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@ Copyright, 1969, by the American Chemical Society
VOLUME 73, NUMBER 9 SEPTEMBER 1969
Thermodynamics of Aqueous Mixtures of Electrolytes and Nonelectrolytes. VIII.
Transfer of Sodium Chloride
from Water to Aqueous Urea at 25” by J. H. Stern and J. D. Kulluk Department of Chemistry, California State College at Long Beach, California 90801 (Received March 24, 1069)
Enthalpies of transfer of sodium chloride over the mixed-solvent range from pure water to urea mole fraction ( X J = 0.18 (42 wt %) were determined calorimetrically a t 25”. These were combined with free energies of transfer a t constant molality (per 1000 g of mixed solvent) calculated from isopiestic activity coefficients to yield entropies of transfer. The positive free energy is a linear function of X 3 . Since the coulombic part of the free energy is qualitatively predicted to decrease with X 3 ,the structural contributions to the free energy are positive. Both the enthalpy and entropy of transfer decrease with X 3 and are considerably greater in magnitude than the free energy. The negative values are in qualitative agreement with those predicted by the Born equation and also indicate that urea may act as a net water structure breaker across the entire measured composition range. The relationships between this transfer and those of amino acids are briefly discussed.
Introduction Aqueous urea is a most important biophysicochemical mixed solvent (ms) which has long been the subject of investigations ranging in scope from heat capacity1 to ultrasound absorption measurements.2 A frequent aim of recent studies has been to learn more about the complicated effect of urea on water. There is no consensus, but many views favor the concept of urea as a net structure breaker.*-6 The thermodynamic properties of urea-water are very well characterized and have been reviewed by StokesS6 Relatively little is known about the thermodynamics of ternary solutions of aqueous urea-containing electrolytes. Early studies on the solubility of sparingly soluble salts have been reviewed by Cohn and Edsall.’ Recently, the isopiestic vapor pressure study of Bower and Robinson8 has led to an analytical equation for the mean ionic activity coefficient of NaCI, relative to the infinitely dilute standard state in pure water, as a function of both urea and salt concentrations. At zero molality of NaCl in the ms (and also at low concentrations) this equation yields the primary
mediums or ion-ms interaction activity coefficient, based on molality of NaCl per 1000 g of water. Upon transformation, this activity coefficient yields the free (1) F. T . Gucker and H. B. Pickard, J . Amer. Chem. SOC.,62, 1464 (1940); E. P. Egan, Jr., and B. B. Luff, J. Chem. Eng. Data, 11, 192 (1966); see also B. E. Conway, Ann. Rev. Phys. Chem., 17, 481 (1966). (2) K. Arakawa and N. Takenaka, BuEl. Chem. Soc. Jap., 40, 2739 (1967); G. G. Hammes and P. R. Schimmel, J . Amer. Chem. Soc., 89, 442 (1967). (3) If. S. Frank and M . W. Evans, J. Chem. Phys., 13, 507 (1945); see also H. S.Frank and F. Franks, ibid., 48, 4746 (1968). (4) F. Franks, “Physico-Chemical Processes in LMixed Aqueous Solvents,” American Elsevier Publishing Co., Inc., New York, N. Y., 1967, p 50. (5) J. A. Rupley, J. Phys. Chem., 68, 2002 (1964); for a different interpretation, see A. Holtzer and M. F. Emerson, ibid., 73, 26 (1969). (6) R . H. Stokes, Aust. J . Chem., 20, 2087 (1967). (7) E. J. Cohn and J. T. Edsall, “Proteins, Amino Acids and Peptides,” Hafner Publishing Co., New York, N. Y., 1965, p 254. (8) V. E. Bower and R. A. Robinson, J . Phys. Chem., 67, 1524 (1963). (9) G. N. Lewis and M. Randall, “Thermodynamics,” revised by K. S. Pitzer and L. Brewer, 2nd ed, McGraw-Hill Book Co., Inc., New York, N. Y., 1961, p 591.
