8509
J. Phys. Chem. 1992,96,8509-8516 (13) Schwartz, S.E. Chemistry of Multiphase Atmospheric Systems. In NATO AS1 Series Vol. G6; Jacschke, W., Ed.; Springer-Verlag: Berlin, 1986. (14) Pruppacher, H. R. Microphysics of Clouds and Precipitation; D. Reidel Pub. Co.: Dordrecht, Holland, 1978. (15) Solomon, S.Rev. Geophys. 1988, 26, 131. (16) Molina, M. J.; Tso, T.-L.;Molina, L. T.; Wang, C.-Y. Science 1987, 238, 1253. (17) Tolbert, M. A.; Rossi, M. J.; Malhotra, R.; Golden, D. M. Science 1987, 238, 1258. (18) Altv, T. Proc. R . SOC.London A 1931, 131, 554. i19i Sinanvalla. A. M.: Alofs., D. J.:. Carstens. D. J. J. A m o s . Sci. 1975. 32,'592. (20) Levine, N. E. J. Geophys. Res. 1973, 78, 6266. (21) Wenzel, H. Inr. J. Heat Mass Trans. 1969, 12, 125. (22) Schulze, F.-W.; Cammenga, H. K. Ber. Bunsenges. Phys. Chem. 1980.84. 163. (23) Philos. Map. 1933. IS. 82. ,--,.Altv. . .. ,, T.. . . .. . (24) Alty, T.; Nicole, F. H.Can. J . Res. 1931, 4, 547. (25) Narusawa, M.; Springer,G. S.J. Colloid InterfaceSci. 1975,50,392. (26) Kiryukhin, B. V.: Pliude, N. 0. In Studies ojclouds, Precipitation and Thunderstorm Electricity; Vulfson, N. I., Levin, L. M., Eds.; Amer. Meteor. Soc.: Boston, 1965. (27) Yamamoto, G.; Miura, A. J. Meteor. Soc. Jap. 1949, 27, 257. (28) Pruger, W. Z . Phys. 1940, 115, 202. (29) Delaney, L. J.; Houston, R. W.; Eagleton, L. C. Chem. Eng. Sci. 1964.19, 105. (30) Hammecke, V. K.; Kappler, E. Z . Geophys. 1953, 19, 181. (3 1) Hirth, J. P.; Pound, G. M. Condensation and Euaporation; Progress in Materials Science, Vol. XI; Pergamon: Oxford, 1963. 132) Pound. G. M. J . Phvs. Chem. Ref. Data 1972. 1. 135. (33j Tanner, D. W.; Pop& D.; Potter,k J.; West, D. jnf. J . Heat Mass Trans. 1968, 11, 181. (34) Tamir, A.; Hasson, D. Chem. Eng. J . 1971, 2, 200. (35) Mills. A. F.: Seban. R. A. Int. J. Heat Mass Trans. 1967,10, 1815. (36) Bryson, C. E., 111; Cazcarra, V.; Levenson, L. L. J . Yac. Sci. Technol.
-
~
--1
~
1974, 11, 411.
(37) Fujikawa, S.;Akamatsu, T.; Yahara, J.; Fujioka, H. Appl. Sci. Res. 1982. 38, 363.
(38) Tolbert. M. A.: Middlebrook. A. M.J . Geoohvs. . _ Res. 1990.. 95.. 22423.(39) Hollcnberg, J. L.; Dows, D. A. J . Chem. Phys. 1961, 34, 1061. (40) Groner, P.; Stolkin, I.; Gunthard, H. H. J . Phys. E 1973, 6, 122. (41) Sugawara, K.; Yoshimi, T.; Okuyama, H.; Shirasu, T. J. Electrochem. Soc. 1974, 121, 1233. (42) Rosetti, R.; Brus, L. E. J. Chem. Phys. 1980, 73, 572. (43) Olson, G. L.; Kokorowski, S.A,; McFarlane, R. A.; Has, L. D. Appl. Phys. Left. 1980, 37, 1019.
(44)Murahm, K.; Tohmiya, Y.; Takita, K.; Masuda, K. Appl. Phys. Lett. 1984,45, 659.
(45) Haynes, D. R.; Helwig, K. R.; Tro, N. J.; George, S.M. J . Chem. Phys. 1990, 93, 2836. (46) Tro, N. J.; Arthur, D. A.; George, S. M. J . Chem. Phys. 1989, 90, 3389. (47) Tro, N. J.; Nishimura, A. M.; George, S.M. J . Phys. Chcm. 1989, 93, 3276. (48) Tro, N. J.; George, S.M. Surf. Sci. 1988, 199, L246. (49) Poppa, H.; Moorhead, D.; Heinemann, K. Thin Solid Films 1985, 128. 251. (50) Tro, N. J.; Haynes, D. R.; Nishimura, A. M.; George, S.M. J . Chem. Phys. 1989, 91, 5778. (51) Winkler, A.; Rendulic, K. D.; Wendl, K. Appl. Surf.Sci. 1983, 14, 209. (52) Eisenberg, D.; Kauzmann, W. The Structure and Properties of Wafer;Oxford University Press: New York and Oxford, 1969. (53) Hare, D. E.; Sorensen, C. M. J . Chem. Phys. 1986, 84, 5085. (54) Hare, D. E.; Sorensen. C. M. J . Chem. Phys. 1987, 87, 4840. (55) Gupta, P.; Coon, P. A,; Koehler, B.G.; George, S . M. J. Chem. Phys. 1990,93, 2827. (56) Gupta, P.; Mak, C. H.; Coon, P. A.; George, S.M. Phys. R w . E 1989, 40, 7739. (57) Weinberg, W. H. In Kinetics of Interface Reactions; Grunze, M., Kreuzer. H. J., Eds.; Springer-Verlag: New York, 1987; p 94. (58) King, D. A.; Wells, M. G. Proc. R . SOC.London 1974, Ser. A 339, 245. (59) Cassuto, A,; King, D. A. Sutf Sci. 1981, 102, 388. (60) Faraday, M. Philos. Mag. 1859,17, 162. (61) Hobb, P. V. Ice Physics; Oxford University Press: Oxford, 1969; p 397. (62) Steinbruchel, Ch.; Schmidt, L. D. J . Phys. Chem. Solids 1973,34, 1379. (63) Head-Gordon, M.; Tully, J. C.; Rettner, C. T.; Mullins, C. B.; Auerbach, D. J. J . Chem. Phys. 1991, 94, 1516. (64) Thiel, P. A.; Madey, T. E. Surf.Sci. Rep. 1987, 7, 211. (65) Madey, T. E.; Yatcs, J. T., Jr. Surf. Sci. 1978, 76, 397. (66) Stulen, R. H.; Thiel, P. A. Surf.Sci. 1985, 157, 99. (67) Ehrlich, G. J . Chem. Phys. 1962, 36, 1499. (68) Fletcher, N. H. The Chemical Physics ofIce; Cambridge University Press: 1970. (69) Goodman, J.; Toon, 0. B.; Pueschel, R. F.; Snetsinger, K. G. J . Geophys. Res. 1989, 94, 16449. (70) Toon, 0. B.; Turco, R. P.; Jordan, J.; Goodman, J.; Ferry, G. J . Geophys. Res. 1989,94, 11359. (71) R a n , J. M.; Kjome, N. T.; Oltmans, S.J. Geophys. Res. Lctf. 1991, 18, 171.
