890
J . Phys. Chem 1990, 94, 890-894
trostatic interactions to the effect of line tension. Experimentally, however, r is nearly constant for a given lipid system at fixed temperature. By use of fluorescence film balance the microstructure of the two-dimensional lipid system is typically observed along a pressure-area isotherm. In the fluid/gel coexistence region, the fraction of molecules in the gel phase increases as the total average area per molecule is decreased. As this happens, the gellike domains generally increase in size, rather than in number.1° The theoretical analysis presented in this work can be applied to the behavior of equilibrium shapes observed experimentally as the area of the domains is varied. Recall that as the domain area is increased, the value of r at each predicted equilibrium shape transition decreases continuously. This implies that, at fixed I’,shape transitions occur as the area is varied, with higher order modes becoming more stable with increasing area. In other words, the variation with domain area of the mode and relative amplitude of the shape with lowest free energy is qualitatively similar to the variation with I?, as shown in Figure 4. For example, at some critical domain area, a regular undulating domain with mode m and relative amplitude e * , can coexist with a domain with mode m I and relative amplitude e*,+,. Over the range of areas where a given mode is the most stable, the relative amplitude increases with increasing area. Note that circular domains become the most stable as the area is decreased. This is in agreement with experimental observations of small domains observed just after nucleation;I0 these are generally circular. Based on the analysis above, the following shape transitions are then predicted: as the domains are further grown, the circular domains will become unstable and undergo a continuous second-order transition to an m = 2 mode domain. At progressively larger areas, the domain undergoes a series of first-order shape transitions between successive modes. The area corresponding to a given shape transition between two successive modes is, of course, a function of r. The larger the r, the smaller the area at each predicted shape transition. While higher order modes become continuously favored as the area is increased, the relative amplitude of the equilibrium domain shape asymptotes to a constant value. Thus as their sizes increase, the domains sprout additional protrusions while their shapes remain qualitatively similar. Furthermore, a domain shape that
+
is pinched off at the center never has the lowest free energy. We also investigated the possibility of mode coupling. In all cases examined, the free energy of mixed-mode domains formed by combining two successive modes was equal to or higher than that for regularly undulating (pure-mode) shapes. The analysis also revealed that the lowest free energy of mixed-mode shapes is in general obtained for a combination of shapes where one amplitude equals zero. We note that we have examined only one, albeit large, class of domain shapes: those that can be described as regularly undulating (and those formed by a linear combination of two such shapes). Hence, we cannot account for, nor predict, domain shapes and associated shape transitions that do not fall within this classification. On the other hand, many lipid systems form domains with shapes that are qualitatively similar to those considered. We also note that the calculations presented herein are strictly valid for systems in equilibrium and in the low-temperature limit.’4 Experimentally, the attainment of thermodynamic equilibrium is quite questionable; in fact, it has been demonstratedlothat the nucleation and growth of domains depend on the compression speed, temperature, and impurity content. Yet, with respect to their shape, it appears domains can reach a metastable state,15 and within this context the analsyis presented herein may be applied. Finally, recall that the shapes and associated shape transitions of the dipolar lipid domains, as predicted by the free energy theory, are governed by a dimensionless number, r, as defined in eq 17. r accounts for the competition between the long-range repulsive electrostatic interactions and the effect of line tension. Unlike line tension, the excess dipole moment normal to the interface is experimentally accessible from surface potential measurements. Thus, the analysis presented here offers a means to indirectly measure the line tension from quantitative measurements of the various shapes and corresponding shape transitions that are predicted. This is currently one focus of experimental work in our laboratory and will be the subject in a forthcoming publication.
Acknowledgment. The authors gratefully acknowledge a NATO Postdoctoral Fellowship awarded to T.K.V. and the Deutsche Forschungsgemeinschaft for financial support of this research.
