THERMODYNAMICS OF ADSORPTION OF CARBON DIOXIDE ON ZINC OXIDE
pi" =
pM(so1n)
+ b(so1n)
(1-4)
and
17
we obtain
+
+
+
~ ~ ( s o l n )d s o l n ) = P M * ( g ) RT In aM(g) cc2*(1) RT In az(l) Ap"
where po refers to the pure condensed phase, and Apo is the standard free energy change of reaction 2. Equating eq. 1A and 2A and substituting for ~ M O PM"
:=
m*(g)
+ RT In aM(g)
(34
and for pzo PZ" =
+
PZ*(~) RT In a2(U
+
+
(5A)
Now substituting pt(soln) = pl*(soln)
+ RT In at(soln)
(6A)
into (5a) for ~ ~ ( s o l nand ) p2(soln), recognizing that p2*(1) = pz*(soln), and rearranging terms, we find pM*(soln) - C(M*(g) = RT In aM(g)/aM(soln)
RT In aMx,(l)/aMx,(soln)
(44
+ Apo
+ (7A)
Thermodynamics of Adsorption of Carbon Dioxide on Zinc Oxide
by R. J. Kokes and Rimantas Glemza Department of Chemistry, The Johns Hopkins Univereity, Baltimore, Maryland (Received December 69, 1969)
81618
Chemisorption of carbon dioxide on zinc oxide has been studied between 473 and 588'K. for pressures ranging from lo-' to 1 atm. Partial molal enthalpies and entropies of adsorption were computed by application of the Clausius-Clapeyron equation at fixed coverage; corresponding molar thermodynamic quantities were computed by application of the Clausius-Clapeyron equation at fixed spreading pressure. Experimental data show that the isosteric heat of adsorption a t low coverage is about 25 kcal. ; at coverages above half a monolayer, the isosteric heat is approximately equal to that for the formation of bulk zinc carbonate even though the bulk phase is thermodynamically unstable. Analysis of the data suggests that adsorption of carbon dioxide on zinc oxide results in the formation of surface carbonate groups.
Introduction Although a large number of studies of adsorption on semiconductor oxides have been carried out, very few of these deal with equilibrium adsorption. In this paper we present such data for the ZnO-CO2 system. This system was chosen for three reasons. (a) It yields reproducible data. (b) Bulk zinc carbonate (which is unstable under the conditions of our measurements) is a well-known compound, and it is possible to conipare the properties of the bulk phase with those of the
adsorbed phase. (c) Recent infrared studies' suggest the mode in which carbon dioxide is bound to the surface. Adsorption data for this system have been reported in the past12tabut these data are in conflict.
(1) J. H. Taylor and C. H. Amberg, Can. J . Chem., 39, 535 (1961). (2) P. M. G . Hart and F. Sebba, Trans. Faraday Soc., 5 6 , 557 (1960). (3) T. Kwan, T. Kinuyama, and I;. Fuiita, J . Res. Inst. Catalysis, Hokkaido Uniu., 3 , 31 (1953).
Volume 69,Number 1
January 1906
R. J. KOKESAND RIMANTAS GLEMZA
18
Experimental The carbon dioxide was prepared by fractionating the thermal decomposition products of sodium bicarbonate. The zinc oxide used in this experiment is the zinc oxide SP500 pigment manufactured by the Kew Jersey Zinc Co. Prior to each set of runs, the catalyst was degassed at 515’, a procedure known to reniove all but a trace of the residual carbon d i ~ x i d e . ~ Two different vacuum systems were used in these experinxnts, one covering a pressure range from to 10 mm. and the other covering a range from 1 to 800 mrn. The results in the low-pressure region were obtained for a sample of zinc oxide weighing approximately 5 g.; the results in the high-pressure region were obtained for a different sample of zinc oxide weighing approximately 25 g. Surface areas were measured by nitrogen adsorption on both samples. Wherever necessary, connections for thermal transpiration were made with equations suggested by Bennett and Tompkirm5 The temperatures for corresponding isotherms in the high- and low-pressure regions were the same to about il o , but the nominal variation in the highor low-pressure region alone was .t0.2’. Kormally, equilibriuni was achieved within an hour, except a t the lowest pressure and temperature for which data were obtained, Equilibrium was assured by the fact that data for adsorption and desorption agreed within experimental error. The raw data are presented in Figure 1 expressed in terms of the nominal 0, the ratio of the volume of carbon dioxide adsorbed to the B.E.T. um, and the measured pressure, i.e. , uncorrected for thermal transpiration. To avoid confusion, all of the experimental points are shown only for the data a t 200”. Calculations of the thermodynamic quantities are based on these smoothed curves. Calculations Partiul Molal Quantities. Isosteric heats of adsorption were computed by least-squares methods on the assumption that the data are represented by the following equation for a fixed 8
In p
=
-(“;R 1 + I=} s’-s 81-
T
I
R
,
I n this equation P is the equilibrium pressure (corrected for thermal transpiration) for a given e, R is the gas constant, T is the absolute temperature, H , - R, (E& is the isosteric heat of adsorption, R, and s’, are the molar enthalpies and entropies for C02gas, and H , and S,are the partial molal enthalpies and entropies for the adsorbed phase. This equation is valid provided the quantities in braces are independent of temperature in the range covered. The validity The Journal of Physical Chemistry
0
5
10
Ln Pb).
