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T.VRBASKI
Kinetics of Processes Occurring on the Catalyst Surface during the Oxidation of o-Methylbenzyl Alcohol over Vanadia
by T. Vrbaiiki Sinclair Research, Inc., Harvey, Illinois 60.4b6 (Recsiwd April 6, 1966)
~~
Rates of oxidation (disappearance) of o-methylbemyl alcohol were measured in the temperature range from 300 to 350" using a fused vanadium pentoxide catalyst. The Hinshelwood treatment suggests that the steady-state conditions are established on the catalyst surface. That the specific rate constant for the adsorption of oxygen agrees within one order of magnitude with those reported in the literature for the oxidation of other organic compounds over the same catalyst supports this conclusion. In this treatment it is supposed that only oxygen is chemisorbed on the catalyst surface; the rate of adsorption of oxygen and the rate of the chemical reaction are of the same order of magnitude; all the processes are first order in the concentration of each reactant, and the rate of desorption of oxygen from the catalyst is negligible. Application of the Hughes-Adams expression implies that the o-methylbenzyl alcohol is first chemisorbed on the catalyst and then reacts with the oxygen of the catalyst; the oxidation products desorb, and finally the catalyst is very rapidly reoxidized. The chemisorption of the o-methylbenzyl alcohol is probably a reversible process.
Introduction A study of the vapor phase oxidation of o-methylbenzyl alcohol over vanadium oxide as catalyst was recently carried out in this laboratory.' This study clarified the reaction mechanism and provided, in addition to kinetic data, clues for the interpretation of processes which occur on the catalyst surface. o-Tolualdehyde was found to be the main product in the temperature range between 280 and 350". Carbon oxides along with o-toluic acid and small quantities of per-o-toluic acid and phthalide were also formed. No correlation was made, however, of the rate data of disappearance of o-methylbenzyl alcohol in terms of kinetic expressions which were previously found by others to be useful in describing catalytic oxidation processes. Kinetic studies of naphthalene, toluene, and benzene in the presence of vanadium oxide as catalyst were made by Shelstad, et aL12Downie, et aL13and Hayashi, et aL4 The rate equation based upon a reaction scheme proposed by Hinshelwood5 provided good characterization of the oxidation rates of individual compounds. Alternatively, Hughes and Adams6 developed an The, Journal of Physical Chemistry
expression for unimolecular surface reactions which proved useful in the interpretation of data obtained in their study on the vanadia-catalyzed vapor phase oxidation of phthalic anhydride. Using this expression they determined the heats of activation for both the irreversible adsorption of phthalic anhydride and the desorption of the oxidation products. In addition, they were able to calculate the heat of the reversible adsorption for phthalic anhydride, along with the entropy change for the reversible process. In view of the potentially useful information which comparative studies may provide, it was of interest to attempt to apply both treatments to the oxidation of o-methylbenzyl alcohol. In the present communica(1) T. VrbaZki and K. W. Mathews, J . Phya. Chem., 69,457(1965). (2) K.A. Shelstad, J. Downie, and W. F. Graydon, Can. J. Chem. Ena., 38, 102 (1960). (3) J. Downie, K. A. Shelstad, and W. F. Graydon, ibid., 39, 201 (1961). (4) R. Hayashi, R. R. Hudgins, and W. F. Graydon, ibid., 41, 220 (1963). (5) C. N. Hinshelwood, "The Kinetics of Chemical Change," The Clarendon Press, Oxford, 1940, p. 207. (6) M. F. Hughes and R. T. Adams, J . Phya. Chem., 64,781 (1900).
KINETICS OF OXIDATIONOF O-METHYLBENZYL ALCOHOL OVER VANADIA
tion the corresponding rate constants for the oxidation of a-methylbenzyl alcohol in both the Hinshelwood and the Hughes-Aciams expressions were determined. These data were correlated with available literature values for naphthalene, toluene, benzene, and phthalic anhydride, and the nature of the reactions occurring on the catalyst surface was discussed.
