Temperature Effects on Methanol Dissociative Chemisorption and

Sep 30, 2000 - Temperature Effects on Methanol Dissociative Chemisorption and Water Activation at Polycrystalline Platinum Electrodes. Dawn Kardash an...
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Temperature Effects on Methanol Dissociative Chemisorption and Water Activation at Polycrystalline Platinum Electrodes Dawn Kardash and Carol Korzeniewski* Department of Chemistry & Biochemistry, Texas Tech University, Lubbock, Texas 79409-1061 Received April 11, 2000. In Final Form: August 11, 2000 A thermostated IR spectroelectrochemical cell was used to probe the effects of temperature (25 - 70 °C) and methanol concentration (1.5 × 10-2 - 1.0 M) on methanol dissociative chemisorption at platinum electrodes in 0.1 M HClO4. The measurements provide molecular level evidence for mechanisms derived from electrochemical studies of methanol oxidation at above ambient temperatures. For fixed methanol concentrations, increasing the temperature lowered the quantity of detectable adsorbed CO and increased the rate of CO2 formation. The results can be understood in terms of the thermal activation of water and CO desorption. At fixed temperature, raising the methanol concentration in solution increased the integrated intensities of the adsorbed CO vibrational bands between 0.2 and 0.7 V (versus reversible hydrogen electrode, RHE). Below 0.9 V, the rate of CO2 formation was faster in 1.0 × 10-1 M than in 1.0 M solutions of methanol, reflecting the concentration dependence of surface poisoning. In 1.0 M methanol solutions, adsorbed CO was observed at hydrogen adsorption and double layer potentials up to the high-temperature limit of the experiments (70 °C).

Introduction The electrochemical oxidation of methanol has been widely studied mainly because of its potential for use in fuel cells.1-5 Considered as an alternative to hydrogen as the anode gas in polymer electrolyte membrane (PEM) fuel cells, methanol can overcome many of the storage and purity limitations of hydrogen in fuel cells designed for transportation and field applications.6 The standard thermodynamic potential of a methanol-O2 cell is nearly identical to that of a H2-O2 cell.2 However, the multistep, six electron process that converts methanol to carbon dioxide (CH3OH + H2O f CO2 + 6H+ + 6e-) suffers severe kinetic limitations. The partial oxidation product carbon monoxide can be particularly deleterious. CO adsorption on the Pt-based catalyst materials typically employed in fuel cells can block sites for methanol adsorption. Compared to the lowest potentials that enable the formation of CO from methanol on Pt, potentials at least 0.1-0.2 V more positive are required to activate water and produce the surface oxides needed to convert CO to CO2.1,3 Major steps in the electrochemical oxidation of methanol are described by the following equations:

CH3OHsol a CH3OHads f COads + 4H+ + 4e- (1) H2O f OHads + H+ + e-

(2)

COads + OHads f CO2 + H+ + e-

(3)

* Corresponding author. E-mail: [email protected]. Phone: (806)742-4181. Fax: (806)742-1289. (1) Hammett, A. In Interfacial Electrochemistry. Theory, Experiment, and Applications; Wieckowski, A., Ed.; Dekker: New York, 1999; p 843. (2) Jarvi, T. D.; Stuve, E. M. In Electrocatalysis; Lipkowski, J. Ross, P. N., Eds.; Wiley-VCH Publishers: New York, 1998; Chapter 3, p 75. (3) Beden, B.; Leger, J.-M.; Lamy, C., In Modern Aspects of Electrochemistry; Bockris, J. O. M., Conway, B. E., White, R. E., Eds.; Plenum: New York, 1992; Vol. 22, p 97. (4) Parsons, R.; Vandernoot, T. J. Electroanal. Chem. 1988, 257, 9. (5) Chrzanowski, W.; Wieckowski, A. In Interfacial Electrochemistry; Wieckowski, A., Ed.; Dekker: New York, 1999; p 937.