Figure 1. Isotherms for the adsorption of carbon dioxide on zinc oxide. The pressure, measured in M , is not corrected for thermal transpiration. Temperatures are A, 473°K.; B, 510°K.; C, 548°K.; D, 588°K.
of this approximation is attested to by the fact that the average variance for both the entropy and enthalpy terms was about 1.5% of the values, Le., 0.4 kcal. in enthalpy and 0.8 e.u. in entropy. Values of R, - R, are plotted in Figure 2. The least-squares variance for each point is indicated by a perpendicular line if its value is greater than the diameter of the point symbol. Values of S, were put on an absolute basis from fl, - S,with the help of the tabulated entropies for carbon dioxide in the “JAKAF Tables.”6 I n this computation it was assumed that the value of s’, - S, obtained from the leastsquares computations was that corresponding to T,’ 5283°K. A plot of 3, us. e is shown in Figure 3 ; the absolute variance for s’, - S, is indicated if its value is greater than the diameter of the point symbol. Molar Quantities. The data were adequate to permit computation of molar enthalpies, entropies, and energies of adsorption by the procedures outlined by HiL8 For the sake of completeness, these calculations were carried out. The error inherent in such calculations is uncertain because extrapolation to zero coverage is required. Details of this extrapolation will not be given here, but analysis of the error thus introduced suggests the error is less than 0.5 kcal. for the energy terms and ranges from 2 e.u. a t 0 = 0.1 to 0.4 e.u. a t e = 0.5 for the entropy. (4) R. J. Kokes, J . Phys. Chem., 66, 99 (1962).
(5) M .J. Bennett and F. C. Tompkins, Trans. Faraday Soc., 53, 185 (1957). (6) D. R. Stull, et al., “JANAF Interim Thermochemlcal Tables,”
Dow Chemical Co., Midland, Mich. (7) T. L. Hill, P. H. Emmett, and L G. Joyner, J . A m . Chem. Soc., 73, 5102 (1951). (8) T. L. Hill, J . Chem. Phvs., 17, 520 (1949); 18, 246, 791 (1950).
THERMODYNAMICS OF ADSORPTION OF CARBON DIOXIDEON ZINC OX ID^
I
I
0
0.40
0.20 9.
Figure 2. Heats of adsorption of carbon dioxide on zinc oxide. Solid points represent qst. Plain open circles represent molar enthalpies of adsorption computed from the Clausius-Clapeyron equation a t fixed spreading pressure.
19
herein do not correspond to bulk reaction. Estimates based on data a t 298°K.,g with the assumption that AH and AS do not vary appreciably with temperature, indicate that the dissociation pressure of zinc carbonate is 20 to 650 atm. in the temperature range studied. Since the highest pressures used were -1 atm. we can exclude formation of bulk zinc carbonate. The formation of solid solutions of zinc carbonate in zinc oxide is regarded as unlikely.) The over-all fall in isosteric heats of adsorption with increasing coverage is often associated with heterogeneity of sites.1°-12 Although in this present case there is a fall in qat from about 25 kcal. a t e = 0.1 to 16.5 kcal. a t 0 = 0.5, this does not seem to be the result of a heterogeneity of sites in the usual sense. In the statistical, mechanical treatment of heterogeneity it is assumed that different sites have different integral energies of adsorption, AB, a t O°K.la This results in a fall of A E with coverage a t 0°K. Values of A 8 a t 529°K. shown in Figure 4 are essentially constant a t 23 f 1 kcal. up to e 0.4. Thus, by the usual standards, the zinc oxide surface is reasonably homogeneous in carbon dioxide chemisorption. The absolute entropy values plotted in Figure 3 represent the value of the entropy of zinc oxide with adsorbed gas minus the entropy of zinc oxide in the absence of adsorbed gas. To compare this with the
-
\
40 ti
-I
3
\
>; 30 c
B
20 0
0.20 8.
