J. Phys. Chem. 1985, 89, 80-85
80
(B’) and 1545 cm-’ (D). These are assigned to linear P t N O on an oxidized Pt surface and to a bidentate nitrate species Pt0,N=O, respectively. NO2 at low coverages dissociates on Pt to yield a spectrum which is similar to that of NO on Pt/Si02. However, at higher coverages the bands B’ and D assigned above intensify and eventually become major spectral features. When N O is adsorbed on Pt/SiO, which had been pretreated with excess 02,the sharp band B’ is observed first and remains themost intense feature, although accompanied by weaker A, C, and D bands, during subsequent dosages of NO up to saturation. The band B’ is also observed as a weak shoulder at saturation coverage for NO on Pt/Si02 and our results clearly show that
this species is associated with N O and oxygen on Pt. Therefore, we conclude that a small quantity of N O dissociates on Pt/SiO, at saturation coverage at ambient temperatures. The final surface compositon at 325 K is very similar for the following systems using Pt/Si02: (a) PtNO (maximum coverage) O2gas (excess); (b) Pt NO2 gas (excess); (c) PtO (maximum coverage) N O (excess).
+
+
+
Acknowledgment. We are grateful to the National Sciences and Engineering Research Council of Canada and Imperial Oil Ltd. for financial support. Registry NO. NO*, 10102-44-0; NO, 10102-43-9; 02,7782-44-7; Pt, 7440-06-4.
Reduction-OxMation and Catalyttc Properties of 12-Moiybdophosphoric Acid and Its Alkali Salts. The Role of Redox Carriers in the Bulkt Noritaka Mizuno, Tetsuji Watanabe, and Makoto Misono* Department of Synthetic Chemistry, Faculty of Engineering, The University of Tokyo, Hongo, Bunkyo- ku, Tokyo 1 1 3, Japan (Received: September 12, 1984)
The catalytic oxidation of CO and H2over 12-molybdophosphoricacid and its Cs and Na salts has been studied at 350 “C in a closed circulating system. By the quantitative comparison of these rates with the rates of stoichiometric reduction of catalysts by CO and H2 and reoxidation by 02,it was demonstrated that (i) these catalytic oxidations proceeded by a redox mechanism, and (ii) in these reactions the difference in the redox carriers in the bulk was closely reflected, as was found previously for stoichiometric reactions. For example, the rate of catalytic oxidation of CO was proportional to the surface area of catalyst, while the rate of catalytic oxidation of H2 depended very little on the surface area. In the oxidation of CO, the surface polyanions are mainly involved in catalysis, owing to slow diffusion of the lattice oxide ion, Le., the redox carrier. On the other hand, in the latter reaction, the redox carriers are protons and electrons which migrate in the bulk, so that all polyanions in the bulk can take part in the redox cycles. Therefore, although the order of the catalytic activity differed considerably between the two reactions, it decreased monotonously with the alkali content (both Na and Cs) when the rate of CO oxidation was normalized to the surface area and the rate of Hzoxidation to the weight. It was further found that linear relationships exist respectively for CO and H2 between the rate of catalytic oxidation and the rate of stoichiometric reduction. These results were discussed on the basis of a redox mechanism.
Introduction Catalytic oxidation reactions over mixed oxides are often believed to proceed by the repetition of reduction and reoxidation of catalysts. This mechanism is called “redox” or the Mars-Van Krevelen mechanism. In this mechanism, the lattice oxide ion is directly involved in the oxidation reaction, so that the reactivity and mobility of the lattice oxide ion are usually very important. Since the participation of the lattice oxide ion in the catalyst bulk was demonstrated by Keulks et al.’ and Hockey et aL2 for the oxidation of propylene over bismuth molybdate catalyst, quite a few studies have been published about the role of the lattice It was reported that the rate of supply of lattice oxide ion to the surface active site played important roles in catalytic oxidation of propylene over bismuth molybdate4 and multicomponent bismuth molybdate catalysts6 and some role in the oxidation of C O over La,-,Sr,CoO, catalyst^.^ In these cases, the lattice oxide ion was incorporated into the products by catalytic oxidation (oxygen-addition reactions), and the lattice oxide ion is the redox carrier in the catalyst bulk. For heteropoly acids, protons and electrons as well as the lattice oxide ion can be redox carriers in the bulk? We previously found that the redox carrier differed depending on the kind of reactions for stoichiometric (noncatalytic) reduction of heteropoly compounds and their reoxidation.lOJ1 The polyanions not only on the surface but also in the bulk were able to take part in the Catalysis by Heteropoly Compounds. VIII.
