THERMODYNAMICS OF AQUEOUB MIXTURES OF ELECTROLYTES AND NONELECTROLYTES species charge is high. Such ionic association would cause the calculated d values to be smaller than in the solid state. Apparently, the forces of interionic attraction are even further reduced in the complex salt solutions than in LaCL. Of course, one cannot overlook the possibility that numerous effects are contributing to the activity coefficient and are being absorbed by the single d parameter. However, Poirier's calculations serve to emphasize that this single variable is the most
167
significant one in solutions of low ionic strength. Many additional data are required in order to substantiate this point and to determine the effect of size and structure upon the activity coefficients of polyvalent electrolytes.
Acknowledgment. The authors wish to thank Mr. L. R. Wolf for his assistance in the computer calculations.
Thermodynamics of Aqueous Mixtures of Electrolytes and Nonelectrolytes.
IX. Nitromethane in Pure Water and in 1 rn Potassium Chloride from 15 to 35"
by J. H. Stern and J. T. Swearingen Department of Chemistrv, California State College at Long Beach, Long Beach, California 90801 (Received April 17, 1960)
Calorimetric enthalpies cf solution of liquid nitromethane to low concentrations (0.01-0.02 m) in pure water and in 1 m potassium chbride, respectively, are reported from 15 to 35". The enthalpies are positive, with the former giving rise to a third-srder dependence on temperature illustrating the complicated nature of interactions in the purely aqueous environment over a 20" span. The latter appears to be linear with temperature and is lower, showing the striking change in the solution as a result of the added water-structure-breaking electrolyte. The temperature derivatives of the enthalpies yield the appropriate heat capacity differences between the dissolved and pure liquid nitromethane. While the heat capacity difference to pure water changes rapidly and ranges from 73 to 115 cal/deg mol between 15 and 35" with a minimum of ca. 10 cal/deg mole near 24", that to 1 m ICCl remains constant a t 25 cal/deg mol. The positive heat capacity differences,especially when their values are high, may be assumed to be due to net structure making by nitromethane. In pure water near 25", complicated changes in nitromethane-water interaction caused by water-structure shifts may be taking place, while in KCl, specific ion-nitromethane and ion-water interactions may predominate over the entire temperature range. Heat capacities of transfer of dissolved nonpolar gases from water to electrolyte solutions, calculated from solubility measurements, are briefly discussed.
Introduction This contribution reports on the calorimetric enthalpy of solution of liquid nitromethane in pure water, AH3", and in 1 m KCl, AH3, at 2 to 5" intervals between 15 and 35". These data also yield the enthalpy of transfer, Z a , for CH3NOz(HzO) = CH3NOz(1 m KCl), where A T 3 = AH3 - AHSO, and the heat capacity differences for nitromethane between the following final and initial states: aqueous and pure liquid, ACpao,in 1 m KC1 and pure liquid, AC,,, and the heat capacity of = AC,, - AC,,". Heat transfer, KC,,, where capacity differences have proved to be valuable in the complicated problem of water structure by showing the part of the solute in local structure changes. For
a,,
example, large positive heat capacity differences observed for noneleotrolytes containing nonpolar groups have been attributed to their net structure-malting.' No other calorimetric study of this type, in the presence and absence of an electrolyte a t temperatures other than 25", appears to have been reported to date. Many nonelectrolytes are not suitable since their low solubility precludes calorimetric measurements. Heat capacities may also be calculated from the second derivatives of analytical equations of the solubility with (1) (a) J. H.Stern and A. Hermann, J . Phvs. Chem., 72,364 (1968); (b) G.C.Kresheck and L. Benjamin, ibid., 68,2476 (1964); (c) L.A. D'Orazio and R. H. Wood, ibid., 67, 1435 (1963); (d) H.9. Frank and W. Wen, Discussions Faraday Soc., 24, 133 (1957). Volume 74- Number 1 January 8, 1970
168
J. H. STERN AND J. T. SWEARINGEN