0.40
0.60
Figure 3. Entropies of adsorbed carbon dioxide on zinc oxide. Open circ!es represent 3,; solid points represent SP,molar entropies.
Discussion Comparison of the isosteric heats in Figure 2 with those conflicting results2a3reported earlier in the literature show that these data are in essential agreement with those reported by Hart and Sebba.2 The combined data suggest that the isosteric heat levels off a t a value of about 16.5 kcal. The value of the enthalpy change ZnO(s) C02(g) is for the reaction ZnCOB(s) +17.0 kcal. at 298°K.9 If we make the crude approximation that the heat capacities do not vary with temperatures, we find the decrease in AH between 298 and 529°K. is less than 1 kcal. Thus, the AH for dissociation of ZnC03 a t 529°K. is probably about 16-17 kcal., i e . , essentially equal to the isosteric heat a t higher coverages. (It should be emphasized that the data reported,
-
+
0
Figure 4.
0.20
e.
0.40
0.60
The integral energy of adsorption
of carbon dioxide on zinc oxide.
(9) F. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine, and J. Jaffe, National Bureau of Standards Circular 500, U. S. Government Printing Office, Washington, D. C., 1952. (10) A. Clark and V. C. F. Holm, J . Catalysis, 2, 21 (1963): A. Clark; V. C. F. Holm, and D. M . Blakburn, ibid., 1, 244 (1962). (11) L. E. Drain and J. A. Morrison, Trans. Faraday Soc., 48, 316; 840 (1952); 49, 654 (1953). (12) D. H. Everett, ibid., 46, 453 (1950). (13) T. L. Hill, J . Chem. Phys., 17, 762 (1949).
Volume 69, Number 1
January 1965
R. J. KOKESAND RIMANTAS GLEMZA
20
bulk phase, we must compare these entropy values to that of zinc carbonate minus that of zinc oxide. At 298°K. this value is 9.2 e.u.,2but the necessary thermodynamic data a t 529°K. are not available. A crude estimate of entropies a t 528.8”K. can be made if we assume the heat capacities are independent of teniperature. This yields a value of 14.7 e.u. for Szncos SZ,,O a t 528°K. If this is subject to the same sort of error as in the similar AlgC03-RIg0 system,6 a better estimate for this quantity is 17.5 e.u. Comparison with Figure 3 shows that even near $ = 0.5 (where the contributions of the configurational entropy should be the least) the entropy of the adsorbed phase is about 10 e.u. greater than that for the bulk phase. Let us assume for simplicity that the fall in the heat of adsorption is due to interaction between adsorbed carbon dioxide molecules. Then & a t $ 0 5 (27 e.u.) depends primarily on the mode of adsorption of carbon dioxide on a bare surface. The entropy of’ gaseous carbon dioxide a t 529°K. is 56.7 e.u.‘j We can use the conventional approaches of statistical thermodynamic^^^ and data given in ref. 6 to divide the entropy into rotational, vibrational, and translational entropy, The two degrees of rotational freedom contribute 14.2 e.u., the four degrees of vibrational freedom contribute 2.4 e.u., and the three degrees of translational freedom, 40.1 e.u. On adsorption without .change in geometry the three translational degrees of freedom are replaced by three center-of-mass vibrational degrees of freedom; moreover, although the internal vibrational degrees of freedom will be little affected, on adsorption the rotational freedom will probably be curtailed. The center-of-mass vibrational entropy can be estimated by a variety of means, but for chemisorption it is probably small, L e . , -2 e.u. If this be so, we find an entropy for adsorbed carbon dioxide between 4.4 (no rotation) and 18.6 e.u. (free rotation), values far lower than that actually observed. Infrared datal suggest that adsorbed carbon dioxide is riot linear but bent with one of the structures
-
0-0
0
0
~\ ,/
‘\,?