0022-3654/85/2089-0080$01.50/0
oxidative dehydrogenation reactions, as the redox carriers were protons and electrons. On the other hand, polyanions only on or near the surface were utilized when the lattice oxide ions (and electrons) are the carriers, e.g., oxygen-addition reactions, due to the slow diffusion of oxide ion. In the present work, catalytic oxidations of CO and H2 over 12-molybdophosphoric acid and its alkali salts were chosen as typical examples, and we investigated how the above differences in the redox carriers found for stoichiometric reactions were reflected in the catalytic reactions and attempted to clarify the mechanism of catalytic oxidation over heteropoly compounds and the role of the redox carriers in the (1) Keulks, G. W. J . Caral. 1970, 19, 232. (2) Wragg, R. D.; Ashmore, P. G.; Hockey, J. A. J . Catal. 1971, 22, 49. (3) Keulks, G. W.; Krenzke, L. D. Proc. Inr. Connr. Catal., 6th. 1976, 1977, 2, 806. (4) Christie, J. R.; Taylor, D.; McCain, C. C. J. Chem. SOC.,Faraday Trans. 1 1976, 72. 334. Pendleton. P.; Taylor. D. Ibid. 1976. 72. 1114. ( 5 ) Sakata, K.; Nakamura, T.; Misono, M.; Yoneda, Y. Chem. Lett. 1979,
273. (6) Ueda, W.; Moro-oka, Y.; Ikawa, T. J . Catal. 1981, 70,409. (7) Grasselli, R. K.; Burrington, J. D. Adu. Caral. 1981, 30, 133. (8) Sleight, A. W. ‘Advanced Materials in Catalysis”; Academic Press: New York, 1977; p 181. (9) Misono, M.; Sakata, K.; Yoneda, Y.; Lee, W. Y. Proc. Inr. Congr.
Catal., 7rh, 1980 1981, 1047. (10) Komaya, T.; Misono, M. Chem. Lett. 1983, 1177. (1 1) Misono, M.; Mizuno, N.; Komaya, T. Proc. In?. Congr. Caral., 8th, 1984 1984, 5,487.
0 1985 American Chemical Society
The Journal of Physical Chemistry, Vol. 89, No. 1, 1985 81
Oxidation Catalysis by Heteropoly Compounds H2-reduction timelmin
5
10
I
u
-
0 50 100 150 CO-reduction or 02-reoxidation timdmin
Figure 1. Reductions of H3PMoI2OU, by H2 (line a) and CO (line b), and its reoxidation by O2(line c) at 350 O C . At the dotted line reduction or reoxidation was interrupted for 3 h. This period is not included in the reduction or reoxidation time given on the abscissa. Broken lines indicate the time course without interruption.
catalyst bulk. This paper is an extention of the work which we preliminarily reported before.I2 Heteropoly acids are noted as selective oxidation catalysts in practical processes like oxidation of methacrolein and isobutyric acid. Fundamental studies of the redox m e c h a n i ~ m ~ and * ~ ~the ,'~ correlation between the oxidizing power and the catalytic activity have also been reported.+I6 It is still difficult to find a clear correlation between the catalytic activity and composition (particularly the effects of the metal species in the ~ a l t s ~ - l " ' and ~ ) to describe the oxidation reaction mechanism on a molecular level. We also expected that the present work would shed some light on these subjects.