Catalytic Oxidation and Isotopic Exchange of Hydrogen over 12-Molybdophosphoric Acidt Noritaka Mizuno,* Tetsuji Watanabe,t and Makoto Misono* Department of Synthetic Chemistry, Faculty of Engineering, The University of Tokyo, Hongo, Bunkyo- ku, Tokyo I13, Japan (Received: June 22, 1989) The reactions of H,, such as H2-D2 isotopic equilibration and exchange, reduction of catalyst by H, (noncatalytic oxidation of H? by a catalyst), and catalytic oxidation of H,, have been studied mostly at 573 K over 12-molybdophosphoricacid (PMoI2) and its Na salt. The isotopic equilibration and exchange were apparently very slow, although the uptake of H2 proceeded at a considerable rate. This was in marked contrast to 12-tungstophosphoricacid (PW,,) studied previously, for which the equilibration and exchange were very rapid, while the H2 uptake was negligible. The present results were explained by numerical simulation, which revealed that the accumulation of protons and subsequent formation of water from proton and polyanion took place nearly uniformly in the bulk of PMoi2at a rate comparable with the H, uptake. Due to this, the rates of reduction and the catalytic oxidation of H2 were almost independent of their specific surface area (thus called “bulk-type catalysis, type 11”). Introduction Heteropoly compounds are useful catalysts in both acid and oxidation catalysis. In particular, 12-heteropoly compounds are
practically used in several industrial processes.I In oxidation catalysis, the dynamics of the reduction-oxidation of catalysts is important. We have PreviouslY reported that heteropoly compounds have interesting reduction-oxidation
‘Part 15: Catalysis by Heteropoly Compounds. *Present address: Central Research Laboratory, Nippon Oil Co., Ltd., Yokohama 23 I , Japan.
( I ) Misono, M. Catal. Reu. 1987, 29, 269-321. and references cited therein.
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0 1990 American Chemical Society
Catalytic Oxidation and Isotopic Exchange of H2 over PMo12 properties; that is, in some cases, the reduction and oxidation take place rapidly throughout the solid bulk due to the fast diffusion of protons and electrons. Due to this behavior, the oxidation catalysis of 12-heteropolycompounds are classified into two groups from the standpoint of the dependency of the rate (catalytic and noncatalytic oxidations) on the specific surface One is the surface-type catalysis, in which the rate is proportional to the surface area as in the case of ordinary solid catalysts. The second group is remarkable. The rates of the reactions belonging to this group are almost independent of the surface area but instead depend on the weight (or volume). So we called this “bulk-type catalysis, type 11” (cf. ref 1 for bulk-type catalysis, type I-“pseudoliquid”). Oxidation of H2 and oxidative dehydrogenation of cyclohexene belong to this group. We assumed that in the bulk-type catalysis the reduction-oxidation cycle of catalyst proceeded not only on the surface but also in the bulk of the catalyst, due to the rapid diffusion of proton and electron (redox or charge carriers). We have been studying the mechanism of the bulk-type oxidation catalysis and attempted to rationalize the above assumption. The importance of this mechanism may be understood, if one considers, for example, that the two industrially important processes for the production of methacrylic acid, i.e., oxidation of methacrolein and oxidative dehydrogenation of isobutyric acid, belong to the two different groups, respectively.2 In the previous paper,4 we investigated H2-D2 reactions over 12-tungstophosphoric acid and were able to quantitatively demonstrate on the basis of the two-step reduction mechanism previously proposed (eq 1 and 2,5 where M = W or Mo) that the
-
H2 + PM,20403-
2H+ + PMI2OM5-
(1)
2H+ + PM,20405-
H 2 0 + PM,20393-
(2)
rate-determining step of the reduction by H2 was the formation of water in the solid bulk (eq 2) and, therefore, the rate depended little on the surface area. In the present work, we extended the study of the H2-D, reactions to 12-molybdophosphoric acid, which is a much more important oxidation catalyst, and for this and the salt of this, we found the above-mentioned two groups of oxidation catalysis. We also compared the results with those previously obtained for 12tungstophosphoric acid with respect to the activation of H 2 and the reduction process.