Experimental The reaction system, the method of analysis, and the general experimental procedure used in this study were described in a previous communication.'
Results In the simplest steady-state treatment by Hinshelwood it is assumed that only oxygen is adsorbed on the catalyst surface, that its rate of adsorption and the rate of the chemical reaction are of the same order of magnitude, that the rate of desorption of the oxygen from the catalyst surface is negligible, and that all the processes are first order in the concentration of each reactant. According to this interpretation, the rate constant for the adsorption of oxygen should be independent of the organic compound employed at a given temperature and in the presence of the same catalyst. The rate of adsorption of oxygen should, therefore, be the common step for all reactions. If N is the number of moles of oxygen consumed per mole of o-methylbenzyl alcohol converted, then at the steady state the rate of adsorption of oxygen is equal to the rate of chemical reaction multiplied by the factor N. Thus kaCo(l - 01) = NkrC$l rr =
krC&
(1) (2)
Alternatively, at a steady-state situation, the rate of reaction of oxygen is also equal to the rate of oxidation of the organic species multiplied by the factor N
k,CrA = NkrCrOi (3) Solving (1) for 01 and substituting it into (2) the rate of a-methylbenzyl alcohol oxidation is
(4) Inverting and also keeping the oxygen concentration constant, eq. 4 becomes 1_ - _1 _1 k r Cr
Tr
+ KI
(5)
where KI is equal to N/k,Co. The values of the rate constants in eq. 5 were calculated from experimental points given in Figure 7 of
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ref. 1, by the general method of regression analysis. The results are listed in Table I. The solid lines in Figure 7 of ref. 1 were calculated from eq. 4 by using the values for the rate constants listed in Table I. The scatter of experimental points at 340 and 350' is probably due to the slightly higher temperature gradient in the catalyst bed under these conditions. The ratio N is the average of individual values determined from the product distribution data in each series of experiments at constant temperature. The logarithms of IC, and kr are linear functions of the reciprocal absolute temperature. The calculated Arrhenius temperature coefficient of k, is 28.0 kcal./mole, and that of kr is 9.4 kcal./mole.
Table I: Constants of Eq. 4 and 5" Temp.,
l/S, g. aec./l.
Ki X 10-6, g. sec./mole
ka X 108,
OC.
300 310 320 330 340 350
2993.9 3005.5 2792.5 2264.9 2037.9 1675.0
109.8 72.6 49.3 34.7 23.2 14.8
0.44
a
l./g. aec.
i+ X 104, l./g. sec.
3.34 3.33 3.58 4.42 4.91 5.97
0.67 0.99 1.40 2.10 3.29
Average Co = 9.3 X 10-8 M. Average N = 4.53.
In the Hinshelwood expression a firsborder dependence of the rate on the concentrations of both the oxygen and the organic reactant is assumed. Experimental data,'-&? however, show that the reaction orders vary from 0.5 to unity depending on the compound employed. In fact, only the concentration dependence for xylene in the oxidation of a-xylene and that for oxygen in the oxidation of naphthalene were found to be first order?12 The adoption of an arbitrary firsborder rate dependence, as proposed by Hinshelwood, was found useful by various authors in correlating the results obtained in different system^.^-^ It is therefore believed that the adoption of the same h b o r d e r rate dependence in this study was permissible although not absolutely correct in view of the experimental results in ref. 1. The k. values of o-methylbenzyl alcohol determined from a modified Hinshelwood expression in which the dissociative adsorption of the oxygen is taken into account (K1 = N/kaCo'/' ~
(7) H. Clark, G. C. Serreze, G . L. Simard, and D. J. Berets, unpublished results of the American Cyanamid Co. presented at the Gordon Research Conference on Catalysis, June 1956. See J. I(.Dixon and J. E. Longfield, "Hydrocarbon Oxidation in Catalysis," Vol. VII, P. H. Emmett, Ed., Reinhold Publishing Corp., New York, N. Y . , 1960, p. 183.