Experimental Section Catalysts. 12-Molybdophosphoric acid, H3PMol2OW(PMoI2 or H), was obtained commercially from Kanto Chemical Co., Inc. Alkali salts were prepared by titrating aqueous solutions of PMo12 with alkali carbonates as described previously.'oJ1 The alkali salts, MxH3-,PMo12040(x = 1-3, M = Na, Cs) will be abbreviated as MxPMoI2or M,. The water content of the starting materials was determined by a quartz spring balance for the precise adjustment of stoichiometry. The surface area was measured by N2adsorption (flow method) after the catalysts were dehydrated in circulating O2 with a dry ice-ethanol trap at 350 "C. Reduction and Reoxidation of Catalysts, and Catalytic Oxidation. These reactions were carried out in a closed circulation system (about 250 cm3) equipped with a Baratron pressure gauge. The standard pretreatment was as follows: the catalysts (500 mg) were evacuated at room temperature and treated a t 350 OC for 2 h in circulating O2(50 torr; 1 torr = 133.3 Pa), where the water evolved was condensed with a dry ice-ethanol trap. Then the catalysts were evacuated for 5 min and reduced by H2 or C O at 350 "C, the H 2 0 and C 0 2formed being condensed with a liquid nitrogen trap. After the reduction, the reduced catalysts were evacuated for 5 min and then reoxidized by O2 (with a dry iceethanol trap). The rate of reduction and reoxidation was measured by the pressure decrease. In the catalytic oxidation, after the (12) Mizuno, N.; Misono, M. Chem. Lett. 1984, 669. (13) Mizuno, N.; Katamura, K.; Yoncda, Y.; Misono, M. J . Cutul. 1983, 83, 384. (14) Eguchi, K.; Toyozawa, Y.;Yamazoe, N.; Seiyama, T. J. Coral. 1983, 83, 32. Eguchi, K.; Aso, I.; Yamazoe, N.; Seiyama, T. Chem. Left. 1979, 1345. (15) Akimoto, M.; Tsuchida, Y . ;Sato, K.;Echigoya, E. J. Cutul. 1981, 72, 83. (16) Ai, M. Appl. Cuful. 1982, 4, 245.
Reduction or reoxidation timelmh Figure 2. Reductions of H3PMol2O4by CO (a) and H2 (c), m d reoxidation by O2(b) in the presence of water at 350 OC. At the dotted line reduction or reoxidation was interrupted for 3 h. This period is not included in the reduction or reoxidation time on the abscissa. Broken lines indicate the time course in the absence of water (the same as those in Figure 1).
standard pretreatment, a H2-02or C M 2 mixture was introduced and the pressure decrease was followed, H 2 0 or C 0 2 being trapped. In the above runs, the initial pressures of H2 and C O were both 98 torr and that of O2 was 50 torr. Degree of Reduction in the Catalytic Oxidation. In some cases, after a certain period of catalytic oxidation, the catalyst was evacuated for 5 min and reoxidized by 02.The uptake of O2by reoxidation was' measured by the pressure decrease. It was confirmed by separate experiments" that the reoxidation was reversible under the reaction conditions; the amount of O2uptake after reduction by H2was one-half the amount of the H2uptake. After the catalysts were reoxidized, the color of the catalysts was yellow, the same as that just after the standard pretreatment.
Results Reduction by H2 and CO, and Reoxidation by 02.Figure 1 shows the time course of reduction of PMolZby H2 and CO, and its reoxidation by 02.The ordinate is the degree of reduction of polyanions (DR) averaged over whole bulk expressed by the number of electrons introduced per anion ( n e/anion). Prior to the reoxidation, the anion was prereduced by H2 to DR = 1 elanion. At the dotted lines in Figure 1, the reduction or reoxidation was interrupted by evacuating the system and the sample was kept under vacuum for 3 h. After the interruption, the rates of reduction by H2 and CO, and that of reoxidation by 02,which will be denoted by r(H2), r(CO), and r(02), respectively, were measured again. It was noted that r(02) and r ( C 0 ) became several times greater than that just before the interruption (lines b and c). In these plots, the broken lines are those obtained in experiments which were carried out without the interruption. In contrast to r ( C 0 ) and r ( 0 2 ) ,little change by the interruption was observed in r(H2) (line a). Figure 2 shows the time courses for reduction of PMoI2by C O and its reoxidation by O2 in the presence of water. In these experiments, about 0.8 torr of water was initially present in the gas phase. No trap was used in these runs, while in the other experiments a trap for C02 or H 2 0 was used. The dotted lines indicate the same procedures as in Figure 1, and broken lines are the data in the absence of water already shown in Figure 1. In the presence of water vapor, in contrast to the data in Figure 1 (shown by the broken lines in Figure 2), r ( C 0 ) and r ( 0 2 )as well
82
The Journal of Physical Chemistry, Vol. 89, No. 1 , 1985
Mizuno et al.