Experimental Section Catalysts. 12-Molybdophosphoric acid ( H , P M 0 , ~ 0 ~ . 2 7 . 4 H ~ 0 , abbreviated as PMo,, hereafter), which was commercially obtained from Kanto Chemicals, K. K., was used. The IR spectrum and the XRD pattern agreed with those of the standard sample used in the previous work^.^,^ Na2HPMoI2O4,-,.9.7H20 was prepared as described previ~usly.~ Catalysts having different specific surface areas were obtained by changing the drying process or the rate of temperature increase of the pretreatment as in the previous work of 12-tungstophosphoric acid (PW12).4 Before and after the reaction, the surface areas of the catalysts were measured by the BET method, using N, adsorption. Reactions. The reactions such as (i) noncatalytic reduction of catalyst by H2, (ii) H2-D2 equilibration in the gas phase, (iii) H-D exchange between the gas phase and catalyst, and (iv) catalytic oxidation of H2 by 0, (H2-02 reaction) were carried out in a closed recirculation system as in the previous work.4 In some cases, reduction by CO and catalytic oxidation of C O were also performed. The initial pressures of H 2 and D, were about 12 Torr (0.15 mmol) and 20 Torr (0.25 mmol) (1 Torr = 133.3 (2) Komaya, T.; Misono, M. Chem. Lett. 1983, 1177. (3) Mizuno, N.; Watanabe, T.; Misono, M. J . Phys. Chem. 1985,89,80. (4) Mizuno, N.; Misono, M. J . Phys. Chem. 1989.93.3334; Chem. Lett.
1987, 1195. ( 5 ) Misono,
M.; Sakata, K.; Yoneda, Y.; Lee, W. Y. Proceedings of the International Congress on Catalysis, 7th. 1980; Kodansha and Elsevier: Tokyo and Amsterdam, 1981; p 1047.
The Journal of Physical Chemistry, Vol. 94, No. 2, 1990 891 I
. 3
4 0.8
Time I min
Figure 1. Reduction of H3PMoI2Oa catalysts having different surface areas by H2 at 573 K. Lines 1-4 represent time courses for the catalysts, of which the specific surface areas were 0.6, 1.1, 1.9, and 2.5 m2@, respectively. Conditions: catalyst, PMo12,1 g; initial pressure of H2, 100
Torr. I
I
.
.
Surface area / m2 g-’
Figure 2. Rates of reduction of PMo12catalysts having different surface areas at 573 K (observed and calculated). 0 and 0 indicate the experimental values for the first 3 min and at 16 min, respectively. Bars indicate the changes in the surface areas caused by reactions. The solid lines were obtained by simulation (see text).
Pa), respectively. For the H2-02 reaction, the initial pressures were H2, 100 Torr, and 02,53 Torr. The initial pressure of H 2 for noncatalytic reduction was usually 100 Torr. The water evolved was condensed by a cold trap. The catalysts (0.5-1 g; 1 g = 0.43 mmol for PMo12)were pretreated in 0, at 573 K for 2 h. The degrees of reduction and the rates of the reoxidation of the catalysts during the catalytic oxidation were determined from the amount and rate of 0, uptakes after the evacuation for 3 min of the system a t the stationary states of the catalytic oxidation of H,, as described bef0re.j
Results Reduction by H2. Time courses for PMo12)sthat have different specific surface areas are shown in Figure 1. The rate of the reduction ( = H 2 uptake) gradually decreased and was almost independent of the surface area in the range from 1 to 3 m2@. The degree of reduction after 16 min was 0.7 electrons/anion for all samples in this range. In Figure 2, the rates of reduction calculated from the slopes of the curves in Figure 1 (denoted hereafter by r(H,)) are plotted against the specific surface area. The bars in the figure express the extent of changes in the surface areas before and after the reaction. The surface area decreased by about 20% after the reaction. It may be clear that r(H2) depends little on the surface area when it is greater than 1 m 2 g 1 . This is in contrast to the reduction by CO, in which the rate was proportional to the surface area.3 r(H2) was almost proportional to the H2 pressure and was 1.1-1.2 times that of D,. The apparent activation energy of H, uptake was 54 kJ.mo1-I (473-623 K). Figure 3 shows the amount of evolved water, which was condensed and determined volumetrically, as a function of the amount of H, uptake. A part of these data have already been reported
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The Journal of Physical Chemistry, Vol. 94, No. 2, 1990
Mizuno et al.