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T. V R B A ~ K I
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in eq. 5 ) were found to be about one-tenth of those listed in Tables I and 11. The E, values, however, remained unaffected.
0 W IJ.
4
$ 0
Table II : Constants from o-Methylbenzyl Alcohol, Benzene, Toluene, and Naphthalene Oxidations over Vanadium Pentoxide Catalyst a t 350" .k X 106, & X 108, E., kcal./ Er, kca1.l
Reactant
N
o-Methylbenzyl alcohol 4.53 Benzene",' 3.07 Benzene' 3.07 TolueneQsd 1.35 NaDhthalene" -2.7
L/g. 8ec.
3.3 3.3 4.4 13.9 9.3
sec.
mole
mole
59.7 0.9 12.8 23.8 1700'
28.0 7.7 8.7 29.4 28.1
9.4 27.2 28.1 26.4 24.6
l./g.
Downie3 and Hayashi' observed an appreciable initial cat% lyst deactivation in the oxidation of toluene and benzene, respectively. Therefore, their data were based on a catalyst activity at a i age of 20 hp. 'See ref. 4. See ref. 9. dSee ref. 3. * See ref. 2. The kr value for naphthalene differs from that reported for toluene by a factor of 70 (ref. 3).
'
the moduct distributions from which the N values are determined were available. The rate data were also correlated by the expression proposed by Hughes and Adams for surface reactions in which the rate-controlling step is either the adsorption of the organic reactant or the desorption of the oxidation moducts. The assumed reaction mechanism in their study of phthalic anhydride oxidation involves adsorption of phthalic anhydride followed by reaction with the oxygen of the catalyst, desorption of the oxidation products, and very rapid reoxidation of the catalyst. In this treatment two limiting cases are considered: (i) irreversible adsorption when k-I = 0 and (ii) equilibrium adsorption when k-l >> kz. The integrated rate equation is 2.3031 log (1
- Z) - B p a
--At
(6)
in which the conversion is given in terms of the initial partial pressure of organic reactant (po), the contact time (t), and the rate constants A and B. Four isotherms were obtained from experimental data and plotted in Figure 1 M conversion of o-methylbenzyl alcohol against initial pressure of o-methylbensyl alcohol in the gas mixture at a constant oxygen concentration and 0.6-sec. contact t h e . The values for The Joural of Physical Chsmdstry
L
W
J
0
I
I
I 4.0
I
1.0 2.o 3.O f NlTlAL OMBA PRESSURE (ATM.x to3)
I
I
Figure 1. Conversion of o-methylbenzyl alcohol (OMBA) as a function of the initial partial pressure of the OMBA in the gas mixture. Average oxygen concentration 9.3 X M and contact time 0.6 sec.
Table III: Constants of Eq. 6O A / B , atm.
Temp.,'C. 300 310 320 330
A,
8ec.-1
1.6 1.8 2.2 3.6
B , atm.-1
1580 1169 983 1227
aec.-1
x
10'
9.9 15.1 22.2 29.4
' Contact time 0.6 sec.