TABLE I: Rates of Reduction by CO (r(CO)), Its Reoxidation by O2 ( ~ ( O Z ) )Catalytic , Oxidation (r(cat.)), and the Degree of Reduction (DR) at the Stationary State of 350 OC
catalyst r ( C 0 ) (initial)" r ( 0 2 ) (initial)b
r(cat.)' r = r(C0) = r ( 0 2 ) o $ c DRd obsdc calcd' surface ared
PMO12 8.5 x 10-3 1.8 X lo-' 2.4 x 10-3 2.1 x 10-3 0.038
CSIPMO~~ x 10-3 1.6 X lo-' 1.6 x 10-3 1.1 x 10-3
3.1
0.064 0.072 0.77
0.08 1 1.o
CS~PMO~~ 1.4 x 10-3 8.8 X 6.0 X IO4
5.4 x 10-4 0.065 0.091 0.45
Csz.esPMol2 8.5 x 10-3 2.9 X lo-' 2.9 x 10-3 3.1 x 10-3
CS~PMOI~ 3.1 x 10-3 2.9 X 10-1 1.2 x 10-3 1.8 x 10-3
0.050
0.054 0.037 116
0.043 104
"e anion-' m i d . bDegreeof reoxidation after 10 (e anion-'). Catalysts were prereduced by 0.3 e/anion. CCrossingpoints in Figure 4. de anion-'. CExperimental.fm2 g-I. a
Timelmin
Time/ min
Figure 3. Reduction by CO (a) and reoxidation by O2 (b) at 350 OC.
PMo12,CsIPMo12,Cs2PMoI2,and Cs3PMol2; Reoxidation was carried out after prereduction by about 0.3 e/anion.
U
0 , 0 , 0 , and 0 correspond to
OO
t
100
-
200
J
960
Tiinelmin
Figure 5. Pressure decrease (a) and the degree of reduction of catalyst at 350 OC. The (b) in the catalytic oxidation of CO over H3PMoIZ04,, broken line indicates the reduction by CO as shown in Figure 1.
Degree of reductionle anion-'
Figure 4. Rates of reduction and reoxidation of
(x =
0-3) as a function of the degree of reduction (350 "C). Lines 1, 2, 3, and 4 represent 0.3 e/anion redox cycles of PMo12,CslPMo12,C%PMoI2, and Cs3PMo12,respectively. The dotted line indicates the rate of reoxidation of Cs2PMoI2prereduced by 0.13 e/anion (see text). Solid line; rate of reduction by CO. Broken line; rate of reoxidation by 02. as r(H2) did not change by the interruption. It was further noted that a t the initial stage of the reactions the rates in the presence of water vapor were almost the same as those in its absence, but, as the reactions proceeded, r ( C 0 ) and r(Oz)in the presence of water exceeded those in the absence of water (Figure 1, b and c) . Redox Cycles of C S , H ~ - ~ M O ~Time ~ O courses ~ ~ . for the reduction of ( x = 0-3) by CO to DR = 0.3 elanion and their reoxidation by O2at 350 OC are shown in Figure 3. As confirmed by the IR spectra, the Keggin structure of the anion was mostly kept after reduction by CO as in the case of reduction by H2."J3 In Figure 4, r ( C 0 ) and r(Oz)calculated from the slopes of the curves in Figure 3 are plotted as a function of DR. r ( C 0 ) decreased rapidly at the initial stage, but the rate of decrease slowed at the later stage. With increasing Cs content, x , r(C0) decreased from x = 0 to 2 but increased for (not shown in Figures 3 and 4) and Cs3PMol2. r ( 0 2 )decreased with the decrease in DR and the order of r ( 0 2 )was Cs3PMoI2
Figure 6. Relative rates for catalyticoxidations of H2 (a) and CO (b and (0)and NaxH3-xPMo12040 ( 0 )at 350 OC. c) over CS~H~-~PMO~~O,,, Flags attached indicate different lots or pretreatments.
-
Csz,ssPMo12>> PMo12> CslPMo12> Cs2PMoI2when compared at the same value of DR. r(C0) and r ( 0 2 )data at the initial stage are collected in Table I. Catalytic Oxidation of CO and H2. The results of catalytic oxidation of CO (CO + COz) over H3PMo120, at 350 OC are shown in Figure 5. Figure 5a shows the pressure decrease with time and Figure 5b shows the DR values of PMoI2during the catalytic oxidation. After the catalyst was reduced rapidly at the initial stage (note that the reduction was much slower than the reduction by CO alone; the latter is shown by the broken line in Figure 5b), DR remained almost constant (0.035 e/anion). The oxidation reaction proceeded at a nearly constant rate for a prolonged period after the initial rapid reduction of catalyst. So we regard those state as the "stationary state" in the present work. The pressure change was not large, so that it little affected the rate at the stationary state. The rate was reproducible by repeated runs, indicating no deterioration of catalyst, if the initial CO or H2 consumption by the catalyst reduction in the first run was subtracted.