Degree of reduction / e anion-’ 0
Hydrogen uptake / lo4mol g-’ Figure 3. Amount of water evolved during the reduction of PMo12by H2 at 523 and 573 K. The dotted line indicates that expected when the amount of water evolution agreed with the amount of H, uptake. Key: 0 , 573 K; 0,523 K. u -
$
$7
E
al
L
Figure 5. Time courses of catalytic oxidation of Hzon PMo12at 573 K. (a) Lines 1-3 correspond to the time courses of the first, second, and third runs, respectively. (b) The line shows the difference between the first and second runs. Condition: catalyst, PMoIz(surface area, 1.9 m2.g-’), 1 g. I
I
OE?
0
gas phase
-30
v
-20 Y
c
. 58 r
in water 0
10
20
30
40
50
Time I mh
Figure 4. Reactions of H2-D2 over PMo12 at 573 K (observed and calculated): (a) isotopic composition in the gas phase and the pressure decrease; (b) values of (HD)’/(H2) (DJ. the H contents in the gas phase and in water evolved. Marks and broken line are experimental data: 0, H,; 8, HD; 0 , D,; V, (HD)*/(H2)(D2);A, H content in the gas phase; 0,H content in water; broken line, indicates the pressure decrease. Solid lines are those obtained by simulation (see text). Condition: catalyst, PMolz (surface area, 0.6 m2-g-I),I g. previously.6 The amounts of water initially formed were 60% (at 523 K) and 66% (at 573 K ) of the amounts of the H2 uptake, but in the later stage of the reduction, the increments of water formed agreed with the increments of the H2 uptake. H 2 - 4 Isotopic Equilibration and Exchange Reactions. Results obtained at 573 K for PMolz are shown in Figure 4a,b (the marks are experimental data and the lines are from simulation; see Discussion). In marked contrast to the case of PW,2,4the H content in the gas phase increased little, and the H D formation was very small. The H-D reactions proceeded at a comparable rate in the presence and absence of 0,. Catalytic Oxidation of H2. Time courses of the catalytic oxidation of H, (H2 I/,O2 H 2 0 ) over PMo12’sare shown in Figure 5. These are the results obtained by repeated runs over the catalyst having the surface area of 1.9 m2.g-I. Curves 1-3 correspond to the first, second, and third runs, respectively. The difference between curves 1 and 2 is due to the reduction of
+
-
(6) Misono, M.; Mizuno, N.; Komaya, T. Proceedings ofthe International Congress on Cutulysis, 8th, 1984; Verlag Chemie: Weinheim, 1984; Vol. 5, p 487.
0
OO
1
2
3
m a c e area I m2.g-1
Figure 6. Rates of catalytic oxidations of H2 and CO over catalysts having different surface areas: (a) PMoIzat 573 K, (b) Na2HPMo120U, at 623 K; 0 , H2-02,A, CO-0,. catalyst in the first run as discussed previously.6 Curves 2 and 3 almost agreed, indicating the absence of the deactivation of the catalyst. The decrease in the surface area after these runs was less than 20%. The pressure difference between curves 1 and 2 is plotted in Figure 5b. It is noted that the catalyst was reduced at the initial stage, and then the catalytic oxidation proceeded stationarily at a constant degree of reduction of the catalyst. The degree of reduction in the presence of 0, (Figure 5) was much smaller than those observed in the absence of 0, (Figure 1). Figure 6a shows the rates of the catalytic oxidation of H, by 0, (denoted by r(cat.)) over PMo12)shaving different surface areas (solid circles). The averaged rates for the initial 3 min (obtained from the slopes of the curves as in Figure Sa) are plotted in this figure. It may be noted that r(cat.) was a little dependent on the surface area of the catalyst. Similar results were obtained also for Na2HPMoI2Om(Figure 6b). These results are in marked contrast to the results obtained for the catalytic oxidation of C O that are also shown in the figure (open triangles). r(cat.), r(H2), and the rates of O2 uptake (denoted by ~ ( 0 , ) ) as well as the degrees of reduction of the catalyst under the conditions of the catalytic oxidation are plotted in Figure 7 . The latter two values were obtained as described in the previous w ~ r k . ~ . ~ Since they were almost independent of the surface area, r(H2) was shown by one solid line.