A Comparison of Electrochemical and Gas-Phase Decomposition of Methanol on Platinum Surfaces K. Franaszczuk; E. Herrero,t P. Zelenay,t A. Wieckowski,* Department of Chemistry, University of Illinois, Urbana, Illinois 61801
J. Wang, and R. I. Masel* Department of Chemical Engineering, University of Illinois, Urbana, Illinois 61801 (Received: April 10, 1992) By using electrochemical and ultrahigh-vacuum (UHV) techniques, combined with an isotope substitution method, it is found that the mechanism of methanol decomposition on platinum in the electrochemical environment is different than that in the UHV. In the UHV the fmt step in the decomposition proc*ul is the scission of an 0-H bond to yield a methoxy intermediate, whereas in the electrochemical environment, the first step is the Scission of a C-H bond. The difference in the decomposition mechanism is discussed in terms of differences in the local electric field at the surface and in terms of methanol hydrophobic/hydrophilic interactions in solution. The latter affect methanol-water near-surface conformation and predetermine the destiny of the individual methanolic bonds in the catalytic splitting.
Introduction Oxidation of methanol on polycrystalline platinum electrodes has been studied extensively.'-I8 Bagotzky et ala2postulated that the rate-determining step involved the rupture of the C-H bond 'On leave from the Department of Chemistry, Warsaw University, Warsaw, Poland. *Onleave from the Department of Physical Chemistry, University of Alicante, Alicante, Spain. 'Send correspondence to these authors.
-
in a methyl group to yield a CHzOH intermediate:
CH30H
CH20H
+ H+ + e-
(1)
In the gas phase, and at low temperature, methanol was found to adsorb molecularly on group VI11 metals and to decompose with an increase in temperature via the mechanism shown in Figure l.19 First, methanol undergoes 0-H bond scission to yield a methoxy (CH30) intermediate. Then, the methoxy sequentially decomposes to yield carbon monoxide and hydrogen.
0022-3654/92/2096-SS09$03.00/00 1992 American Chemical Society
Franaszczuk et al.
8510 The Journal of Physical Chemistry, Vol. 96, No. 21, 1992
TABLE I: Tafel slopg pad Rates of the Oxidnth (Currents) of Methutd-h and -din the Light and Heavy Sulfuric Acid Electrolyte at Pt(ll1) pad Pt(ll0) Electmks in tbe Potentid Range 0.OSo.uH) Va current at C V current, pA current, pA maximum: pA Tafel slope, mV (E = 0.200V) ( E = 0.400 V)
CH3OH/H20 CDjOH/H20 CH3OD/D20
(111) 16 8.8 8.8
(110) 810 690 800
(111) 106 128 125
(110) 123 110 136
(111) 5.94 1.86 3.96
(1 10)
(111)
152 51.0 88.8
69.8 22.1 nmc
(110) 508 444 528
“The current of 500 pA corresponds to the turnover number of 12.5 molecules per Pt site per second. b C V = cyclic voltammetry. cnm = not measured. TABLE Ik W t i c Isotope Effats in Owidnth of Methmol-L d -d in the U@t pad Hm~yS d f d c Acid Eltctrolyte 8t Pt(ll1) pad Pt(ll0) Electrodes Methanol
Methoxy
Formaldehyde
Formyl
voltammetry”
>2mK
Flgm 1. Mechanism of methanol decomposition on group VI11 and Ib metals under UHV conditions as diecwrPad by Davis and Barteau.19 The results here show that methanol decomposition follows this mechanism on (ZXl)Pt(llO).
