NOTES
from the experimental data. The Eberhart S parameter is a first-order approximation to the more precise parameters as shown by the derivation of eq 4 from eq 2. Comparison of S to exp f illustrates this firstorder nature of S. The C parameter used in eq 2 has been estimated and is shown in column 8 of Table I. The surface area, 2, was calculated assuming a spherical molecule from the power of the liquid molar volume for the component with the lower surface tension. Comparison of expi [(a1 - ~ ) / R T ] zto ) the S parameter indicates that this surface area calculation amounts to about a first-order approximation for hydrocarbon systems. From the f parameter, one can calculate the.difference in the adsorption potentials of the two components a t the liquid-vapor interface. This difference is given in the last column of the table, The ability to calculate t,he adsorption potential difference of a third system from measurements on two related systems by eq 7 is fair in most cases and is limited by the uncertainties in the adsorption potential differences. Acknowledgments. The author wishes to thank Dr. H. L. Clever for his helpful discussions in preparing this note, the Emory Data Processing and Analysis Center for use of the computer, and Dr. Louis Homer for discussions concerning the nonlinear leastsquares program. This work was supported by National Science Foundation Grant GP-5937.
Solubilities of Alkyammonium Iodides in Water and Aqueous Urea by F. Franks' and D. L. Clarke Department of Chemical Technology, University of Bradford, Bradford 7, England (Received Noumber 1 1 , 1966)
Attention has recently been drawn to the effect produced by urea on the solubilities of hydrocarbons,2 micelle formation by surfactants,*and the behavior of amino acids4 in aqueous solution. The view is held that urea lowers the tendency of nonpolar groups to form hydrophobic bonds by some modification of the structural properties of water. Aqueous solutions of hydrocarbons no doubt provide the simplest systems in which these effects can be studied, but as the size of the solute molecule is increased, the solubility soon becomes so low that its temperature dependence must be subject to a considerable experimental uncertainty.
1155
One method of introducing reasonable concentrations of alkyl groups into aqueous solution consists of the employment of tetraalkylammonium halides. These compounds possess several properties which distinguish them from normal univalent electrolytes and in some respects they behave almost like hydrocarbons, e.g., they form interstitial clathrate hydrates.6 The hydrocarbon-like behavior becomes more pronounced as the size of the alkyl groups is increased and it was therefore believed that useful information regarding the influence of urea on hydrocarbon-water interactions might be obtained from a study of the solubilities of tetra-n-butyl- and tetra-namylammonium iodides. Experimental Section Solubility determinations in water and 7 M urea were carried out Over a temperature range 2-40", care being taken to avoid supersaturation of the solutionso Samples were taken in duplicate, weighed, and the iodide was titrated against silver nitrate with an automatic potentiometric method which enabled the rate of addition of silver nitrate to be varied according to the total amount of iodide in the sample. The a~kylammoniumiodides were recrystallized from ethanol-chloroform mixtures containing small amounts of water. Analysis showed them to be >99% pure. Since urea slowly decomposes on standing,6 all freshly made solutions were purified by ion exchange on a mixed column of Amberlite CG-400 and Amberlite IR-lBO(H). Results and Discussion The mole fraction solubilities of (n-Bu)rNI and (nAm)4NI in water and 7 M urea are shown in Table I. The log z ( T - l ) plots can be fitted satisfactorily by parabolic functions, showing that at low temperatures solution occurs exothermally. For the aqueous solutions, the enthalpy of solution, AH,, passes through zero at 12 and 16' for ( ~ - B U ) ~ Nand I (n-Am)rNI, respectively. For the solutions in 7 M urea, the minima in the log z (T-l) curves are displaced to much lower temperatures. The results indicate that for (n-Am)4NI this temperature is near 0'. The temperature de(1) To whom all correspondence should be addressed at Unilever Research Laboratory, Colworth House, Sharnbrook, Bedford, England. (2) D. B. Wetlaufer, S. K. Malik, L. Stoller, and R. L. Coffin, J . Am. Chem. Soc., 86, 508 (1964). (3) M. J. Schick, J . Phgs. Chem., 6 8 , 3585 (1964). (4) G. C. Kresheck and L. Benjamin, ibid., 68, 2476 (1964). (5) E.Q.,R. McMullan and G. A. Jeffrey, J . C h m . Phys., 31, 1231 (1959). (6) H. B. Bull, K. Breese, and G. L. Ferguson, Arch. Bwchem. Biophys., 104,297 (1964).