temperature. Such data may, however, be among the most uncertain of therm3dynamic properties since repeated differentiation of smoothed equations causes losses in precision. For example, studies in the field of aqueous weak electrolyte ionization indicate that calorimetric AC, for ionization is frequently a complicated function of temperature. Such results would not be attainable from the second derivative of the ionization constant with temperature.' Nitromethane was chosen for two main reasons. The low enthalpies of solution in water and in 1 m KCI have in a previous study3at 25' yielded accurate enthalpies of transfer = -150 cal/mol, AH3O = 630 cal/mol). The transfer was later experimentally shown to take piace in the limiting region of ion-nonelectrolyte intera ~ t i o nthat ; ~ is, enthalpies of transfer consist of additive contributions from the cation and anion and that nitromethane solute-solute or self interaction eff ects5 are negligible at the experimentally low nitromethane concentrations (ca. 0.01-0.02 m). Vaporization of nitromethane at these low concentrations is negligible. The electronic absorption spectrum in water shows large solvent shifts relative to the spectrum in n-heptanea which are of the order of magnitude of hydrogen bond enthalpies. Since both equilibria CH3NOZ ~2 H+
Table I : Enthalpies of Solution of Nitromethane in Pure Water
a
(a3
+
OH CH2N02-(K S 1O-l1) and CHJY02
+/ ~2 CH2=N
\
0lie very far on the left si de,'^* the principal form of liquid and aqueous nitromethane appears to be the
/"+\
resonance hybrid CH3-N
*
0
Experimental Section Measwements and Materials. The calorimeter has been described previously. Measurements were carried out within ca. 0.1" of the reported temperatures and were initiated a t balanced Wheatstone bridge decade settings corresponding to these temperatures, checked by a calibrated mercury-in-glass thermometer accurate to 0.01". The experimental procedure and materials were described el~ewhere.~ Results and Discussion All enthalpies of solution are summarized in Tables I and I1 with uncertainty intervals equal to the standard deviations (a) of the tabulated values. The over-all concentration of nitromethane in all solutions ranged from 0.01 to 0.02 m. Both AH3' and AH3 increase rapidly with temperature, and the former doubles in magnitude approximately with each 10" rise in temperature. The AHa" curve in The Journal of Physical Chemistry
T,OK
No. of runs
AH8', cal/mol
288 291 293 295 298 303 308
5
350 10 550 f 20 590 =!= 10 600 f 20 630 =!= 10" 790 & 10 1150 i 10
4
3 4
4 7 4
Data from ref 3.
Table I1 : Enthalpies of Solution of Nitromethane in 1 m KC1 T,O
No. of runs
X
3 4 7 4 4
288 293 298 303 308 a
AH,, oal/mol
240 360 480 570 750
i 10 f 10 f loa f 20 f 10
Data from ref 3.
Figure 1 represents the least-square fit of enthalpies in Table I with their weighting factors of 1 / c 2 to a power series with temperature AH3" = 650
+ 13.3(T - 298) + 1.05(T - 298)' -I- 0.269(T - 298)'
(1) while that for AH3 is based similarly on the least-squares equation
AH3
=
488
+ 25.4(T - 298)
(2) The experimental data from Tables I and I1 are shown in Figure 1 by circles and triangles for AH3' and AH3, respectively. The estimated uncertainty of enthalpies calculated from eq 1 and 2 is *30 cal/mol. Values of AC,," and AC,, are obtained directly from the temperature derivatives of eq 1 and 2, with an estimated error of *3 cal/deg mol, and are shown in Figure 2 piotted against temperature. There are great differences in AC,,O and AC,, reflecting changes of interaction in both media. The former has remarkably (2) E. J. King, "Acid-Base Equilibria," The Macmillan Co., New York, N. Y., 1965, p 192. (3) J. H. Stern and A. Hermann, J . Phya. Chem., 71, 309 (1967). (4) J. H. Stern, J. Lazartic, and D. Fost, ibid., 72, 3053 (1968). (5) G. N. Lewis and M. Randall, "Thermodynamics," revised by K. S. Pitzer and L. Brewer, 2nd ed, McGraw-Hill Book Go., Inc., New York, N. Y., 1961, p 587. (6) G. C. Pimentel and A. L. McClellan, "The Hydrogen Bond," W. H. Freeman and Co., San Francisco, Calif., 1960, p 158. (7) J. D. Roberts and M. C. Caserio, "Basic Principles of Organic Chemistry," W. A. Benjamin Inc., New York, N. Y., 1964, p 690. (8) J. E. Prue, "Ionic Equilibria," Pergamon Press, Ltd., Oxford, 1966, p 88. (9) J. H. Stern and C. W. Anderson, J. Phys. Chem., 68,2528 (1964),
THERMODYNAMICS OF AQUEOUS MIXTURESOF ELECTROLYTES AND NONELECTROLYTES
AH3
A
2 U6 0 0
4001
200
oL15
30
20-
35
t:"C
Figure 1. Enthalpies of solution of nitromethane.