It is not, at present, clear to what extent the electrochemical and gas-phase data are com~arab1e.l~Except for Pt(1 1 l), the mechanism in Figure 1 is well established for many transition metals studied in the UHV. The Pt(ll1) surface shows anomalous behavior; methanol adsorbs molecularly and decomposes to CO and H2, and no one has observed any intermediates between methanol and C0.B22 To the contrary, the electrochemicalwork is not conclusive, and the Bagotzky data can formally be explained via the mechanism shown in Figure 1. Our primary plan was to examine and compare the effects of isotopic substitution on the rate of methanol decomposition at solid/liquid and solid/gas platinum interfaces in order to identify the corresponding rate-determining steps. We have used wellordered Pt(110)and Pt( 1 11) electrodes in electrochemistry and the Pt(110) surface in the UHV. In the UHV, such a surface has enabled to isolate the key decomposition intermediates for the UHV-electrochemistry comparison. To our knowledge, this is the fmt report which considers the effects of isotopic substitution of the methyl group on the rate of decomposition of methanol at a platinum/solution interface. However, there are previous data on light methanol oxidation in H2S04and D2S04.23
Experimental Section Electrochemistry. Spherical small platinum single crystals of 0.2 cm in diameter were built and oriented to (1 10) and (1 11) crystallographic planes by the Lipkowski group at Guelph and the Feliu group at Alicante, respectively. The surface orientation was better than l o in each case. The real surface area and the number of surface Pt atomsze26were determined from hydrogen desorption charges.’ All the measurements were carried out in 0.1 M sulfuric acid electrolyte. Before the measurements, the electrodes were prepared by annealing the crystals in hydrogen flame, cooling down to ca. 200 OC, and quenching in ultrapure When the voltammograms of such prepared singlecrystal planes were taken, they were characteristic of clean and ordered surfaces (Figure 2A-D).27-29 (Previous work has shown that when Pt( 1 1 1) and Pt( 110)samples were prepared this way and transferred into the UHV system, they displayed correct LEED patterns for these surfaces.30) Since we observed that the oxidative methanol voltammetry was changing with the number of cycles, the surfaces were flame-treated and quenched before every kinetic measurement at a chosen potential. They were also treated the same way prior to recording every voltammogram presented in this paper. The chemicals used were Millipore water (18 cmMQI, 99.9% heavy water (Sigma Chemical Co.), reagent grade sulfuric acid (Fischer), reagent grade methanol4 (Fischer), and 99% methanol-d (Aldrich). The Pyrex glass, two-compartment cell was equipped with a Luggin capillary. The iR drop was compensated for by the use of a positive feedback option of the PAR 273 potentiostat and the Tektronix 465 oscilloscope (25 MHz). The
chronoamperometry CH30H/CD3OH H20/D20 D20 0.200 V 0.400 V 0.200V 0.400V 1.8 3.2 3.2 1.5 nmb 1.0 3.0 1.1 1.7 1.0
CD3OH Pt(ll1) Pt(ll0)
1.8 1.2
Obtained from the peak current. nm = not measured. experiments were carried out at temperature of 20 f 1 OC. The electrodes were kept in a meniscus position to avoid an exposure of other than desired parts of the crystals to the electrolyte. All potentials are referred to the Ag/AgCl reference electrode in 1 M NaCl solution. UHV M e J m “ & The UHV measurementswere conducted only with the Pt(ll0) surface with which the stabilization of the methoxy radical was po~sible.~’ The apparatus and procedures used were described previo~sly.~~*~’ A single-crystal sample was cut from a Metron single-crystal rod. The sample was polished with diamond pastes and then cleaned via repeated cycles of oxidation, sputtering and annealing until no impurities could be detected by AES. A sharp (2x1) LEED pattern was seen at this stage. Next, the sample was exposed to a measured amount of CH30H, CH30D, CD30H,or CD30Dthrough a capillary array doser. The decomposition process was then examined with TPD, EELS, AES, and LEED. Further details can be found elsewhere. I 9 v 3 ]
Results Vdtammetry. Cyclic voltammetric (CV) profiles for methanol oxidation obtained in “light” (H-substituted) and ‘heavy” (Dsubstituted) supporting electrolytes are shown in Figures 3 and 4 for Pt(1 11) and Pt( 1 lo), respectively. In agreement with earlier data,16932the maximum oxidation current densities are significantly higher for Pt(110)than for Pt(ll1) (Table I, columns 1 and 2). For Pt( 1 1 l), the light methanol oxidation current measured in D 2 0was approximately 2 times lower than that taken for the light methanol in H 2 0 ; see the voltammetric isotope effect in Table 11. (One should notice that an addition of light methanol, CH30H, to D 2 0results in a formation of CH30Ddue to a rapid H/D exchange; see ref 23 and references therein.) When deuterated methanol was added to the heavy water electrolyte, the oxidation currents dropped even further, and the methanol interference was barely noticeable on the voltammogram as referenced to that taken in the clean electrolyte. In general, such voltammetric isotope effects between CH30H and CD30H are smaller with Pt(110)than with Pt(1 11)and are practically nonexisting between H 2 0 and D 2 0 (Figures 3 and 4, Table 11). There are also considerable differences in voltammetric morphology between individual CV profiles (Figures 3 and 4). On Pt( 1 1 l), noticeable events are the development of an asymmetric surface redox system around 0.2 V, the formation of the main oxidation maxima at about 0.4 V, and the appearance of a shallow maximum on the pitivagoing sweep between 0.5 and 0.6 V. The observations at 0.2 and above 0.5 V are associated with the blank current behavior (the current measured without methanol); the first represents remnants of the anomalous wave (Figure 2A), and the second reflects the reduction in a small blank current above
The Journal of Physical Chemistry, Vol. 96, No. 21, 1992 8511
Decomposition of Methanol on Pt Surfaces I
I
I
I
I
II - 1
A
I
I
I
0.4
0.6
0.8
-0.2 0.0 I
I
0.2
0.4
0.6
I
I
I
0.8
1 I
I
0.2
-0.2 0.0
l l
E/V (Ag/AgCI)
E N (Ag/AgCI) F i e 3. Cyclic current-potential curves of the Pt(ll1) electrode in (A) 0.2 M CH,OH in 0.1 M HzSO4, (B) 0.2 M CD,OH in 0.1 M HZSO4, and (C) 0.2 M CH,OH in 0.1 M DZSO4. Sweep rate 50 m V d .
I
I
-0.2 0.0
I
I
I
0.2
0.4
0.6
I
0.8
EN (Ag/AgCI)
EN (Ag/AgCI)
Figure 2. Cyclic current-potential profiles of platinum single-crystal electrodes: the Pt(ll1) electrode in 0.1 M H2SO4 (A) and in 0.1 M D2S04(B); the Pt(ll0) electrode in 0.1 M HISO4 (C) and in 0.1 M 0.1 M D2S04(D). Sweep rate 50 m V d .