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J. H. STERNAND J. K. KULLUK
energy of transfer bFz at constant molality per 1000 g of ms. This contribution reports on the calorimetric enthalfrom pure water to aquepies of transfer of NaC1, ous urea ranging in composition from dilute solutions to ms of urea mole fraction X3 = 0.18 or 42 wt %. The enthalpies of transfer were obtained by the difference of the enthalpies of solution of crystalline NaCI in the ms and in pure water, AH2 and AHzO,respectively
m2,
a
= AH2
- AHz’
(1) Combined with free energies, the enthalpies of transfer yield entropies of transfer. The latter two thermodynamic properties have in previous ms studies’O proved to be more revealing than in showing the complicated effect of changed environment on the transferred solute.
a
Experimental Section Calorimeter. The general experimental procedure and the calorimeter have been described elsewhere.’oJl Materials. AI1 materials were AR grade, and the ms were freshly prepared with distilled and deionized water (450 ml per run).
Results and Discumion Summaries of the enthalpies of solution of XaC1 to
m2molal in pure water ( A H , ” ) and in the ms (AH2) are shown in Table I. All uncertainty intervals associated
I
-1800 0
I
I
0.04
1
I
I
I
0.12
0.08
I
I
0.16
I
0
x3
T , a ,t , and TTSZwith XI. Figure 1. Variation of A Table I : Enthalpies of Solution of NaCl in Water (AHa’) and Aqueous Urea (AHz) No. of runs
m2 X 102
6
0.7-1.5
6 4 4 3 3
1.0-1.3 0.5-1.2 0.8-1.0 0.6-1.2 0.9-1.0
x,x
10%
0.00 (AHz’)
1.77 5.13 8.26 12.6 17.8
AH20 and AB2,cal/mol
985 i 15 805 ct 30 425 i: 30 195 i 15
-10 ct 30 -185 i 30
with mean values are their standard deviations multiplied by factors necessary to give 90% confidence levels. The value of AH,” is within 15 cal of that calculated from the enthalpy of dilution combined with the “best” value of the enthalpy of solution of NaCl to infinite dilution.12 It is of interest to estimate how closely i\Hzrepresents the transfer of NaCl at infinite dilution. The differences in enthalpy of dilution from m2 approximately 0.01 m to r n 2 = 0 in water and in the ms, respecwould yield the enthalpy of tively, combined with transfer at infinite dilution. In pure water the enthalpy of dilution is about -40 cal/mol.12 Extrapolationla in the ms requires values of the dielectric constant D and its temperature coefficient. The latter is
x2,
The Journal of Physical Chemistry
subject to high uncertainty (discussed by Harned and Owenz4), and since neither quantity appears to be available above X 3 = 0.05,l6 direct meaningful estimates are not possible. Turning to other related studies for estimates, the enthalpy of dilution of 0.01 m KC1 in 5% by weight of urea (X,= 0.016) appears to be nearly equal to that in pure water,14while in formamide (D = 109.5) the calculated enthalpy of dilution for univalent salts from cu, 0.01 rn to infinite dilution is 1.5 cal/mo1.16 The estimated dielectric constant of the most concentrated solution in the present study is less than 1OO.l’ The enthalpy of dilution (ma = 0.01 to m2 = 0) in urea solutions may thus be between the (10). (a) J. H. Stern and J. Nobilione, J . Phys. Chem., 72, 1064 (1968); (b) zbid., 72, 3937 (1968). (11) J. H. Stern and C. W. Anderson, ibid., 68, 2528 (1964). (12) V. B. Parker, “Thermal Properties of Aqueous Uni-Cnivalent Electrolytes,” NSRDS-NBS2, National Bureau of Standards, U. 8. Government Printing Office, Washington, D. C., 1965. (13) See ref 9, p 339. (14) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” 3rd e d , Reinhold Publishing Gorp., New York, N. Y., 1958, p 344. (15) Landolt-Barnstein, “Zahlenwerte und Funktionen,” Vol. 11, part 6, Elektriache Eigenschaften, Springer-Verlag, Berlin, 1959, p 777. (16) G. Somsen and J. Coops, Rec. Trav. Chim., 84, 985 (1965). (17) See ref 7, p 144.
