman's cosphere modelz9 provides a useful framework with which to rationalize the trends in the data. In this model, a cosphere may be thought of as a layer of water molecules around a solute particle which is perturbed by the presence of the solute. The distinction is made in considering these cospheres between hydration of the first and second kind.29While hydration of the first kind arises from ionic or polar field effects, hydration of the second kind refers to a perturbation produced in water by the presence of a solute particle but without the influence of a directional solute-solvent force. For ions and other polar molecules, two, more or less distinct cospheres might correspond to these levels of hydration. They are type I and type I1 cospheres. We deal only with type I1 cospheres. Urea is assumed to possess a type II& cosphere. The subscript, sb, stands for structyre breaking meaning that the water in this cosphere contains fewer hydrogen bonds than bulk water. The ions we consider fall into three categories. They may have no
Bu4NBr
t 14
- 1000
I
Na2S04
8
-3
-2 s21(0),
-I
0
I
2
GIBBS MOLE-'
Figure 3. Enthalpy-entropy compensation plot for the transfer of urea from water to salt solutions.
type I1 cosphere, may have a type II& cosphere, or may have a type 11,, cosphere. The subscript, rg, stands for rare gas. The structure in this cosphere may be considered analogous to that surrounding a rare gas molecule, i.e., there are more hydrogen bonds than in bulk water. In discussing ion-urea interactions, we assume that overlap of the cospheres of urea with those of the ions leads to destruction of some of the cosphere material of one or both of the ions. Starting with the interactions at the upper right hand corner of Figure 3, overlap of the type II& cosphere of urea with the type 11,, cosphere of Bu4N+ leads to a net destruction of some of the type II,, cosphere, a relaxation of the structured water to bulk water, and a concomitant gain of enthalpy and entropy. Overlap of the type II& cosphere of Br- with that of urea will lead to a net destruction of some of the type II& cosphere, a relaxation of unstructured water to bulk water, and a concomitant loss of enthalpy and entropy. The sum of the effects in the solution leads to a net gain of enthalpy and entropy in the transfer process. A p p a r e n t l ~ cosphere ,~~ I1 for Li+ may also be thought of as the rare gas type. For the interaction of LiCl with urea, one obtains a negative enthalpy and entropy of transfer. The processes involving the urea-Li+ interaction are similar to those discussed above for urea-Bu4Nf but the effect is smaller and (1) the interaction of urea with C1- and (2) the influence of hydration of the first kind lead to negative values in this case. We next consider NaC1. Since Na+ does not appear to give rise to hydration of the second kind, the enthalpy and entropy values are more nearly the sum of contributions of directional interactions and a cavity term in this case than for any of the other salts considered. While the enthalpies and entropies of transfer for KCl are similar to those for NaC1, the values for CsCl are more negative than for NaCl because of the overlap of the type II& The Journal of Physical Chemistry, Vol. 79, No. 14, 1975
M. Y. Schrier, P. J. Turner, and E. E, Schrier
1396
cospheres of the ions with that of urea resulting in the loss of some of the ionic cosphere material. Although the position of Me4NBr may be partially the result of the Br- ion, it does suggest some structure-breaking character to the Me4N+ cosphere. Finally, S042- exhibits either exceptionally large structure-breaking tendencies or a strong H-bonding interaction with urea. We next consider the effectiveness of these electrolytes as protein denaturants. Denaturation by electrolytesg has been thought of in terms of interactions of the electrolytewater medium with groups which are buried (not exposed to the solvent) in the native state of the protein but which become exposed as the protein is denatured. The free energy of denaturation is comprised of opposite contributions from the unfavorable salting out of nonpolar groups and the favorable interactions of the amide linkage with the salt medium. Ying et al.30 have obtained data describing the isothermal denaturation of ribonuclease by different salts a t 25'. They included LiC1, LiBr, LiC104, NaC104, NaSCN, and CaC12. Other alkali chlorides and Me4NBr did not act as denaturants in the concentration range investigated. Although Na2S04 was not studied, it is well known to be an antidenaturant both from thermal denaturation resultsg and general experience. Urea may be thought of as a model for the amide linkages which are exposed to the solvent as the protein denatures. Although we hoped that the values of RT In g21(O) would be helpful in determining the denaturing ability of salts, they provide little assistance. The value of LiCl is more negative than that for the other alkali chlorides but is the same as that for NazS04. There is, of course, the difficulty here of accounting for the larger salting out of nonpolar groups by Naps04 than by LiC1. The case of MesNBr is unambiguous, however. The large negative value of RT In gn(O) for MerNBr suggests that it should be a strong denaturant but it is not. Since Wen and Hung31 have shown that Me4NBr salts-in nonpolar compounds, there would be no loss of denaturing ability through interaction with nonpolar groups. The concentration of effort by various investigator^^^^^ on limiting interaction coefficient^^^ has come about because of the desire to interpret the thermal denaturation of proteins, a process which can be described by parameters obtained at low salt concentrations. However, high salt concentrations are required for isothermal denaturation by salts, e.g., 6.2 m LiCl is the midpoint concentration in the denaturation of ribonuclease at 25°.30 Many workers have assumed that the sign and magnitude of the limiting interaction coefficients determine the salt-nonelectrolyte interaction up to high molalities of the solutes. Examination of available Gibbs free energy data for higher solute molalities than those studied here4J3~25,30 indicate that the situation is more complex. In the systems, M e d N B r - ~ r e a ,B ~ ~q N B r - ~ r e a ,and ~ ~ N a C l - ~ r e a ,a~t ~m l > 1, m2 < 3, the quantity, (a2 log yz/am12)m2,is positive. In NazSO.purea, log y2 even changes sign from negative to positive a t N a ~ S 0 4molalities greater than 0.75 m. In contrast, in the LiC1-urea system (a2 log y2/arn1~),~is positive up to 3 m LiCl but is negative thereafter. This observation may be significant both for the thermodynamic description
The Journal of Physlcal Chemistry, Vol. 79, No. 14, 1975
of isothermal protein denaturation and with regard to the controversy which has grown up regarding the interaction of lithium salts with amide linkages in aqueous solution. Thermodynamic studies of both the 1,iCl-urea system and other urea systems containing denaturing salts a t high solute molalities will be the subject of a later publication. Acknowledgment. This work was supported in part by Grant No. GM 11762 from the Institute of General Medical Sciences, U S . Public Health Service. Supplementary Material Available. Tables l M , 2M, and 3M will appear following these pages ih 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 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, D.C. 20036. Remit check or money order for $4.00 for photocopy or $2.50 for microfiche, referring to code number JPC-751391.