Figure 4. Cyclic current-potential curves of the Pt( 110) electrode in (A) 0.2 M CH30H in 0.1 M H2SO4, (B) 0.2 M CDIOH in 0.1 M H2SO4 electrolyte, and (C) 0.2 M CH,OH in 0.1 M D2S04. Sweep rate 50
0.5 V (Figure 2B). However, the key issue in the voltammetric considerations of this study is an interpretation of the main pair of the maxima at 0.4 V. Clearly, the drop in the methanol oxidation current during the positive-going scan coincides with the development of a small oxidative wave at 0.4 V in the clean supporting electrolyte (Figure 2A,B). Since the wave position cannot be correlated with the threshold of bisulfate ad~orption,~~ and its charge is too small to originate from the surface oxidation process, this wave most probably results from a partial discharge of water (see Discussion). On the other hand, the increase in the methanol oxidation current on the negative-going scan (Figure 3) coincides with the development of a negative current counterwave at 0.4 V (Figure 2A,B). Although not truly symmetric, the current-potential maxima on the negative and positivegoing scans are not significantly separated from each other. Consequently, the methanol oxidation peaks observed on the negativeand positive-going scans practically overlap (Figure 3). On the Pt( 1 10) electrode, the voltammetric current increases linearly with potential up to approximately 0.5 V and drops to a very small current level at about 0.6 V (Figure 4). The drop in the oxidation current coincides with the platinum oxide formation as documented in Figure 2C,D. On the reverse run, the current commencement requires a well-developed oxide film reduction (Figure 2C), and the latter process exhibits a pronounced
irreversibility with respect to the oxidation. It is therefore clear that the methanol oxidation current is lower on the negative than on the positive run because of the large irreversibility involved in the surface oxidation and oxide reduction process. CbronoamperometricRate Data Acquisitions. The catalytic rates of methanol decomposition and successive oxidation were measured from current-time plots (chronoamperometry). The sequences of potential s t e p were programmed using the ASYST computer software on IBM PC/AT interfaced to a PAR 273 potentiostat. The sampling rate was 10 kHz. The data were collected for 0.2 M CH30H (0.2 M CD30H) in 0.1 M H2S04 and for 0.2 M CH30H (0.2 M CD30H) in 0.1 M D$04 (in heavy water). Before the rate measurements, activating potential steps were applied, as shown in Figure 5. One activating step was uscd for the Pt(111) and three for pt(1 lo), and the latter program is shown in Figure 5. The positive potential limit of the step, 0.800 V, which coincided with the minimum of the current potential curves in Figures 3 and 4, was low enough to avoid crystallographic rearrangement of the surface atoms. The negative limit of the steps was -0.250 V, and the waiting time at each potential extreme was 1 s. Immediately before recording a current transient, platinum was reductively cleaned at -0.190 V (Figure 5 ) . The latter value was chosen to avoid interference with methanol oxidation process of a small amount of hydrogen that could be
mV&.
8512 The Journal of Physical Chemistry, Vol. 96, No. 21, 1992
Franaszczuk et al.
l
start of
1
8ook 60C -
P
.3
-0.2
“ I
400 -
-0.6
-1.o
0
I
I
I
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I
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I
1
2
3
4
5
6
7
200 -
8
Timelseconds Figure 5. A potential step program used in the chronoamperometric
experiments. 200°
20
40
60
ao
io0
A
160
“0
20
60
40
80
roo
ffms Figure 7. Current-time plot for oxidation of methanol on the Pt( 110) electrode in 0.2 M solution of CHIOH in 0.1 M H2S04at E = 0.400V. Shown is the background-correctedcurrent and the fitting least-squares
line; see text. 7. Parts A and B of Figure 6 show the total and the background currents, respectively, taken at 0.200 V, and Figures 6C and 7 give the difference currents approximated by the least-squares curves (the solid lines).34 In view of a practically unlimited diffusion delivery of methanol to the surface for Cmethanol = 0.2 and of a relatively small current measured (Figures 6 and 7), the drop in the electric current is exclusively accounted for by some form of surface poisoning. According to the previous data, the poisoning is predominantly due to chemisorbed CO f~rmation:~’ M35936
+
i 40
CH30H
I
I
0
20
40
60
80
100
t/ms Figure 6. Current-time plots for Pt(ll0) in 0.1 M H2S0, at E = 0.200 V: (A) the oxidation current with 0.2 M CHJOH in solution, (B) the background current (without methanol in solution), and (C) the difference current (circles) and the fitting least-squarescurve (solid line); see text.