C
C
I
S
I
1
-
0
I
AI
The first species is postulated because the spectrum of the strongly held carbon dioxide closely resembles that obtained for carboxylate ions; the nature of the The Journal of Physical Chemistry
surface atom is unspecified. The second species is regarded as possible, but not as likely, because the observed spectrum does not quite agree with that found for bicarbonate ions. The latter structure corresponds approximately to a surface carbonate ion. If such a species does form, there is the possibility of essentially free rotation around the 11-0-C axis, a mode of motion that is unlikely for the bulk carbonate. The entropy contributed by this type of motion can be coniputed from the equation Srot= R/2 In (nTe/a2$,)15s’6 wherein all symbols have their conventional meaning. If this is the principal difference between the surface and bulk carbonate, we find that the entropy of the surface carbonate will be 7.5 e.u. higher than that of bulk carbonate, corresponding to an entropy for the adsorbed phase of 25 e.u., in fair agreement with the observed value of 27 e.u. Although the above analysis is undoubtedly oversimplified, it does suggest an attractive picture for the adsorption of carbon dioxide on zinc oxide involving the formation of a surface carbonate. For convenience, let us compare adsorption and bulk carbonate formation. At low coverages, $ < 0.5, adsorption is more favorable because of the higher isosteric heat and the higher entropy term. The higher heat term occurs in this region because an isolated carbonate group is more firmly bound than one interacting with its neighbors; the higher entropy term presumably results from the rotation of the carbonate group and configuration effects. At $ 0.5 the slope of In P us. 1 / T plots is the same as that for the bulk reaction (the isosteric heat of adsorption is the same as that for the bulk phase reaction) ; the more favorable entropy term is the only factor that makes adsorption occur under conditions that no bulk reaction occurs. Data are not available for higher coverages. In line with the above, it would appear that a t $ 2 1 the isosteric heat of adsorption is still that of the bulk reaction. At this point, moreover, the entropy approaches that for bulk reaction since the high surface coverage would curtail free rotation of the surface carbonate group. Accordingly, plots of In P us. 1 / T would be the same for bulk phase reaction and adsorption because the partial molal properties are the same; this would occur even though the integral heats of reaction were different for the bulk and surface reaction.
(14) T. L. Hill, “Statistical Thermodynamics.” Addison-Wesley Publishing Co., Inc., Reading, Mass., 1960, pp. 164-166. (15) This yields a value different from that obtained using the general relation suggested by Kemball. 18 The difference in partition function is a factor of ?r, < . e . , a difference of 2 e.u. (16) C . Kemball, Aduan. Catalysis. 2 , 233 (1950).
KINETICS OF MONOMOLECULAR FILMS OF CHLOROPHYLL a
Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered
21
by the American Chemical Society, for support of this research.
Reaction Kinetics of Monomolecular Films of Chlorophyll a on Aqueous Substrates
by Morton Rosoff and Carl Aron I B M Watson Laboratory, Columbia University, New York, NEW York
(Received February 8 , 1964)
The kinetics at constant area of the monolayer reaction of chlorophyll a to pheophytin a on an acidic substrate were studied. The rate constants were found to depend upon pH, the initial surface pressure, the presence of O2at the interface, and divalent metal ions such as Ca+*and YIg+2in the subphase. The half-life of chlorophyll a at 23' on a pH 4 substrate and at an initial surface pressure of 6 dynes/cm. was about 6 min. Evidence was obtained for the participation of H 2 0 in the reaction and the probable existence of a charged intermediate.
Introduction of chlorophyll spread a t an Film balance aqueous interface afford a unique way of approximating a molecular state of organization and environment which is closer to that present in vivo than the one which exists in simple solutions of photosynthetic substances in organic solvents. In the course of investigating the stability, reproducibility, and physical properties of chlorophyll monol a y e r ~it, ~was found that pheophytin was the primary product of chlorophyll decomposition a t the surface. This work reports on the utilization of the techniques of monolayer reactions to study the kinetics of the conversion of chlorophyll a to pheophytin a in the presence of H + ion. Since the relationship between surface properties and time depends upon the interaction between reactant and product, the properties of mixed monolayers of these substances were also examined.
Experimental The preparation of pigments, the apparatus, arid spreading techniques used in these experiments have
been described elsewhere. Briefly, the chlorophyll a was prepared from fresh spinach leaves by the method of Jacobs, et aL6 Pheophytin was formed from chlorophyll by the addition of HC1 and rechromatographing. The surface balance was of a semiautomatic Wilhelmy type employing a Teflon trough and was sensitive to 0.1 dyne/cm. Surface potentials were measured by means of a Ra226ionizing source and a Keithley 6lOH electrometer. Concentrations of solutions of pigments were obtained from measurements of optical density and previously determined extinction coefficients. Solute was delivered to the substrate surface by means of a Hamilton microliter syringe using benzene as the M , was used carrier solvent. Phosphate buffer, (1) A. E.Alexander, J. Chem. Soc., 1813 (1937). (2) H. J. Trurnit and G. Colmano, Biochem. Bwphys. Acta, 36, 447 (1958). (3) W. 0.Bellamy, J. L. Gaines. Jr.. and A. G. Tweet, J . Chem. Phys., 39, 2528 (1963). (4) M.Rosoff and C. Aron, hrature, to be published. ( 5 ) E. E. Jacobs, A. E. Vatter, and A. S. Holt, Arch. Biochem. Biophys., 53, 278 (1954).
Volume 69, Number 1 January 1965