-
The Journal of Physical Chemistry, Vol. 89, No. 1, 1985 83
Oxidation Catalysis by Heteropoly Compounds x ~ ~ 3
1 I
I
I
2
I
I
I
I
I I
I
I
I
I
Z
Surface area/m*g-1
Figure 7. Change of the rates in the catalytic oxidations of CO and H2 with surface area. Catalytic oxidations of CO and H 2 were carried out at 350 "C. 0, A, 0,and 0 correspond to oxidation of CO over PMo12, NalPMolz, Na2PMolz,and Na3PMolz, respectively. corresponds to oxidation of H 2 over Na2PMo12.
0.03 0.04 0.05 006 0.07 ~egreeof reduction/e.anion-l Figure 9. Rates of reduction by CO and reoxidation by O2over Na2HPMol2Omas a function of degree of reduction (350 "C). Lines 1, 2, and 3 represent 0.1 1 e/anion redox cycles of NazPMo12,of which the surface areas are 2.9,2.2, and 1.0 m2/g, respectively. Solid line; rate of reduction. Broken line; rate of reoxidation.
7: x,63
I
TC
3
2
1
I
1
I
'
I
I
*-
0
c
'
:
a
5: o o
5
IO x ~ ~ 3
c
0.1 0.2 0.3 0.4 0.5 Degree of reductionlea-tion-' Figure 8. Rates of reduction by H z and reoxidation by O2over NazHPMoI20@as a function of degree of reduction (350 "C). Lines 1,2, and 3 represent 0.5 e/anion redox cycles of Na2PMolz,whose surface areas are 2.9, 2.2, and 1.0 m2/g, respectively. Solid line; rate of reduction. Broken line: rate of reoxidation.
0
r(cat.) data, the rate of catalytic oxidation for C O at the stationary state for Cs salts, are given in Table I. Similar results were obtained for catalytic oxidation of Hz to H 2 0 and C O to COz over MxH3-xPMo12040 ( M = Na, Cs) a t 350 OC and the relative rates are plotted against x in Figure 6, where the rates over PMo12are taken as unity. r(cat.) for H2 oxidation normalized to catalyst weight decreased monotonously with the alkali content except for an increase with Csz,85PMolz(Figure 6a). For the catalytic oxidation of CO, there is no simple or monotonous relationship between r ( C 0 ) and the amount of alkali (Figure 6b). Moreover, the rate over NalPMolz and Na2PMo12varied from one catalyst lot to another, while those for the catalytic oxidation of H2 varied little between catalyst lots. In Figure 7, r(cat.) for CO and Hz are plotted against the surface area of each catalyst which was measured after the same pretreatment. The surface area of PMo12changed less than 10% during the catalytic oxidation of H,. This figure shows that r(cat.) for CO was nearly proportional to the surface area and r(cat.) for H2 was almost independent of the surface area. Therefore, the scatters of the data observed for r(cat.) for CO over Na,PMo12
Initial rates of reduction by CO/e.aniorilmiril
Figure 10. Correlation between the initial rates of reduction by CO and rates of catalytic oxidation of CO (350 "C). Naz-l, Na2-2, and Na2-3 are NaZHPMol2OM of different lots, of which the surface areas are 2.9, 2.2, and 1.0 m2/g, respectively. Csz-l and Csz-2 are CsZHPMol2Oaof different lots, of which the surface areas are 0.45 and 0.31 m2/g, respectively.
and Na2PMolz(Figure 6b) were likely due to the difference in the surface area. In Figure 6c, r(cat.) for CO normalized to the surface area is plotted against x . In this case, the rate decreased monotonously. These results indicate that r(cat.) for each catalyst composition was determined by the catalyst bulk for Hz oxidation and by the catalyst surface for C O oxidation. In Figures 8 and 9, r(H2),r(CO), and r(0,) for Na2PMolzwith different surface areas are plotted by the same procedure used in Figure 4. The relationship between these data and r(cat.) will be discussed in a later section. In Figures 10 and 11, the rates of catalytic oxidation are compared with the rates of stoichiometric reaction for both CO and H2. Good correlations are notable in both cases.