generated at -0.250 V. We found that, if a period of 10 ms was chosen for the Pt(ll0) reduction at -0.190 V, no noticeable methanol chemisorption occurred, as concluded from the absence of hydrogen suppression in the potential range -0.190 to 0.050 V. For Pt( 11l), this reduction time may be as long as several minutes due to a very slow methanol chemisorption at this crystallographic plane of platinum,32especially at the negative potentials. In practice, a 1-s waiting time was used after which the hydrogen desorption process was unaffected by the methanol presence in solution. These observations show that the platinum oxide was completely reduced using the procedure described above, and a contamination-free platinum surface was obtained. That is, the initial rates of methanol oxidation reported below were measured on the electrode surfaces virtually free of both surface oxides and methanol chemisorbate. The blank current transients, due to adsorbed hydrogen oxidation and double-layer charging, were systematically recorded in the clean supporting electrolyte under the same conditions as those used for the methanol oxidation. The methanol oxidation kinetic profiles were obtained after correcting for the background current. Any uncertainty involved in this procedure was negligible since the background current-time plots decayed rapidly and could affect only the shortest time in which the methanol oxidation kinetics was measured (see below). Kimtic Anlysia Figure 6 shows a representative current-time response for oxidation of CH30H in the light acid on Pt( 110) at 0.200 V. The corresponding data for 0.400 V are given in Figure
+
CH30H CO 4H+ 4e(2) The chemisorbed CO formation is postulated in so-called “dual path mechanism” of the electrochemical methanol oxidation,’ in which the main branch is the carbon dioxide formation:
+ HzO* COz + 6H++ 6 2
(3) If so, the total current measured (in the absence of the mass transport limitations), it,,,,, is itotal
= ibackg + fox + iads
(4)
where ibW is the background current, i, is the oxidation current (reaction 3). and i d is the adsorption current (reaction 2). After correcting for the background, the corresponding rate equation at a constant methanol bulk concentration is i = io, + iads= it=o(l- e) + Q, dtl/dt (5) where i,* is the initial current, 0 is the fraction of electrode awered by surface CO, and Q- is the electric charge required to saturate surface with CO. Qmaxcan be expressed as follows: = 4eN~t(m)-’emax (6) where 4 is the number of electrons produced in reaction 2, e is the elementary charge, NR is the total number of surface Pt sites, Omax is the maximum coverage of the surface CO, and m is the number of surface sites occupied by one CO molecule. Further, dO/dt, the rate of the surface CO formation is Qmax
dO/dt = kad(emax - 0)’
(7) where kad is the rate constant for the site blocking process. The surface reaction order was found to be 2 on the basis of the least-squares curvefitting results (seebelow) showing that, of four likely reaction orders (from 1 to 4, assuming m to be integral), the reaction order of 2 consistently provided the best fit at all potentials. Using eq 6 and both the differential and integrated forms of eq 7, the final expression for the methanol-related current, derived from eq 5 , can be written as follows:
The Journal of Physical Chemistry, Vol. 96, No. 21, 1992 8513
Decomposition of Methanol on Pt Surfaces I
V.“
I
I
I
0.4-
li
0, 0.3i al 5
Mcthanol(32,Oamu)
3 o.20.1 -
0.0‘ 0
/I
1
I
I
I
2
3
4
Wt=dW Figwe 8. Electrode potential vs log (i,+) (see cq 4 and text) for methanol oxidation on the Pt( 1 10) electrode in 0.2 M CHoOH in 0.1 M H2S04 (the Tafel plot).
Four quantities shown in eq 8, i,,@ kad,OM, and m, were treated as adjustable parameters and determined from the experimental i-t data by using the least-squares method (Simplex algorithm,j4 typically 300 data points per fitting procedure). The obtained i,* values are given in Table I. Significantly,the same calculation shows that the adsorption current contribution to the total current is negligible; that is, the second term in eq 8 is much smaller than the first one. More details on the modeling of the methanol oxidation kinetics will be given elsewhere. The initial current increases exponentially with potential to yield the Tafel sl0pe3~of 120 mV.(decade)-l (Figure 8; Table I, columns 3 and 4). Above 0.200 V, the rate of the current rise continuously decreases, and the Tafel slope increases. The observed deviations from the Tafel slope from 120 mV.(decade)-’ imply that some changes in the surface properties occur at high potentials (see Discussion). That is, the studied potential ranges below and above 0.2 V are not catalytically equivalent in oxidizing methanol on platinum. Tables I and I1 summarize the electrochemical rate data obtained with the Pt(ll1) and Pt( 110) surfaces and the light and heavy reaction components. Columns 5 and 6 in Table I, and corresponding isotope effects in Table I1 (chronoamperometry), show the data obtained at 0.200 V. For Pt( 1 1 l), it is seen that the rate of methanol oxidation in light water is 3.2 times higher for CH30H than for CDjOH (Table 11). The corresponding value for Pt(ll0) is 3.0. These are typical values of the primary isotope effect in electrochemical system^.^^"^ For Pt( 1 1 1) and Pt( 1 lo), and at 0.200 V, the substitution of H 2 0 by D 2 0 changes the oxidation rates only by a factor of 1.5 and 1.7, respectively (Table 11), which shows that the C-H vs C-D isotope effects are twice as high as those observed between light and heavy water. Based upon the theoretical background contained in ref 39, and using the earlier results,a the D20/H20 effects can be viewed as solvent isotope effects. To further support the proposed discrimination between the solvent and the kinetic isotope effects, the light methanol in the heavy electrolytewas replaced by heavy methanol. As expected, this c a d a further significant drop in the oxidation current. The isotope effects obtained at 0.400V are also shown in Table 11. Notably, the effect measured with Pt( 1 10) at this potential, 1.1, is lower than that observed in the low potential range, 3.0. This observation is consistent with the values of the voltammetric isotope effects taken at, or above, 0.400 V, which are also lower for Pt(ll0) than for Pt(lll), 1 . 1 vs 1.8. It is not excluded that some D/H isotope scrambling takes place on the more catalytic Pt( 110) electrode, additionally activated by the high positive polarization. If so,the catalytic attack on the methanol molecule is preceded by such an isotope scrambling. These precursor methanol states in methanol electrocatalysis will further be investigated. The highest initial current measured from the current-time plots was obtained with the Pt( 110) electrode for light methanol in the
I 100
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I
1
400
500
I EM
7M
600
900
loo0
Temperature (K)
Figure 9. A composite TPD spectrum taken by exposing a clean ( 2 X l P t ( l l O ~surface to 1.16 lannmuirs of methanol and then heating at 14K.s-’.