The same authors showed that the constants A and B in eq. 6 were equal to klkdc'/(k-l k2) and kl/(k-l kz), respectively, in a rate expression derived by applying a steady-state treatment to reactions occurring on the catalyst surface. The constants k , k, and kz are specific rate constants of adsorption and desorption of organic reactant and desorption of oxidized reactant, respectively. The conversion factor L' changes the fraction of the surface occupied by the organic reactant to its true surface concentration. The ratio A / B consequently equals kzk' and is independent of the nature of the adsorption step. The Arrhenius plots of constants A , B , and A / B
+
+
KINETICSOF OXIDATION OF 0-METHYLBENZYL ALCOHOL OVER VANADIA
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using data reported by Ioffe and Lyubarskig in their study on the vanadia-molybdena-catalyzed oxidation of benzene. Both the rate constants k, and k, a t 350" and activation energies E, and E, were calculated. These results compare well with those reported by Hayashi except for the k, value. The reasons for this discrepancy are not fully clear although the difference in the catalyst and potential dficulties inherent in comparing results of such complex reactions can, at least in part, account for it. The observed difference, however, appears to be of little significance, for it does not affect the consistencyof the k, values. The k. value for the oxidation of o-methylbenzyl alcohol obtained in the present study compares with the corresponding data reported for other aromatic compounds although there are some differences between values obtained by various workers for benzene, toluene, and naphthalene. According to the assumptions made above, k. should be independent of the organic compound employed at a given temperature and in the presence of the same catalyst. The adsorption of oxygen should be the common step for all reactions. It is not possible to say at present how meaningful these differences are until information on the rate constant k, of a greater number of organic compounds becomes available. However, it appears that the k. values for the organic species vary within a Discussion relatively reasonable range and yet are within the same The Hinshelwood Treatment. In the Hinshelwood order of magnitude. rate equations (4 and 5) it is assumed that in heteroOn the other hand, data in Figure 7 of ref. 1show that geneous catalytic reactions where the temperature is the rate of disappearance of o-methylbenzyl alcohol is relatively low the rate of adsorption of the reactant is dependent upon the alcohol concentration in the carrier of the same order of magnitude or even smaller than gas at low values of the alcohol concentration and at the rate of the chemical reaction (r, 9 rr). It is noted the higher temperatures. At higher concentrations that Taylors reported a slow adsorption on the catalyst and a t lower temperatures the rate tends to become surface under these conditions. Consequently, the independent of the concentration of the alcohol. The values for k, should be in this case independent of order of reaction was found to be 0.48, and the activathe organic compound employed if the same catalyst tion energy for the disappearance of alcohol, 20.0 kcal./ is used. mole. The most logical explanation of this result is The Langmuir approach, in which the concentrations that the o-methylbenzyl alcohol undergoes oxidation in the gas phase are related to the surface concentrapredominantly in the adsorbed state, and its rate of tions (adsorption equilibrium), proved to be most adsorption on the catalyst surface tends to control successful only for high- temperature reactions in which under given conditions the over-all oxidation rate. red-hot metal wire catalysts were employed. In this If this is so, the Hinshelwood assumption that only treatment a necessary requirement is that r, and raf oxygen is adsorbed on the catalyst surface does not be much greater than r,. appear to be completely correct. However, even with In Table I1 a summary of data obtained from the the organic compound adsorbed, the steady-state present study and those reported by other a ~ t h o r s ~ - ~approach will remain valid if rat 7 r,. This modiiicaor calculated from kinetic results found in the literation seems to be permissible, for it is inherent to the tureg is presented. Although the k, values for toluene and naphthalene at 350" differ only by 30%, Hayashi's (8) H. S. Taylor, J. Am. Chem. SOC.,53, 578 (1931). value for benzene4 is about one-fourth that for (9) I. I. Ioffe and A. G. Lyubarski, Kinetics Catalysis (USSR), 3 , 223 toluene. Further information was obtained by (1962). listed in Table I11 are straight lines. The slope of the line for A / B gives the heat of activation for the desorption of oxidized o-methylbenzyl alcohol AHz* = 31 kcal./mole. In the case of irreversible adsorption of o-methylbenzyl alcohol, the constant k-I is zero and A = k&'. The temperature coefficient of A gives the heat of activation of the irreversible adsorption of o-methylbenzyl alcohol AHl* = 15 kcal./mole. In the case of equilibrium adsorption and assuming a slow desorption of oxidation products from the catalyst (kz is small) B = k&~, which is the adsorption equilibrium constant K. The slope of the line for B in the Arrhenius plot represents the heat of adsorption of o-methylbenzyl alcohol A H 1 = -16 kcal./mole. Finally, by using the conventional thermodynamic equation which combines the equilibrium constant K with AF, AH, and AS, the entropy change A& for the reversible adsorption of o-methylbenayl alcohol was found to be - 13 cal./(mole deg.). It is not possible to interpret the dependence of the rate upon oxygen pressure with the Hughes-Adams expression in its present form. This could possibly be done by developing a modified equation in which the above relationship would be included. However, more experimental data than those presently available (Figure 10 of ref. 1) would be required.