Discussion Processes for Reduction by CO and H2, and Reoxidation by 4. We have previously reported that the reduction processes of 12-molybdophosphates were phenomenologically classified into two groups, depending on the kind of reductant and the type of reaction.I0J1 The first group is oxygen-addition reactions like oxidation of methacrolein (MAL) to methacrylic acid (eq l ) , in
84 The Journal of Physical Chemistry, Vol. 89, No. 1, 1985 R R'H,
+ -- ++ ++
+A0
+A0
RO
A
Mizuno et al.
(1)
R' 2H+ A02(2a) R' H20 A (2b) A 0 = polyanion (0 = oxygen), R, R'H, = reactants which oxygen of the catalyst is directly taken into the products. The second group is reactions in which the catalyst oxide ions are not necessarily removed. An example is dehydrogenation of isobutyric acid (IBA) to methacrylic acid (eq 2). This classification is based on the presumed difference in the mobility of redox carriers in the catalyst bulk relative to the rate of reduction of the surface. In the first group, the redox carriers are oxide ions (and electrons) which migrate very slowly, so that the reduction of polyanions only on or near the surface is possible and the rate for each catalyst composition becomes proportional to the surface area. On the other hand, in the reduction by IBA (eq 2a), redox carriers are protons and electrons. Both of them formed near the surface and migrate into the bulk as discussed previouslyg and by this migration whole anions can be reduced at a significant rate, so that the rate does not strongly depend on the surface area. In eq 2b, oxygen is removed, but the desorption or migration of water formed is probably rapid at this temperature and may not limit the rate. Absorption of reductant molecules into the catalyst bulk, as was observed for molecules like alcohols at low temperat~res,~J' is not likely in the present case, since the temperature is much higher and the classification does not appear to reflect the absorptivity of the reductant. Among the redox processes studied in the present work, the reduction by CO and the reoxidation by 02,in which redox carriers are oxide ions, belong to the former group, and the reduction by H2 to the latter As expected, the rate of reduction by H 2 (r(H2))changed little with surface area (Figure 7) and r(H2) normalized to catalyst weight showed a monotonous change with alkali content (see Figure 6b) as in the reduction by IBA.lo*llSalts with high Cs contents were sometimes exceptional. This may be due to the low mobility of the proton, electron, or water in the bulk of these salts and/or to their very high surface area. In the reduction by CO, anions only near the surface would be readily involved in the reactions, so that the rate was proportional to the r(02) surface area (Figure 7 plus Figure lo), as in exhibited essentially the same behavior as r ( C 0 ) . For example, r ( 0 2 )for Na2PMo12was nearly proportional to the surface area. This is reasonable, since diffusion of oxygen is also necessary for the bulk polyanions to be reoxidized. It is interesting to note that r ( 0 2 )for C ~ , H ~ , P M O was , ~ Oin~the order of the surface area: x =3 2.85 >> 0 > 1 > 2. The contrast observed by the effect of the interruption in Figure 1 confirms the presence of these two groups in the redox processes. According to the above discussion, for reduction by CO the surface anions were in a much more reduced state than those in the bulk. During the period of interruption, the surface anions were reoxidized through the migration of oxide ion, so that r ( C 0 ) and r ( 0 2 ) after the interruption became greater. On the other hand, owing to the rapid migration of protons and electrons, reduction by H2 proceeds significantly also in the bulk. Therefore, the interruption did not much affect r(H2). Migration of protons and electrons must not be so rapid as to reduce the bulk completely and homogeneously. If so, the surface processes would be rate limiting and r(H2)would have become proportional to the surface area. The presence of two redox processes (r(CO),r ( 0 2 ) ,vs. r(H,)) may be further confirmed by the faster decay of r ( C 0 ) than r(H2) with increasing DR (compare Figure 1 in ref 11 and Figure 1 in this paper). It is worthy to note that in the presence of water (see Figure 2) the effect of the interruption was not observed even in the reduction by CO and the reoxidation by 02.This is probably because water molecules (or their dissociated form) which diffuse rapidly carry the lattice oxide ions. Another possibility is that the water absorbed accelerated the rearrangement of the secondary
-
on'
k u-
0
-1 . -1 Initial rates of reduction by H2/e,anion min
Figure 11. Correlation between the initial rates of reduction by H2 and rates of catalytic oxidation of H2(350 "C). Na2-1,Na2-2,Na2-3,and Cs2-l are the same catalysts as those in Figure 10.