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’
4doo
Energy Loss, cm -I Figure 10. A series of EELS spectra taken by exposing a clean (2x1)Pt( 110) to 0.4 langmuir of methanol and then sequentially annealing to the temperatures indicated. light electrolyte (Table I) and is around 500 MA. This current was converted to reaction turnovers ( T I ) using the standard formula TI = ipo/(6Ne) (9) where TIis given in molecules.(surface atom)-W; ilrO is the initial current in microamperes, N is the number of surface atoms, and e is the elementary charge. Since n = 6, the highest turnover number obtained is 12.5 molecules.(surface atom)-Id. A comparison of our reaction turnover data with the data obtained by others cannot be made since the direct determination of the methanol oxidation current on the well-defined electrodes by chronoamperometry has not been reported yet. UHV Investigations. We have also examined methanol decomposition on (2Xl)Pt( 110) in UHV. Figure 9 shows a composite TPD spectrum taken by exposing a clean (2Xl)Pt(llO) sample to 1.16 langmuirs of methanol and then heating at 14 K-s-I.~O There are methanol peaks at 140 and 215 K,a carbon monoxide peak at 485 K,a 28 amu peak at 145 K,a broad carbon monoxide desorption between about 150 and 300 K,and a hydrogen peak at 290 K. There is also a small hydrogen peak at about 140 K which we associate with desorption of background hydrogen from methanol-covered surface. In a previous paper,” Wang and Masel have detected traces of formaldehyde production at exposures of 2 langmuirs or more. However, the formaldehyde may be an artifact since much more formaldehyde was detected when the vacuum chamber was back-filled with methanol. EELS was used to identify the intermediates during the TPD experiments described above. Figure 10 and Tables 111 and IV summarize the results of this study. When methanol first adsorbs
Franaszczuk et al.
8514 The Journal of Physical Chemistry, Vol. 96, No. 21, 1992
TABLE III: Comparison of the Peak Positioap in the EELS Spectrum of Methanol Adsorbed on 95 K (1x1)- and (2Xl)Pt(llO) !hutplea to the Peak Positions in Previous EELS Spectra of Molecularly Adsorbed Methanol C H 3 0 H on C H 3 0 H on CH,OH CH30H CHpOH assignt (2XI)Pt(llO) Pt(l11)2' liquid4* vitreous solid48 crystals48 760 680 655 730 790 0 - H out-of-plane bend 1050 910 1029 1032 1029 C-O stretch 1135 1114 1124 1142 CH3 rock CHI bend 1460 1430 1455 1452 1455 2960 2930 2822, 2934 2928, 2951 2829, 2955 CH stretch 3330 3320 3235 3284, 3187 3327 OH stretch unassigned (overtone) 2590 CDBODon CD,OD CD3OD CDiOD CD,OD CDiOD CD,OD CDiOD assiant ((2Xl)Pt(llO) ~xijpt(iio) liquid48 vitreous solid4' crystals48 crysials48 0 - D out-of-plane bend 560 483 533 578, 495 C-O stretch 990 979 975 968 780 818 831 862 CD3 rock 1100 1097 1100 1080 CD, bend 2070, 2235 2082, 2225, 2250 2075, 2213 2075, 2212 CD stretch 2410 2474 2384 2361, 2432 OD stretch TABLE I V Comparison of the Peak Positions in the EELS Spectrum Taken by Adsorbing Methanol onto a 95 K (ZXl)Pt(llO) Sample and Then Annealing to 175200 K and the Peak Positions in Previous EELS Spectra of the Methoxy Intermediate on Various Transition-Metal Surfaces C H 3 0 H on CH,O(ad) CH,O(ad) mode 175 K Pt(ll0) on Pt(lll)49 on Pd(110)50 M-O stretch 330 370 280 C-O stretch 1025 1000 1010 CH3 rock 1130 CH3 bend 1480 1430 1460 CH3 stretch 2990 2910 2800, 2975 M-H stretch 710 unassigned 770
on (2Xl)Pt(llO) at 100 K, one observes an EELS spectrum which is very similar to that for solid methanol, as indicated in Table 111. These results show that methanol adsorbs molecularly at 100 K. Upon heating to 200 K, the 0-H stretching mode at 3300 cm-' disappears, while the C-O stretching mode at 1025 cm-l is slightly enhanced. There is also a new mode at 330 cm-' which is the position expected for a platinum-oxygen stretch. A comparison of the 200 K EELS spectra in Figure 10 to those in the previous literature as shown in Table IV indicates that the adsorbed methanol decomposes to produce a methoxy intermediate and presumably adsorbed hydrogen at 200 K. (The hydrogen is not detected in EELS because is such a weak scatterer.) The data are more complex upon further heating to 220 K. TPD shows that some of the methoxy recombines with hydrogen to yield methanol. Concurrently, a series of transients are seen in EELS. Initially, the EELS spectrum looks as though one would expect if formaldehydeand a formyl species form. However, the formaldehyde and formyl species disappear after a 2-min anneal at 220 K. Simultaneously, two other peaks grow into the EELS spectrum corresponding to linear- and bridge-bound CO. As a result, we conclude that methanol adsorbs molecularly on (2X 1)Pt(110) at 100 K. It then decomposes via the mechanism in Figure 1. First, the methanol loses its hydroxyl hydrogen to yield a methoxy intermediate. Then, the methoxy intermediate sequentially loses hydrogens to produce CO. We have also examined the effects of isotopic substitution on our TPD spectra. In these experiments, a clean (2Xl)Pt(llO) sample was dosed with varying amounts of CH30H, CD30H, CH30D, or CD30D, and then the sample was heated at 14 K d and a TPD spectrum was recorded. Figure 11 shows how the area of the CO desorption peak varied with exposure during these experiments. Generally, it was found that the CO peak area varies linearly with exposure up to 0.4 langmuir and then becomes saturated. Interestingly, at low coverages, 60% more CO production is observed with CH30H and CD30H than with CH30D and CD30D. We attribute this difference to a kinetic isotope effect. Evidently, at low coverages, the mechanism, in the irreversible adsorption of methanol to eventually yield CO, is the
CO From Deuterium-Labeled Methanol On (2xl)Pt(llO)
CHPD CD@D
V
0
I 0.1
02
0.3
0.4
, 1
2
3
4
Exposure, L Figure 11. Variation in the area of the CO TPD peak with methanol exposure, for CH30H, CH,OD, CD,OH, and CD30D decomposition on (2Xl)Pt(llO).
removal of a hydrogen from the OH group and formation of the methoxy intermediate (see Discussion).