Volume 69, Number 9 September 1966
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T. VRBASKI
steady-state concept as outlined above. Furthermore, AHz*, AHI, AHz+, MI, and AF1 obtained from the it was found that k., > .k However, under steadypresent study and those reported for phthalic anhystate conditions and also depending upon the actual dride is presented. Although the values for o-methylconcentrations of the components in the gas mixture, benzyl alcohol agree in their main outlines with those ra may become greater than ra/. This may well exfound for phthalic anhydride, there are some significant plain why a large excesa of oxygen is always required dfierences. No conclusion, however, can be drawn in the reaction mixture during the oxidation process. upon the exact nature of the initial adsorption step. Data on the isotope exchange of Ole with CO on This is because the thermodynamic values appear oxidic catalysts10 indicate that the lattice oxygens of reasonable for both the assumed irreversible and the the catalyst do not participate to a larger extent in equilibrium adsorption of o-methylbenzyl alcohol. the process of oxidation of hydrocarbons at low temperatures. Under these conditions, when the stepTable IV: Thermodynamic Quantities Derived from the wise oxidation of o-methylbenzyl alcohol is taking Experimental Rate Constants A and B place, the reaction probably occurs mainly on account of the loosely chemisorbed oxygens on the surface. Irreversible -Equilibrium adsorptionAt higher temperatures when the direct oxidation adsorption AS>, AR1*, AH,*, AH1, AH?,+ cal./ process is taking place, the top layers of the catalyst kcal./ (mole AF1, kcal./ kcal./ kcal./ were shown to participate in the reaction.lrll This Reactant mole mole mole mole deg.) cal./mole result and the evidence on the adsorption of 0-methyl o-Methylbenzyl benzyl alcohol make it questionable that the Hinshelalcohol 15 31 -16 31 -13 -8100 at 320' Phthalic wood treatment is applicable to the conditions of the anhydride 11 44 -33 44 -29 -9240 at 542" direct oxidation pro cess. Inspection of the activation energy data in Table I1 shows that the E, value for 0-methylbenzyl alcohol (a) Irreversible Adsorption. The heat of activacompares with those reported for toluene and naphtion for the irreversible adsorption of o-methylbenzyl thalene, whereas the value for benzene is consistently alcohol (AH1* = 15 kcal./mole) compares reasonably three to four times lower. The r e w n s for this diswell with the value reported for phthalic anhydride crepancy are not fully clear. A tentative explana(AHl* = 11 kcal./mole). The Arrhenius activation tion, however, is that the catalytic sites which particienergy for the formation of o-tolualdehyde from opate in the oxidation of benzene are different from those methylbenzyl alcohol, calculated from the speciiic in the oxidation of other compounds studied. Hayashi4 The rates, was found to be 30.4 kcal./mole.' came to a similar conclusion in his study of the catacomparison of this value with AH2* = 31 kcal./mole lytic deact,ivation in the oxidation of benzene and for o-methylbenzyl alcohol appears to be permissible toluene. Further experimental work will be necessince o-tolualdehyde was found to be the principal sary to clarify this point. oxidation product in the temperature range frqm 300 Experimental data reported by Hughes and Adams6 to 350". The observed Arrhenius activation energy in their study on the vanadia-catalyzed oxidation of therefore represents either the heat of activation for phthalic anhydride between 472 and 575" were used to the desorption of o-tolualdehyde, A H 2 * , or that for calculate the activation energies E. = 36.4 and ET = the adsorption of o-methylbenzyl alcohol, AH1*, 22.6 kcal./mole. Furthermore, the value E, = 24.6 whichever is higher, unless the reaction is solely conkcal./mole for naphthalene was determined by using trolled by the adsorption of oxygen. the available data from Shelstad's work2and approxiIn an attempt to obtain information on the extent mating the k, value at 350" as 17.0 X (ref. 3). to which the presence of by-products such as carbon The results agree well with those reported for the oxides, maleic and phthalic anhydrides, and o-toluic majority of the other compounds studied and are acid in the oxidation product affects the AHz* value presented in Table 11. for o-methylbenzyl alcohol, plots of the rate of formaThe low activation energy E, = 9.4 kcal./mole for tion of individual products against the initial alcohol o-methylbenzyl alcohol, compared with the average value of 26.0 kcal./mole for other compounds, appears to be in good agreement With the much higher SU& (10) L. Ya. Margolie and S. Z. Roginskil, Probl. Kinetiki i Katdiza, AkQd. Nauk SSSR, 9 , 107 (1957). ceDtibility of alcohols toward oxidation. (11) T. Vrbagki and K. W. Mathews, Symposium on Heterogeneous T h iugheeA&ms Treatment. In Table IV a Catalysis, 150th National Meeting of the American Chemical Sosummary of the thermodynamic qUahtitiW mi*, ciety, Atlantic City, N. J., Sept. 1965.