structure. The rapid diffusion of water and the accelerated reaarrangement of secondary structure by water were previously indicated in the isotopic exchange of oxygenI3J8 and in the isomerization of butene.lg The presence of water also made easier the reversible redox cycle of anion, as confirmed by the color, as well as by ESR and IR spectra. Mechanism of Catalytic Oxidation. If the catalytic oxidation proceeds by repetitive reduction and reoxidation of catalyst, that is, the =redoxnmechanism, the rate of catalytic oxidation agrees, to the first approximation, with the rate of reduction of the catalyst and also with the rate of its reoxidation. In other words, the rates of reduction and reoxidation become identical at the stationary state. This stationary rate has been estimated for each Cs salt from the crossing points of the reduction and reoxidation curves in Figure 4 and compared in Table I with the rate of catalytic oxidation. Within the limit of the above approximation, the oxidation state of the catalyst at the stationary state can also be estimated from the abscissa of the crossing points. These values are compared with the data obtained experimentally in the same table. General agreements between estimation and experiments may be seen for both the rate of catalytic oxidation and the stationary oxidation state. Similar agreements have been observed for the catalytic oxidation of H 2 over these catalysts." Fair agreement was found also for N a salts. For example, Figures 8 and 9 predict that catalysts are in a much more reduced state for catalytic oxidation of H2 than that of C O . This was confirmed by experiments; the DR of Na2PMo12(2.2 m2 g-I) was 0.23 e/anion for H2 oxidation and 0.04 e/anion for C O oxidatation. Here it must be considered that the abscissa of Figure 4 is the oxidation state averaged over whole bulk. As discussed in the previous section, in the reduction by CO and reoxidation by 02, the oxidation state of the surface was different from that of the bulk, owing to slow diffusion of the oxide ion. On the other hand, the catalyst bulk must be homogeneous at the stationary state of catalytic oxidation. Therefore, in order to compare the redox rate with the catalytic rate, it is desirable to use the redox rate measured when the catalyst is homogeneously reduced. Nevertheless, the comparison in Table I is essentially correct, and the agreement between the estimation and the experiment can be good evidence for a redox mechanism, as discussed below. r ( C 0 ) and r ( 0 2 )at the stationary state of catalytic oxidation were directly obtained for some catalysts as follows. After the catalytic oxidation reached the stationary state, the systems were evacuated and the initial rates for the reduction of the catalyst by CO and the reoxidation by O2 were measured in separate (18) Sakata, K.;
(17) Okuhara, T.; Hashimoto, T.; Misono, M.; Yoneda, Y.; Niiyama, H.; Saito, Y.; Echigoya, E. Chem. Le??. 1983, 573.
Misono, M.; Yoneda, Y. Chem. Left. 1980, 151. Y.; Misono, M. J .
(19) Okuhara, T.; Kasai, A.; Hayakawa, N.; Yoneda, Catal. 1983, 83, 12.
Oxidation Catalysis by Heteropoly Compounds
The Journal of Physical Chemistry, Vol. 89, No. 1. 1985 85
TABLE II: Rates of Catalytic Oxidation of CO, Reduction by CO, and Reoxidation by O2 at DR for the Stationary State of Catalytic Oxidation at 350 O C catalyst PMoI~ CSIPMOI~ CSZPMO~~ 2.4 x 10-3 1.6 x 10-3 6.0 x 10-4 r(cat.) r(C0) 3.1 X 8.3 X lo4 7.7 x 10-4 r(CO)a.c 3.0 x 10-3 1.2 x 10-3 2.6 X 1.8 X 6.8 X lo4 r ( 0 2 ) (true)"Vb 2.9 X lo4 5.6 X lo4 9.2 X r(02)a,c 0.065 0.038 0.064 DR(true)b*d
r(H2) depended little on the surface area and r(Hz) changed only moderately with DR near the crossing points, the rates at the crossing points of r(H2) and r ( 0 2 )curves did not change much with the change in surface area. In other words, r(cat.) for H2 only slightly depends on the surface area as found experimentally. As discussed previously," the decreases in r(H2) and therefore r(cat.) with increasing alkali content may be ascribed to the change in the redox potential of the polyanion. Possible reasons for the exceptional values observed for some Cs salts have already been discussed in the preceeding section. On the other hand, in the redox cycle for catalytic oxidation of CO, both r(C0) and r ( 0 2 )increased with surface area as shown in Figure 9. Consequently, the rate at the crossing point of r ( C 0 ) and r ( 0 2 )curves varied greatly, in parallel with the surface area. This trend is in agreement with the experimentally obtained r(cat.) for C O (Figure 7). Thus, it has been demonstrated that two types of reactions, i.e., bulk type and surface type, are also present in catalytic oxidation at a high temperature over heteropoly compounds, although the mechanism is different from that in acid catalysis at low temperatures." Next we consider the reason why r(cat.) was nearly proportional to the initial r ( C 0 ) or r(H2) values (Figures 10 and 11). Linear correlations between catalytic activity and reducibility of the catalyst are often regarded simply as evidence for a mechanism in which the reduction process is rate determining in the redox mechanism. However, the correlation in Figures 10 and 11 may be understood in a more reasonable way as follows: (1) As mentioned above, r ( C 0 ) and r(H2) do not depend much on DR near the crossing points (e.g., Figure 4), so that rate at the crossing point is not much affected by the change of r ( 0 2 ) . In other words, r(C0) and r(H2) curves mostly determine the rate at the crossing point. (2) The r(Hz) and r ( C 0 ) curves vary with DR more or less similarly as seen in Figures 4, 8, and 9. So the initial r(Hz) and r ( C 0 ) have linear correlations with r(cat.)'s as in Figures 10 and 11. The abscissa values in Figures 10 and 11 are about three times greater than the ordinate values. The reason for this may be evident from above discussion.