Discussion Eled"ical Decomposition. The theory generally predicts and experiments confirm3*that if an activation barrier in the rate-detennining step (rds) of an electrode process is approximately symmetric, i.e., the transfer coefficient,a,is about 0.5,5I the Tafel slope of about 120 mV.(decade)-' should be observed and considered as an evidence of a one-electron charge-transfer process. In other cam, in the absence of mass-transfer complications, the theoretically predicted Tafel slopes are always lower than 120 mV-(decade)-'. Therefore, the value of the Tafel slope of 120 mV*(decade)-' found between 0.050 and 0.200 V ascertains that the initial catalytic decomposition of methanol with Pt( 111) and Pt( 110) surfaces is associated with a one-electron-transferprocess. Since we have also identified the primary isotope effect between the C-H and C-D bonds, and not between the 0-H and 0-D bonds (Table 11), the appropriate catalytic mechanism is such that methanol decomposition occurs via the C-H bond scission: CH30H(sol)
CH20H(ad)
+ H+ + e-
(10)
where CH30H(sol) and CH,OH(ad) represent the solute methanol and the adsorbed methanol intermediate, respectively. We do emphasize that the successive, not rate-determining steps, could not be considered in this study due to the lack of data on the intermediate evolution when the reaction proceeds. At potentials higher than 0.200 V, the current rise decelerates and the Tafel slope noticeably increases (Figure 8). The current
Decomposition of Methanol on Pt Surfaces
The Journal of Physical Chemistry, Vol. 96, No. 21, 1992 8515
rise deceleration is also shown by the voltammetry (Figures 3 and 4), since its main feature is the peaked current-potential morphology not related to mass-transfer limitations. A model for the two samples of data, chronoamperometric and voltammetric, may be worked out by emphasizing that the oxidation process is preceded by a substitution reaction:
our results are consistent with the previous literature for methanol decomposition in both gas phase and the electrochemical environment. Still, it is quite interesting that on the same sample methanol decomposition proceeds via a different mechanism in gas phase than in an aqueous electrochemical environment. At present, we do not have a quantitative model for these differences. originally, we considered that water might be a direct CH3OH Pt-H,O Pt-CH3OH H20 (1 1) participant in the reaction.44 For example, water could conceivably react with one of the methanol CH3groups to produce the CH20H or/and intermediate. Note, however, that if this was to occur, one would expect to observe a larger difference between the oxidation rate CH30H Pt-HS04Pt-CH30H HS04- (12) in H 2 0 and in D20 than between light and heavy methanol in H20. This is not what is observed in this study. Further, the rate Reactions 11 and 12 determine the availability of weakly bounded should vary much less with the surface structure than is actually methanol for the oxidation process (reaction 10). If either reaction observed (Tables I and 11). It is therefore clear that the water 11 or 12 is potential dependent, a deviation from the regular is not a direct participant in the rate-determining step of the 120-mV Tafel behavior may result. Below 0.2 V, the potentials reaction. Rather, the main effect of the water is to modify the are not far from the potential of zero charge for platinum, and interaction between the methanol and the electrode through the the interfacial water is relatively loosely bound to the surface.40 solvent effects, as discussed in ref 23. Therefore, the surface water should not significantly affect In fact, this is not completely surprising. Barteau19 notes that methanol interactions with the electrode. Similarly, bisulfate is simple thermodynamics suggests that a methanol molecule would loosely bonded to the surface at such relatively less positive porather break a 94 kcal*mol-' C-H bond than a 104 kcal.mol-' 0-H t e n t i a l ~ .While ~ ~ a compact monolayer of anions may significantly bond. However, the first step in the methanol decomposition on interfere with methanol adsorption, the loosely packed anionic all of the group VI11 and Ib metals in gas phase (or UHV) is structure is not expected to do so. Noticeably, the deceleration always the scission of the 0-H bond, even on those surfaces which of methanol oxidation rise coincides with the end of the curbind oxygen weakly. No one knows why the 0-H bond scission rent-potential profile observed between 0.05 and 0.2 V (Figure dominates on so many transition-metal surfaces. However, one 2), that is, in the potential range where the anion adsorption possibility is that at low coverages the gas-phase methanol binds ~ t a b i l i z e s . ~This ~ * ~ coincidence ' makes the anionic interference with the oxygen end down on most transition-metal surfaces. The a very likely explanation of the deceleration process observed by hydroxyl hydrogen is in close proximity to the surface while the chronoamperometry. The role of anions in methanol oxidation methyl hydrogen is 2-3 A away. It appears that, from the gekinetics at single-crystal electrodes will further be investigated ometry alone, one would expect the scission of the 0-H bond to in this laboratory. be favored over the scission of the 0-H bond. At even higher electrode potentials, the electrode properties of At present, there is no experimental data on the geometry of the interfacial water change significantly, as shown by the voltsurface methanol in aqueous environments before its decomposition a m e t r i c behavior of clean platinum (Figure 2). Namely, water and oxidation begins. However, there are several reasons to suspect discharge clearly begins at 0.35 V (Figure 2). Further develop that, in the presence of water, methanol could orient with its CH3 ment includes water splitting at 0.5 V for pt( 110) and a sustained group pointing toward the surface. Platinum is a fairly electropartial discharge of water at Pt( 111). Both processes suppress positive element:' and in UHV the local electric field near the the voltammetric oxidation of methanol (Figures 3 and 4 and surface is aligned in a way that the methanol would tend to orient Results section). We also notice that the ascending part of the itself with the oxygen end down. However, at the solid/liquid current-potential plot for the methanol oxidation on Pt( 110) has interface, the electric field is reduced because