References and Notes R. H. Stokes, Aust. J. Chem., 20, 2067 (1967). H. S. Frank and F. Franks, J. Chem Phys., 48, 4746 (1968). D. B. Wetlaufer, S. K. Mallk, L. Stoller, and R. I. Coffin, J. Am. Chem. SOC.,86, 509 (1964). (4)A. Ben-Naim and M. Yaacobi, J. Phys. Chem., 78, 170 (1974). (5)(a) W. Y. Wen and C. M. L. Chen, J. Phys. Chem., 73, 2895 (1969);(b) R. 8. Cassel and W. Y. Wen, /bid., 78, 1369 (1972). (6)J. H. Stern and J. D. Kuiluk, J. Phys. Chem., 73, 2795 (1969). (7)N. Desrosiers, G. Perron, J. G. Mathieson, 9.E. Conway, and J. E. Desnoyers, J. Solution Chem., 3 , 769 (1974). (8) R. B. Cassel and R. H. Wood, J. Phys. Chem., 78,2460(1974). (9)P. H. von Hippel and T. Schleich in "Structure and Stablllty of Blologlcal Macromolecules", S. Timasheff and Q. Fasman, Ed.. Marcel Dekker. New York, N.Y., 1969,p 417. E. R. Stimson and E. E, Schrier, J. Chem. €ng. Data, 10, 354 (1974). M. Y. Spink and E. E. Schrler, J. Chem. Thermodyn., 2, 821 (1970). F. L. Wilcox and E. E. Schrler, J. Phys. Chem., 76, 3757 (1971). E. E. Schrier and R. A. Robinson, J. Blol. Chem., 245, 2432 (1970). E. J. Prosen and M. V. Kilday, J. Res. Natl. Bur. Stand. U.S., Sect. A, 77,
561 (1973). E. P. Egan, Jr., and B. B.Luff, J. Chem. Eng. Data, 11, 192 (1966). S. Subramanian, T. S. Sarma, D. Baiasubrarnanian, and J. C. Ahiuwalia, J, Phys. Chem., 75, 815 (1971). See paragraph at end of text regarding supplementary material. Throughout this paper, the subscript 1 will refer to the salt while the subscript 2 will indicate urea. J. H. Stern, J. Lazartic, and D. Fost. J. Phys. Chem., 72, 3053 (1968). The coefficients, g;,1(0, replace the coefficients designated A, B, C- or a, b, c- in earlier H. David Ellerton and P. J. Dunlop, J. Phys. Chern., 70, 1631 (1966). H. F. Gibbard, Jr., G. Scatchard, R. A. Rousseau, and J. L. Creek, J. Chem. Eng. Data, 10, 293 (1973). R. A. Robinson and R. Stokes, "Electrolyte Solutions", 2nd ed, Butterworths, London, 1965,p 465. (24)S. Phang and E. J. Steel, J. Chem. Thermodyn.,6, 537 (1974). (25)V. Bower and R. A. Robinson, J. Phys. Chem., 67, 1524 (1963). (26)A. L. McClellan, "Tables of Experimental Dipole Moments", W. H. Freeman, San Francisco, Calif.. 1963,p 44,45. (27)E. E. Schrier, unpublished calculations. (26)R. Lumry and S. Rajender, Biopolymers, 9, 1125 (1970). (29)H. Friedman in "Water, A Comprehensive Treatise". Vol 3, F. Franks, Ed., Plenum Press, New York, N.Y., 1973,p 50. (30)A. Ylng, D. Blazej, M. Y. Schrier. and E. E. Schrier, unpublished results. (31)W. Y. Wen and J. H. Hung, J. Phys. Chern., 74, 170 (1970). (32)P. K. Nandi and D. R. Robinson, J. Am. Chem. SOC.,94, 1299 (1972). (33)The parameter, g2,(0), is equivalent to the conventional Setschenow parameter when a correction is made for the change of concentration scale.