" e anion-' m i d . Experimentally obtained for the stationary state of catalytic oxidation. CObtained from Figure 4 for DR(true) given in the last line. "e anion-I.
experiments. r(CO), r ( 0 2 ) and , DR thus obtained will be denoted in this section as r(CO)(true), r(02)(true), and DR(true), respectively. These values as well as r(cat.) are summarized in Table 11. It is noted that r(CO)(true), r(02)(true), and r(cat.) essentially agreed with one another, confirming that the catalytic oxidation of CO over these catalysts proceeds by the redox mechanism. In the same table, r ( C 0 ) and r ( 0 2 )which are read in Figure 4 for DR(true) of each catalyst are also given. r ( C 0 ) and r(CO)(true) agreed, probably because the heterogeneity was small when DR was very low. Poor agreement between r ( 0 2 ) and r(O2)(true) may be understood by the fact that the reoxidation rate depended on the extent of prereduction as shown by the dotted line in Figure 4 (Cs2PMo12,prereduced by 0.13 e/anion). In spite of the poor agreement, the value of r(cat.) calculated from the crossing point in Figure 4 changed only a little. This is the reason why r(cat.) values have been well estimated from r ( C 0 ) and r(Oz) curves. Therefore the rate of catalytic oxidation may be reliably discussed on the basis of the reduction-oxidation rates of catalysts. Two Patterns in the Catalytic Oxidation and Summary of the Effects of Alkali Salts. In the catalytic oxidations over alkali salts of 12-molybdophosphates, two patterns which were found in the previous papersI0J1 and in the present work for stoichiometric reduction processes were reproduced as shown in Figures 6 and 7. For example, r(cat.) for Hz oxidation depended little on the surface area, while r(cat.) for CO oxidation was proportional to the surface area. These differences may be explained approximately as follows, by referring to the redox cycles shown in Figures 8 and 9. In the catalytic oxidation of H2 (Figure 8), r ( 0 2 )increased with surface area (line 1 > 2 > 3),20but since
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(20) r ( 0 ) shown in Figure 8 is mainly for the reoxidation process of PMoIzOW,f- (111) PMol2Om3-(I) (cf. Scheme I in ref 11 or Figure 9 in ref 13)." However, in the actual catalytic oxidation of Hzrthe PM012040(3+n)(11) PMO,~O,~,~(I) process may be important as discussed previously." If so, the reoxidation would become less dependent on the surface area, since protons and electrons are redox carriers in the I1 G I redox cycle. In this case, r(Oz) curves 1, 2, and 3 in Figure 8 would come closer, resulting in closer crossing points.
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Acknowledgment. We acknowledge Dr. T. Okuhara for useful discussions. This work was supported in part by a Grant-in Aid for Scientific Research From the Ministry of Education, Science and Culture, and by the Asahi Glass Foundation for Industrial Technology. Registry No. PMo12,12026-57-2; Cs1PMo12,77849-37-7; Cs2PMo12, 80050-14-2; Cs3PMo12,12026-64-1; NalPMo12,77849-35-5; NazPMolz, 55624-58-3; Na3PMo12,1313-30-0; CO, 630-08-0; H2, 1333-74-0.