TRANSFER OF SODIUM CHLORIDE FROM WATERTO AQUEOUSUREA value of the enthalpy of dilution in pure water (-40 cal/mol) and zero. Since the error in AHz in these experiments is 15-30 cal/mol, the values measured at m2 = 0.01 can be understood as applying to infinite dilution of NaC1. Figure 1 shows values of together with hFz calculated a t the experimental compositions of this study, and T S 2plotted as a function of X3. The convention used in Figure 1 is consistent with previous ms studies in this series. Such a plot includes the advantage that AFz is linear and may be expressed by10aJ8
x z
I _
AFz = 2 X 2.303RTKXa
(2)
where K is the ion-ms interaction coefficient. This behavior was also observed for HC1 in aqueous acetic acidlOa and in aqueous ethylene glycollob over wide plotted on the basis of transfer ranges of Xa,while a t constant mole fraction of the solute has been found not t o be linear and thus not as useful. with Xa may be conThe increasing values of trasted with the Born coulombic free energies of transfer.loa I n this system they are predicted to be negative and t o decrease with Xg as the dielectric constant of the ms increases. The difference may be partially and qualitatively accounted for by positive structural contributions t o the free energy. This is, of course, subject to the serious limitations imposed by uncertainties of the Born treatment. Kevertheless, it appears from free-energy considerations that the structural environment of the ms is less favorable t o NaCl than pure water and that this effect predominates over the favorable coulombic contribution. and T a 2 are initially ca. three The changes of times steeper than the increase of hF2 with Xa. The magnitude of these effects points out the large difference in solvent environment to an ionic solute, even though urea has been regarded as a most compatible cosolvent with watern4
z2
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The negative values of A z and T a 2 point toward structure breaking by urea,lob It may be noted, however, that the coulombic entropy of transfer is also negative. The observed results may, of course, be the results of many complicated and opposing effects, and a detailed accounting is not possible. It is possible that similar interactions are in part responsible for the effect of urea on both electrolytes and “small” dipolar zwitterion amino acids, such as glycine. The solubility of glycine in urea-water is lower than in pure water with a positive free energy of transfer.19 Negative enthalpies of transfer (- 1010 cal/mol to ca. X S = 0.13) have also been reported.20 However, amino acids with larger aromatic or aliphatic side chains showed an increase in solubility in the ms with decreasing free energies of transfer from water with increasing side-chain size. It appears that polar groups or ions undergo unfavorable free-energy interaction with ms relative to water, with opposite effects for the nonpolar groups. The solubility of some amino acids may, however, be high enough to yield free energies which include contributions from solute-solute or selfinteraction effects.21 The isopiestic method is preferable, since it also yields free energies at zero solute concentration, as in the case of a-amino-n-butyric acid.22 Studies on the enthalpy of transfer of this amino acid have been initiated.
Acknowledgment. The authors wish to thank Professor R. D. Bauer for fruitful conversations leading t o the study of the urea-water system as a ms medium, and the National Science Foundation for financial assistance. (18) See ref 9,p 593. (19) Y.Nomki and C. Tanford, J . Biol. Chem., 238, 4074 (1963). (20) G. C. Xresheck and L. Benjamin, J . Phys. Chem., 68, 2476 (1964). (21) F.A. Long and W. F. McDevit, Chem. Rev., 51, 119 (1952). (22) E.L. Cussler, Jr., J . Phys. Chem., 71, 901 (1967).
Volume 78, Number 9 September 1069