KINETICS OF OXIDATION OF 0-METHYLBENZYL ALCOHOL OVER VANADIA
concentration in the carrier gas were constructed from experimental data reported in ref. 1. The rate isotherms for the formation of o-tolualdehyde and carbon oxides were found to exhibit the same trends as observed for the disappearance of o-methylbenzyl alcohol (Figure 7 of ref. l),whereas the isotherms for both the maleic and phthalic anhydrides were straight lines. However, the rates of formation of 0-toluic acid showed an inverse linear dependence upon the alcohol concentration, suggesting a superimposed mass-transfer effect. This phenomenon, however, appears to be of little significance in the over-all desorption process due to the small concentration of the acid formed under conditions employed in this study. It was therefore concluded that the desorption of the by-products from the catalyst surface proceeds in a fashion similar to that observed in the case of o-tolualdehyde, presumably with a heat of activation comparable to that found for the desorption of o-tolualdehyde. ( b ) Equilibyium Adsorption. The results indicate that the AHl, AHz*, and AS1 values for o-methylbenzyl alcohol are consistently about one-half those reported for phthalic anhydride. The reasons for this difference are not fully understood. Values of the same constant in similar systems should be comparable and relatively small since they characterize the transition from the activated complex on the catalyst surface to the adsorbedmolecule. The values for the free energy change AF,, calculated from the equilibrium constant K at selected temperatures within the extremes, are comparable and quite negative. Although no conclusion could be made from thermodynamic data on the nature of the initial adsorption step, other evidence obtained in the course of the work tends to support the reversible process. First, the products of oxidation such as o-tolualdehyde, o-toluic acid, phthalide, and phthalic anhydride were shown to desorb readily from the catalyst surface under conditions at which o-methylbenzyl alcohol is being adsorbed. Second, these compounds also appear to adsorb on the catalyst surface under the same conditions since they undergo oxidation, when used singly, in a manner similar to that observed in the case of o-methylbenzyl alcohol.1~6,12,l 3 Therefore, the equilibrium, rather than the irreversible adsorption of o-methylbenzyl alcohol, is the more probable process. It appears, however, when data are treated by the Hinshelwood model, that the equilibrium is not established on the catalyst surface during the actual oxidation process. From average experimental data as treated by an expression derived by Hughes and Adams from the absolute rate theory,6 about 1/30,000 of the surface oxygen sites of the vanadium oxide were found to be
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covered with o-methylbenzyl alcohol molecules in the actual oxidation process.14 This suggests that the surface was very sparingly populated with molecules of o-methylbenzyl alcohol during the oxidation process, Since a portion of the alcohol always remained unreacted and therefore probably unabsorbed on the catalyst surface under conditions employed in this study, it was concluded that the rate of adsorption of omethylbenzyl alcohol was relatively small. This result is in good agreement with the steady-state concept as outlined above. To summarize the conclusions, the Hinshelwood treatment appears to be applicable to the oxidation of o-methylbenzyl alcohol over vanadia under conditions of partial conversion when the stepwise oxidation is taking place if it is supposed that the o-methylbenzyl alcohol is also adsorbed on the catalyst surface. The assumption that steady-state conditions are established on the catalyst surface is confirmed. Supporting evidence for this is that the specific rate constant for the adsorption of oxygen agrees within one order of magnitude with those reported in the literature for the oxidation of organic compounds over the same catalyst. It is, however, questionable that the Hinshelwood mechanism is adequate under the conditions of the direct oxidation process at the higher temperatures. The Hughes-Adams treatment suggests that, even with the milder conditions used in this study, the chemisorbed organic species reacts with the lattice oxygens of the catalyst; the oxidation products then desorb, and very rapid reoxidation of the catalyst follows. The adsorption of o-methylbenzyl alcohol is probably a reversible process.