Catalytic Oxidation and Isotopic Exchange of H, over PMoI2
The Journal of Physical Chemistry, Vol. 94, No. 2, 1990 893 TABLE I: Values of R , , R l , R,, and N , for PMo12and PW,,at
573 K 1.1, I .8,2.5
PMoI2 PWt2
Rl" 3.2 (at t = 0)
R2'
R,'
NpOb
0.5
52
51
1.8 0.6
6.0 8.6
'Units, IO" moI-g-lmin-l. bunits, IO4 mo1.g-I.
PW12
I 0
I
PMo12
I
0.1
0.2 Degree of reduction / e anion-'
Figure 7. Rates of catalytic oxidation of H2, reduction by H2. and reoxidation by O2for PMo,, catalysts having different specific surface areas plotted against the degree of reduction of each catalyst at the stationary state. The numbers attached to marks indicate the surface areas (after reaction) in m2-g-I. Key: 0 , catalytic oxidation; -, reduction; and o, reoxidation. Reaction temperature, 573 K.
Discussion Analysis of H2-D2 Reaction. In the previous paper? we found in the case of PW,, that the H2-D, equilibration and the exchange between H 2 (D2) in the gas phase and proton in the solid took place rapidly. The data were analyzed based on eq 3, where H2,g (D2.g) 2H+S (D+S) 2H'b (D'b) (3) suffixes g, s, and b indicate gas, surface, and bulk phase, respectively. We demonstrated that the proton migration in the bulk was very rapid and the protons in most parts of the bulk phase were isotopically in equilibrium with the protons on the surface. Then, eq 3 becomes eq 4, where it is assumed that the isotopic H2,g (D2,g)
+
(4)
2H+p
concentration is uniform in the hypothetical portion of the solid bulk denoted by p (the number of protons in the portion p is denoted by N, and that in the whole bulk by Nb. See Figure 8). R1 and R , are the rates of the forward and reverse reactions, respectively. During the H2-D, reaction, water is slowly and uniformly formed in the portion p of the bulk by the reaction between protons and oxygens of polyanions (eq 5). In this case, eq 1 corresponds to eq 3 or 4 and eq 2 to eq 5, where 02-is that of the polyanion and R3 is the rate of eq 5 . 2H+p (D+,)
+ 02-5 H 2 0 (D20)
(5)
Since it was already shown that the diffusion of water and proton is rapid in the bulk of PMoI2,' the results of the present study will be analyzed by using the same model as for PW,,. However, there is a difference between PMo12and PWI2. The amount of H plus D atoms in the gas phase, Ng, decreased at a significant rate in the case of PMo,,, and therefore R , is not constant. As the rate of decrease of Ng was nearly proportional to the H 2 pressure, R , can be expressed by R, = klNg ( k , is the first-order rate constant). R2 and R3 may be assumed to be constant for the initial 30-40 min of the reaction, since the relative changes in the number of protons in the bulk during this period were not significant as compared with the rates of H-D exchange and H, uptake. Then, one can derive the following equations, eq 6-9, for the amount of D2. HD, H2, and water evolved, respectively. These equations correspond to eq 8-10 and 12 in ref 4. Here x, y , z , dx/dt = R p 2 - klx (6) dy/dt = 2R2~t(1 - a ) - kly dz/dt = Rz(1 - a)'
- k,z
d(Nwy)/dt = R3 (7) Sakata, K.; Misono, M.; Yoneda, Y. Chem. Letf. 1980, 151
(7)
(8) (9)
Figure 8. Models of H2-D2 reactions for PMoI2(right) and PW,, (left). Values of R l , R2, and R, at 573 K are given in the unit of IOd mol. min-lag-'. Outer circle indicates the surface of PMoI2or PW,,.