all the measurements an unusual linear character, not anticipated from the theory of voltammetric oxidation of neither surface bound s p e c i e ~nor ~ ~ , ~ ~ here have been done near the potential of zero charge, i.e., at 1-1.5 V below the potential of platinum in UHV.45 In addition, water electroactive solutes.35z6 Therefore, at such high electrode pomolecules arrange themselves on the surface with the negative tentials, we may connect such a voltammetric behavior with a poles down toward the surface (see ref 45 and above). This tends reduced availability of the contact-adsorbed methanol (reaction to reduce the surface electric field. A further reduction of the 11) as a precursor for reaction 10. We also conclude that the drop field is due to the high dielectric constant of water. As a result, in surface concentration of the weakly bonded methanol is due there is less of an electric field in an aqueous environment to pull to the increase in the strength of watersurface interactions with the oxygen end of the methanol toward the surface that there is increasing electrode potential. To avoid these complications in in the UHV. the comparison of the solution and vacuum methanol oxidation Hydrophilic/hydrophobic interactions would also be expected kinetics, only the data obtained at, or below, 0.200 V will further to affect the reactivity of liquid/solid methanol. In solution, the be considered. OH in the methanol is thought to be hydrogen bonded to about UHV Decomposition and the Comparisons. Contrary to what three water m o l e c ~ l e s . ~Several * ~ ' other water molecules cluster we observed in solution, the rates of the UHV decomposition of around the OH group. Those water molecules would tend to CH30H and CD30D are the same, and the primary isotope effect prevent the OH group from getting close to the surface. In is between CH30H and CH30D (Figure 9 ) . In fact, our EELS contrast, the CH3 group in the methanol is not protected in sodata in Figure 10 show that the decomposition goes through a lution. Thus, it would be harder for the methanolic OH group methoxy (CH30) intermediate, which is consistent with the isoto bind to the surface in the presence of water and expose the OH topic substitution data. The corresponding mechanism is depicted bond to the catalytic sites. in Figure 1. These combined solid/gas and solid/liquid results show that the rate-determining step of methanol decomposition, Such results raise an issue if the mechanistic data from UHV leading to subsequent oxidation in the electrochemical cell, is can be extrapolated to the electrochemical environment. The different than methanol decomposition in vacuum. That is, the results here demonstrate that water does not directly participate latter process goes via a CH30 intermediate (Figure l), while in in the oxidation of methanol. However, the presence of water and the electrochemical cell it goes through a CH20H intermediate ions moderates the interaction between methanol and platinum (reaction 8). While such results are not unexpected, there was so that the mechanism of methanol decomposition is quite different not any direct evidence prior to this work that the CHzOH inin the UHV than in the electrochemical environment. It is clear termediate forms at the solid/liquid interface. that we need to learn much more about the influence of water It has previously been assumed that the reaction also proceeds on the dynamics of molecules near surfaces before we are able through a methoxy intermediate on platinum in the UHV. The to reliably extrapolate rate or mechanistic results from the UHV results here confirm that assumption on (ZXl)Pt(llO). Hence, to an electrochemical cell.
+
+
+ +
8516 The Journal of Physical Chemistry, Vol. 96, No. 21, 1992
Conclusions We have compared the mechanism of methanol decomposition in UHV to that in an electrochemical environment. In the latter environment, we found that the reaction rate is a factor of 3 slower with CD30H than with CH30H and, at low potentials, the Tafel slope is appropriate for a one-electron process. We conclude that the first step in the decomposition of methanol is the scission of the C-H bond. In contrast, in the gas phase we have observed a factor of 1.6 variation in rate between CH30H and CH30D which shows that the first step in the decomposition of methanol is the scission of the 0-H bond to yield a methoxy intermediate. We have also observed the methoxy intermediate in EELS. These results demonstrate that the mechanism of a simple decomposition reaction can be quite different in the gas phase than in electrochemical environment. The difference in the decomposition mechanism is due to the difference in the local electric field effect between the two studied interfaces, as well as due to the hydrophilic/hydrophobic interactions of methanol in solution vs the absence of such interactions in the gas phase.
Acknowledgment. This work was supported by the National Science Foundation under Grant DMR 89-20538. E.H. acknowledges support via the F.P.I. grant of the Generalitat Valenciana (Spain), and A.W. thanks C. K. Rhee for writing the computer program for the electrochemical data acquisition. Sample preparation was done using the facilities of the University of Illinois Center for Microanalysis of Materials which is supported as a national facility under National Science Foundation Grant DMR 89-20538.Equipment was provided by NSF Grants CPE 83-51648and CBT 87-04667. References and Notes (1) Parsons. R.: VanderNoot. T. J . Electroanal. Chem. 1988. 257. 9. (2) Bagotzky, V. S.; Vassiliev, Yu. B.; Khazova, 0. K. J. Elecrroanal. Chem. 1977,81, 229. (3) Wieckowski, A.; Sobkowski, J. J . Electroonab Chem. 1975.63, 365. (4) McNicol, B. D. J. Electroanal. Chem. 1981, 118, 71. (5) Kadirgan, F.; Beden, B.; Leger, J. M.; Lamy, C. J. Electroanal. Chem. 1981, 125, 89. (6) Clavilier, J.; Lamy, C.; Leger, J. M. J . Electroanul. Chem. 1981, 125, 249. (7) Pletcher, D.; Solis, V. Electrochim. Acta 1982, 27, 775. (8) Lamy, C.; Leger, J. M.; Clavilier, J.; Parsons, R. J. Electroanal. Chem. 1983, 150, 7 1. (9) Ota, K. I.; Nakagawa, Y.; Takamashi, M. J. Electroanal. Chem. 1984, 179, 179. (10) Willsau, J.; Wolter, 0.;Heitbaum, J. J. Electroanal. Chem. 1985, 185, 163. (11) Nicolas, M. J . Electroanal. Chem. 1985, 185, 365. (12) Belgsir, E. M.; Hauser, H.; Leger, J. M.; Lamy, C. J . Electrounul. Chem. 1987, 225, 281.
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