4 3 0
20
t,"c
Figure 2. Heat capacity differences of nitromethane.
high values on both sides of a minimum near 24". The rapidly changing values of AC,," also show that great caution must be used in comparisons and interpretations of heat capacity differences as well as other thermodynamic and transport properties of nonelectrolytes in aqueous solution determined a t a single temperature (usually 25"). The minimum in the present study could be evidence for a structural transition in waterlOvll or for a change in the nitromethane-water interaction, both of which are indistinguishable by the present macroscopic thermodynamic measurements. The value of AC,, remains constant with temperature and except for the region between ea. 20 and 28" is
169
lower than AC,,". Since potassium chloride is a known structure breaker, l2 the aqueous electrolyte solution is less ordered than pure water, with possibly less structure making by nitromethane and consequently a lowered heat capacity change than in pure water. The added KC1 also appears to "smooth out" water structure changes and their effect on specific water-nitromethane interactions by substitution of much stronger and regularly varying specific ion-nitromethane and structure-breaking ion-water effects, with temperature apparently playing a smaller part than in pure water. Thus, the net behavior of the ternary nitromethanewater-KC1 system appears to be deceptively simpler than the binary nitromethane-water system. Measurements at lower KC1 concentrations may show the limit of its structure-breaking effects. Values of the absolute partial molal heat capacities in pure water, CpPo,or in 1 m KC1, C,,, may easily be obtained by adding the constant heat capacity of liquid nitromethane (25 cal/deg moll3) to AC,," and AC,,, respectively. I n contrast to C,,", the variation of the partial molal heat capacity of aqueous strong 1-1 electrolytes, Cp:, with temperature appears to be much smaller. For example C,," of NaCl increases from -28 to -19 between 10 and 4O0.l4 The heat capacity of transfer, E?,,, is negative (except near room temperature), and there are very few other aqueous systems for comparison. That for argon, based on solubilities in water and in 1-1 electrolytes is also negative, and with few exceptions this is generally true for alkanes from methane through butane.16 Thus, the transferred solutes appear to make less structure in the electrolyte solutions than in pure water. However, most of the Ep, decrease with temperature are monotonic and nearly linear. This may be due to the difference between polar nitromethane and nonpolar solutes, and due to the inaccuracy of the solubility method. A calorimetric study with undissociated acetic acid16 between 18 and 35" has been initiated.
Acknowledgment. The financial assistance of the National Science Foundation is gratefully acknowledged (Grants GP-7509 and GP-11271). (10) Benzene is another example of complicated ACpa0behavior; see F. Franks, M. Gent, and H. H. Johnson, J . Chem. rSoc., 2716, (1963). (11) For discussion of thermal anomalies in the structure of water, see W. Drost-Hansen, Advances in Chemistry Series, No. 67,American Chemical Society, Washington, D. C., 1967,p 70. (12) J. L, Kavanau, "Water and Solute-Water Interactions," Holden-Day, Inc., Ban Francisco, Calif., 1964,p 57. (13) The heat capacity of liquid nitromethane remains constant within ca. 0.5 cal/mol deg over wide ranges of temperature. See W. M. Jones and W. F. Giauque, J . Amer. Chem. SOC.,69,983(1947); J. W. Williams, ibid., 47,2644 (1925). (14) See ref 6,p 403. (16) H. L. Clever and C. J. Holland, J . Chem. En'ng. Data, 13, 411 (1968). (16) J. H.Stern, J. P.Sandstrom, and A. Hermann, J . Phys. Chem., 71,3623 (1967). Volume 74, Number 1 January 8, 1070