and Nware the numbers of D,, HD, H2, and water molecules, and a and y are the D fractions of the protons in the catalyst (portion p) and water, respectively. On the basis of these equations as well as the equations of material balance (corresponding to eq 13-1 5 in ref 4), the present results shown in Figure 4 were analyzed numerically. Optimization of the parameters, R , ( =klNg), RZ, R3,and N (Np at t = 0), as in the previous work? gave the values listed in d b l e I. In the table, the previous results for PW12are also shown. The relation R , - R2 = 1.5R3,which was observed in this period of the reaction (see Figure 3), was used in the simulation. The value of Npois about half of the whole protons in the catalyst bulk. This result indicates that a significant part of the catalyst bulk was deeply involved in the reactions. If one considers that the catalyst used in the experiment given in Figure 4 was small (less than 1 m2.g-'), the ratio of Npoto the number of whole protons for the other PMo,,'s must be much greater. The calculation as in the previous paper4 using the above parameters reproduced well the observed data, as shown by the solid lines in Figure 4. The small deviation of the calculated lines is possibly due to the presence of a small isotope effect for R,. It is noteworthy that R I and R2 were much greater for PW,, than for PMoI2. This means that the dissociation and recombination of H2 takes place much faster on PW12than on PMo12. Isotopic exchange of hydrogen was also reported by Yoshida et aL8 and by Hodnett and M ~ f f a t . ~ Mechanism of the Reduction of PMo12 by H2. We have previously shown that the reduction proceeds by the two-step mechanism represented by eq 1 and 2.5 As discussed above, eq 1 and 2 correspond to eq 3-5 and r(H2) = R , - R2. In the case of PW,,," R1- R2was almost equal to R3 (rate of water formation) and therefore few protons were accumulated in the bulk. In other words, PWI2was hardly reduced. On the other hand, in the case of PMo,,, R,- R2 was considerably greater than R3 at the early stage (1.5 times at 573 K; Table I and Figure 8), indicating that the protons were accumulated; that is, PMoI2was reduced to a significant extent. In the later stage of the reaction (after 1-2 h; Figure 3), R , - R2 became almost constant and agreed with R3,showing that the number of protons in the bulk became stationary. Therefore, although the equilibrium strongly favors the left-hand side of eq 1 for PW,2,4 the equilibrium for PMo,, favors the (8) Yoshida, S.;Niiyama, H.; Ekhigoya, E. J . Phys. Chem. 1982.86, 3150. (9) Hodnett, B. K.; Moffat, J. B. J . Cutul. 1985, 91, 93.
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J. Phys. Chem. 1990, 94, 894-900
right-hand side. This difference corresponds well to the higher reducibility of PMo,, observed in the temperature-programmed reduction8J0and the redox potential." Another difference is that R, was greater for PMo,,. This indicates that the reactivity between protons and oxygen atoms of the polyanion was higher for PMoI2,although R , and R2 were much smaller. The reason why the reduction of PMo,, was nearly independent of the surface area may now be obvious. In the initial stage of the reduction, there was an approximate relationship, R , - R , = 1 SR,, and in the very late stage R , - R2 = R3 (Table I). Hence, in any case, the rate of H2 uptake (=r(H,) = R, - R 2 ) was controlled mainly by R3, which is the rate of water formation in the bulk. Since this reaction mostly takes place in the bulk, as discussed above, r(H2) depended to a much smaller extent on the specific surface area than the case when the reduction took place only near the surface. This is what we observed i n the present study. as shown in Figures 1 and 2 . Mechanism of Catalytic Oxidation of H , over PMoI2. If the catalytic oxidation proceeds by the repetition of reduction and reoxidation of the catalyst (a redox or Mars-van Krevelen mechanism), r(H2) and r ( 0 2 ) , which were measured for the catalyst at the stationary state, must coincide. This kind of measurement is possible and meaningful for PW,, and PMo,,, because the reduction and reoxidation can proceed uniformly throughout the bulk, and in fact the results obtained previously proved the redox mechanism.) Similar measurements were performed for PMo,,'s having different specific surface areas. The data are collected in Figure 7. While the degree of reduction at the stationary state differed among the catalysts having different and surface areas, r(H2) (shown by one solid line), r(02) (O), r(cat.) ( 0 )for each catalyst agreed well as shown in Figure 7. (10) Katamura, K.; Nakamura, T.; Sakata. K.; Misono, M.; Yoneda, Y Chem. Lett. 1981, 89. ( I 1) Pope. M. T. Heteropoly and Isopoly Oxometallates; Springer: Berlin, 1983.