Nomenclature k. kat
k,
ko C, Ca
el
Specific rate constant of oxygen adsorption, 1. g.-1 set.-' Specific rate constant of adsorption of o-methylbenzyl alcohol, 1. g.-1 sec.-l Specific rate constant of reaction of organic reactant, 1. g.-1 set.-' Specific rate constant of reaction of adsorbed oxygen, 1. g.-1 sec.-l Concentration of organic reactant in carrier gas, M Concentration of oxygen in carrier gas, M Fraction of the catalyst surface covered with oxygen
(12) C. E. Morrell and L. K. Beach, U. S. Patent 2,443,832 (1948). (13) W. R. Edwards and R. D. Wesselhoft, U. S. Patent 3,128,284 (1964).
(14) r, = K(kT/h)k'Bze-AHz*/RTehsz*/R. rr = 1.23 X 10-*mole/ sec., aHz* = 31 kcal./mole, Bz = 0.6and593'K. Assumingthat K has a value of unity and ASz is zero, k' was found to be 3 X 1014 molecules of o-methylbenzylo alcohol/g. of catalyst. If the oxygen ions in the lattice are all 3.0 A. apart and the vanadium oxide has an area of about 1 m.z/g., then about 10'9 oxygen sites exist in 1 g. of
g.
*
catalyst.
Volume 60,Number 0 September 1966
3098
et
Fraction of the catalyst occupied by o-methylbenzyl alcohol Specific rate of adsorption of oxygen, moles g.+ sec.-l r. Specific rate of adsorption of organic reactant, moles g.-1 r'. sec.-l Specific rate of reaction of organic reactant, moles g.-1 r, sec.-l Activation energy for the adsorption of oxygen, kcal./mole E. E: Activation energy for the reaction of organic reactant, kcal./mole Kl = N/k.Co, g. sec. mole-' A Experimental rate constant, sec.-1 B Experimental rate constant = equilibrium constant K, atm.'1 Partial prwure of organic reactant, atm. po x Conversion of organic reactant, moles/mole t Contact time, sec. AHl* Heat of activation for the irreversible adsorption of organic reactant, kcal./mole
T. VRBASKI
AH2* Heat of activation for the irreversible desorption of prod ucts, kcal./mole A& Heat of reversible adsorption of organic reactant, kcal./ mole A&* Heat of activation for the reversible desorption of products, kcal./mole aS1 Entropy change for the reversible adsorption of organic reactant, d./(mole deg.) A&* Entropy of activation for the desorption of oxidized producta, cal./(mole deg.) k' Conversion factor for changing the fraction of the surface covered to true surface concentration, moles/g. K Transmiasion coefficient k Bolztmanu's constant, 3.295 X lo-*' cal./deg. Planck's constant, 1.563 X lo-*$ cal. sec. h
Acknowledgment. The author wishes to acknowledge the help of Mr. W. K. Mathews, who did much of the experimental work.
