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 effectshere described. (8) F. Franks and H. T. Smith. Collected Abstracts: 16th CITCE Meeting, Budapest, Hungary, 1965, to 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
NOTES
is of great practical importance that data on the activity of catalysts be obtained not only through methods based on measurements of adsorption properties under vacuum.2 Flow methods where conditions of adsorption resemble those dominating industrial processes are desirable. The transient response technique (TRT) probably holds the best promise for the fulfillment of some of the above requirements, since it provides means of measuring heats of adsorption in flow systems, and was used in the present study for the measurement of heat of adsorption of carbon dioxide on zinc oxide preparations of varying semiconducting character. Doped zinc oxides both of n and p type were used here. The TRT was normally applied to physical adsorption. To the best of the authors’ knowledge, it has not hitherto been used to determine heats of adsorption where the interaction is primarily chemical.
Experimental Section The pulse injection technique and measurement of the corrected retention times were as described previou~ly.~~~ Calibration of the desorption peak areas were carried out as foliows. Various volumes of carbon dioxide corrected to STP were injected into a column packed with silica gel. The pulses were injected at temperatures matching those of the chemisorption runs (Table I), Linear relationships were obtained between the carbon dioxide injected and the peak areas. From the calibration curve it was found that the desorptions of carbon dioxide from all the zinc oxide preparations used at the working temperatures were quantitative. The columns in all the runs were flushed with the carrier gas, helium (Air Reduction Co., stated purity 99.99a/,) for 113 hr at 370”. This treatment was found to give reproducible results. Zinc oxide was ’ lithprepared by the oxalate method.’ The 1 mole % iated zinc oxide was prepared by the method described by Cimino, Molinari, and Cramarossa.6 The doped oxide so obtained was yellow even when left in air at room temperature, and its apparent density was 6.5 g/cm3. The 1 mole % galliated zinc oxide was prepared by the method described by Hart and Sebba.’ The color of this preparation was white.
Results and Discussion The logarithms of the corrected retention times were plotted against the reciprocals of the absolute t e m p e r a t ~ r e . ~The retention times were independent of the flow rates. In this study, flow rates of 7-10 ml sec-1 were taken. These rates were found optimal for reasonable elution times. The elution curves for
1157
Table I : Heats of Adsorption of Carbon Dioxide on Various Zinc Oxide Preparations
Abaoi bent
ZnO
ZnO
+ 1 mole % Li
Pulse volume, ml
Temp,
kcal
OC
mole-’
0.5 1.0 1.5
280-310 315-380 320-365
30.8 28.5 26.5
0.1
340-360 320-360 320-360
25.3 21.6 19.5
337-360 332-360 337-360
36.8 32.5 28.7
0.15 0.25 ZnO
+ 1 mole % Ga
1.0 2.0 3.0
AH,
the same flow rate in the doped and undoped zinc oxide were similar, thus indicating the same type of sorption isotherm. The relationships obtained were linear and the calculated heats of adsorption are presented in Table I. It is clearly seen that the heats of adsorption (AH) of carbon dioxide on lithiated zinc oxide are lower than those obtained for undoped and galliated zinc oxide. The highest AH values obtained were for the galliated zinc oxide. The same trend was observed by Hart and Sebba,l using a different method of measurement. The results obtained in this work are in correspondence with the semiconducting character of the oxides under study here. Carbon dioxide, being an electron acceptor, is adsorbed more strongly on n type than on p type surfaces of semiconducting oxides. The heat of chemisorption on the p type lithiated zinc oxide is the lowest of the zinc oxide preparations in this study. The n type galliated zinc oxide and the undoped zinc oxide will tend to absorb carbon more strongly than the p type, and hence the higher values for the heats. It is also observed here that the AH values decrease with the increase in the gas pulse volumes. Preliminary experiments on the adsorption of carbon monoxide on the surface of the above preparations at the same temperatures (Table I) were discontinued. The adsorptions in our flow method were irreversible. The TRT is not suitable for the AH (1) P. N. G. Hart and F. Sebba, Trans. Faraday SOC.,56, 551 (1960). (2) S. J. Gregg, “The Surface Chemistry of Solids,” Reinhold Publishing Corp., New York, N. Y.,1961,pp 122-131. (3) J. J. Carberry, Nature, 189, 391 (1961). (4) 8. A. Greene and H. Pust, J . Phys. Chem., 62, 55 (1958). (5) R. S. Hansen, J. A. Murphy, and T. C. McGee, Trans. Faraday SOC.,60, 597 (1964). (6) A. Cimino, E. Molinari, and F. Cramarossa, J . Catalysis, 2, 315 (1963)
Volume 71.Number .G March 1967
NOTES
1158
calculations for this system at the temperature region 280-360'. The fairly good agreement between the results obtained in our study and those from adsorption isostersl shows that the application of the transient response technique to reversible chemisorptions is worth studying, being a quick technique and pertinent to practical uses.
over the frequency range 60 kc-10 Mc using a Boonton Radio &-Meter (Type 260-A) and a General Radio Schering-type capacitance bridge (Type 716C). The comparison of both sets of data is shown in Table I.
Acknowledgment. The authors thank Dr. F. S. Stone of the Department of Physical Chemistry, The University of Bristol, for his interest in this work.
This work
Density, Viscosity, and Dielectric Constant of
Table I
Density, pZ50 d In v/dTb Viscosity log t) Dielectric constant, c d In c/d In T
Szwarc'
0.883 0.00133 3.670 395/T 1.50 2650/T -1.19
-
+ +
" Determined from -70
to 25".
*v
0,880" 0. 001085" - 3.655 393/T" - 1.495 2659/T" -1.16"
+ +
= molar volume.
Tetrahydrofuran between -78 and 30°1 by Donald J. Metz and Althea Clines Brookhaven National Laboratory, U p t m , New York 11979 (Received November 1 , 1966)
Recently, Szwarc, et aLla published data on the density, viscosity, and dielectric constant of tetrahydrofuran between -70 and 25'. In an independent and simultaneous study, we determined these properties between -78 and 30". We would like to report that our data substantiate and confirm those previously published by both Szwarc2and We determined our densities pycnometrically, using dried and purified ethyl bromide as standard.6 Viscosity was measured in a modified Ubbelohde viscometer us. both ethyl bromide6 and diethyl ethere6 Dielectric constants were determined, us. ethyl bromide,"
T h Journal of Physical Chemistry
With the exception of the temperature dependence of density, the agreement between both sets of data is excellent. When the data of Kuss3 on the variation of density in the region of room temperature are included, there is a suggestion that the larger coefficient of expansion may be more nearly correct.
(1) This work was performed under the auspices of the U. 9. Atomic Energy Commission. (2) C. Carvajal, K. J. Tolle, J. Smid, and M. Szwarc, J . A m . Chem. Soc., 87, 5548 (1965). (3) E.Kuss, 2.Angew. Phys., 7, 376 (1955). (4) F. E.Critchfield, J. A. Gibson, Jr., and J. L. Hall, J. Am. Chem. Soc., 7 5 , 6044 (1953). (5) "International Critical Tables." (6) A. A. Maryott and E. R. Smith, National Bureau of Standards Circular 514,U.S. Government Printing Office, Washington, D. C., 1951.