Volume 71, Number 4
March 1967
NOTES
1156
pendence of AHe shows that solution is associated with large, positive changes in the heat capacities of the systems, but the degree of accuracy of the solubility determinations makes a quantitative evaluation of AC,, values meaningless. ~
~~
Table I: Mole Fraction Solubilities of Tetracn-butylammonium Iodide and Tetram-amylammonium Iodide in Water (G) and 7 M Aqueous Urea (G) Temp,
-(n-Bu)dNI10'tw
O C
2.0 4.0 9.0 15.0 21.0 26.0 31.0 36.0 43.0
-(n-Am)tNI10%a
10""
1.06,1.02 2.01,2.01 1.01,1.02 2.09,2.11 0.99,0.98 2 . 2 5 , 2 . 2 5 0 . 9 7 , 0 . 9 9 2.43,2.50 1 . 0 7 , l . M 2.75,2.69 1 . 1 4 , 1 . 1 0 3.00,2.97 1.26,1.26 3.35,3.40 1.41,1.39 3.98,4.02 1.60,1.62 . . . ,
...
4.79,4.81 4.57,4.50 4.37,4.40 4.17,4.14 4.32,4.39 4.52,4.48 4.76,4.80 5.62,5.40 6.46,6.38
10%"
8.12, 7.98 8.08, 8 . 0 4 8.51,8.46 8.81,8.90 9.29, 9.18 10.4, 1 0 . 2 11.8, 11.7 13.2, 13.7 16.3, 16.0
The thermodynamic functions associated with the transfer at, 25" of 1 mole of solute from an aqueous medium to a 7 M urea solution are set out in Table 11. They reveal a type of behavior which has been reported for the hydrocarbons,2 Le., that the transfer of solute from water to aqueous urea is spontaneous (AGt" < 0) but endothermic (AHt" > 0), indicating that the process is entropy directed and might be interpreted in terms of solvent structure. I n contrast to the behavior shown by the hydrocarbons) AGto decreases in magnitude with an increase in the number of carbon atoms and is not as large as might be expected from a rough extrapolation of the AGt" Values auoted for the lower Daraffins. On the other hand. LHt" and AS,"are of the Same order of magnitude ab the corresponding hydrocarbon values quoted by Wetlaufer, et d 2 Unfortunately, there is at present not enough information regarding the alkylammonium halides to assess the magnitude of the ionic contributions to the transfer functions. Qualitatively, the positive AH," and A s t o results are-in accord with the concept that urea reduces the structuredness of the aaueous medium and hence the tendency of the alkyl chains to participate in the formation of hydrophobic bonds.' It is of course arguable
whether hydrophobic bonds can significantly affect the behavior of systems in which the mole ratio of solute to solvent is of the order 1-106. Partial molar volume studies on dilute solutions of (n-Bu)4NBr have indicated that this substance behaves as a normal univalent electrolyte a t concentrations up to about M (mole ratio 1: 6 X 103)8and that the volume anomalies first observed by Wen and Saitog only occur a t higher concentrations. There is, however, a more fundamental reason why the results cannot be interpreted solely in terms of a shift of the structural equilibrium in water, produced by the addition of urea, at least not on the basis of the currently available water structure models.lOfll For any shift in an equilibrium, the associated enthalpy and entropy changes must cancel in the free energy and hence such a process cannot affect the solute chemical potential, which alone is a measure of the solubility. By an extension of the mixture model"J*'l for water to aqueous solutions, it has been found possible to derive expressions for p z , R2, and % . of hydrocarbons in aqueous and 7 M urea solutions in terms of experimentally accessible quantities12 and it has been found that the experimentally determined AGtovalues of Wetlaufer, et a1.,2can be accounted for. It is intended to apply the treatment to the tetraalkylammonium halides after isopiestic and calorimetric measurements, now in progress on the lower members of the series, have been completed. (7) Although the view has recently been advanced [M. Abu-Hamdiyyah, J . Phys. Chem., 69, 2720 (i965)i that urea increases the degree of hydrogen bonding in water, in the authors' opinion this is incompatible with the physical properties of water-urea mixtures and with the effects here described. (8) F. Franks and H. T. Smith. Collected Abstracts: 16th CITCE Meeting, Budapest, Hungary, 1965, t o be published. (9) W. Y. Wen and S. Saito, J. Phys. Chem., 6 8 , 2639 (€964). (10) H. S. Frank and A. S. Quist, J . Chem. Phys., 34, 604 (1961). (11) G. NBmethy and H. A. Scheraga, ibid., 36, 3382 (1962). (12) H. S. Frank and F. Franks, presented a t the London Chemical Society Anniversary Meeting, Nottingham, England, April 1964.
Heats of Adsorption of Carbon Dioxide on Doped Zinc Oxide by 0. Levy and AI. Steinberg
Table 11: Free Energy (AGt"), Enthalpy (AHt"), and Entropy (A s t o ) of Transfer of Alkylammonium Iodides
Department of Inorganic and Analytical Chemistry, The Hebrew Un$eraity, Jerusalem, Israel (Received November 69, 1966)
from Water to 7 M Aqueous Urea at 25" (cal mole-') AGtO
(n-Bu)rNI (n-Am)rNI
-570 -492
The Journal of Physical Chemistry
AHt'
AStO
860 1240
4.8 5.8
Doping of zinc oxide results in change in the heat of chemisorption of carbon dioxide on its surface.' It