r(H,) depended little on the surface area, while r ( 0 2 ) was proportional to the surface area, if they are compared for a given oxidation state.3 Therefore, the degree of reduction at the stationary state, at which the two rates agree, became smaller for catalysts having greater surface area. Furthermore, the change of r(H2) was rather small in the range of the degree of reduction examined in the present study (0.094.17 electrons/anion), so that r(cat.) did not much depend on the surface area, as observed. Conclusion (1) H,-D2 reactions (equilibration, exchange, and reduction of the catalyst) over PMo,, were very different in features from those for PW,,. Isotopic equilibration of H2 and D2 and exchange between H2 (D,)in the gas phase and protons in the catalyst were much slower than PW,,, while the reduction of catalyst (H2 uptake) was faster. It was quantitatively demonstrated based on numerical simulation that the equilibration and exchange were apparently very slow because the reduction of catalyst (the H2 uptake and subsequent formation of water) proceeded at a rate comparable with the dissociation of H2 (D2). In the case of PW,,, the dissociation of H2 and its recombination were much more rapid than the H, uptake and subsequent water formation. (2) In spite of the marked difference between PMo12and PW,,, the catalytic oxidation of H2 over PMo,, as well as the reduction of PMo,, by H 2 depended little on the specific surface area of PMo,,, if the surface area was greater than 1 m2.g-I. It was also shown quantitatively that this is because the second step of the reduction, that is, the formation of water from proton and polyanion (eq 5 ) , proceeded nearly uniformly throughout the catalyst bulk. Acknowledgment. This study was supported in part by the Grant-in-Aid for Scientific Research from the Ministry of Education, Science, and Culture (No. 01550610 and 62470070). Registry No. H,PMo,,Oa, 12026-57-2; Na2HMo120,0,55624-58-3; HZ. 1333-74-0; CO, 630-08-0.
Coadsorption of Copper with Anions on Platinum(111): The Role of Surface Redox Chemistry in Determining the Stability of a Metal Monolayer J. H. White* and H. D. Abruiia* Department of Chemistry, Baker Laboratory, Cornell University, Ithaca, New York 14853- 1301 (Received: July 25, 1989)
The voltammetry for the underpotential deposition (UPD) of copper from aqueous sulfuric acid solution on flame-annealed Pt( 1 1 1 ) electrodes pretreated with or in the presence of C1-, Br-, I-, or S2- is reported. We find that the potential for the UPD of copper on Pt( 1 11) decreases in the order C1- > Br- > I- > S2-and is very linearly related to the standard reduction potential for the half-cell reaction CuX + e- Cu + X-. This is consistent with the near-edge features of the observed X-ray absorption spectrum and with a model involving partial charge transfer from the adsorbed copper to the coadsorbate, which is most likely initially present on the platinum surface in a nonanionic (i.e., neutral) state. In addition, the kinetics of cooper UPD for the different cases were also found to vary linearly with the aforementioned standard potential.
-
Introduction The underpotential deposition (UPD) of copper on Pt has been the subject of a number of recent publications dealing with the atomic level structure of this and related systems. It has been demonstrated that copper deposition of Pt(l1 l ) , upon which a well-defined iodine adlattice exists, occurs in such a manner so as to produce superlattices of copper and iodine.' Using surface EXAFS,, we have shown that copper UPD on iodine-coated (1) Stickney, J. L.; Rosasco, S . D.; Hubbard, A. T. J . Electrochem. Soc. 1984, 131, 260.
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Pt( 11 1) produces structures that give rise to very-pronounced fine structure, suggesting the presence of copper clusters on the electrode surface at low (