More detailed schemes involving different types of surface oxides7 and fragments from the stepwise dehydrogenation of methanol1,2,7-9 have also been presented. The reactions in eqs 1-3 are basic to all the models.1-5 Additional intermediates are discussed in the sections that follow within the context of the experimental results. The overwhelming majority of fundamental voltammetric and in situ spectroscopic investigations of methanol electrochemical oxidation have been performed at ambient temperatures,1-4,10 even though practical fuel cells frequently operate at 60 °C and above.6,11-13 Recently, there have been reported in situ spectral and electrochemical measurements at temperatures other than ambient involving methanol oxidation at fuel cell catalysts,13-18 Pt-Ru alloy electrodes,19,20 and Ru-modified Pt singlecrystal electrodes.5,21 Voltammetry studies of methanol (6) Gottesfeld, S.; Zawodzinski, T. A. In Advances in Electrochemical Science and Engineering; Alkire, R. C., Gerischer, H., Kolb, D. M., Tobias, C. W., Eds.; Wiley-VCH: New York, 1997; Vol. 5, p 195. (7) Schell, M. J. Electroanal. Chem. 1998, 457, 221. (8) Sriramulu, S.; Jarvi, T. D.; Stuve, E. M. J. Electroanal. Chem. 1999, 467, 132. (9) Herreo, E.; Franaszczuk, K.; Wieckowski, A. J. Phys. Chem. 1994, 98, 5074. (10) Korzeniewski, C.; Childers, C. L. J. Phys. Chem. 1998, 102, 489. (11) Gottesfeld, S. In Polymer Electrolyte Fuel Cells; Gottesfeld, S., Ed.; Pergamon: Oxford, 1995; Vol. 40. (12) Gottesfeld, S.; Fuller, T. F.; Halpert, G. In Proton Conducting Membrane Fuel Cells (2nd International Symposium); Gottesfeld, S., Fuller, T. F., Halpert, G., Ed.; Electrochemical Society: Pennington, 1998; Vol. PV 98-27. (13) Ley, K. L.; Liu, R.; Pu, C.; Fan, Z.; Leyarovska, N.; Segre, C.; Smotkin, E. S. J. Electrochem. Soc. 1997, 144, 1543. (14) Long, J. W.; Stroud, R. M.; Swider, K. E.; Rolison, D. R. J. Phys. Chem. B, in press. (15) Bo, A.; Sanicharane, S.; Sompalli, B.; Fan, Q.; Gurau, B.; Liu, R.; Smotkin, E. S. J. Phys. Chem. B (in press). (16) Day, J. B.; Vuissoz, P.-A.; Oldfield, E.; Wieckowski, A.; Ansermet, J.-P. J. Am. Chem. Soc. 1996, 118, 13046. (17) Tong, Y. Y.; Oldfield, E.; Wieckowski, A. Anal. Chem. 1998, 70, 518A. (18) Tong, Y. Y.; Rice, C.; Wieckowski, A.; Oldfield, E. J. Am. Chem. Soc. 2000, 122, 1123. (19) Gasteiger, H. A.; Markovic, N.; Ross, P. N., Jr.; Cairns, E. J. J. Phys. Chem. 1993, 97, 12020. (20) Gasteiger, H. A.; Markovic, N.; Ross, P. N., Jr.; Cairns, E. J. J. Electrochem. Soc. 1994, 141, 1795.

10.1021/la000544e CCC: $19.00 © 2000 American Chemical Society Published on Web 09/30/2000

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oxidation at well-characterized bulk Pt-Ru alloys over the temperature range from 25 to 80 °C highlighted the strong dependence of reaction rate on the competition between methanol adsorption/dissociation (eq 1) and surface oxide formation (eq 2).19,20 In the present study, the balance between the two pathways on polycrystalline Pt at different temperatures between 25 and 70 °C is probed at the molecular level with IR spectroscopy. The mid-IR vibrational bands for CO2 (2343, 2278 cm-1 for 13CO ) and adsorbed CO (2020-2070 cm-1) were monitored 2 for fixed temperatures as a function of potential and methanol concentration. These bands allow the rate of complete methanol oxidation to be compared to the coverage and spatial arrangement of the adsorbed partial oxidation product CO.22-31 The results demonstrate the effect of temperature on the competing pathways involved in methanol electrochemical oxidation. In particular, spectral evidence is provided for the role of water thermal activation20 in lowering the potential for methanol oxidation. Experimental Section IR spectroelectrochemical measurements were performed with a thermostated cell. The design has been described in detail.32 The cell consists of a glass chamber that has an O-ring seal joint onto which a 60° beveled CaF2 window (Spectral Systems, Hopewell Junction, NY) can be secured. The chamber mounts in an insulated aluminum housing, which provides both thermal and mechanical stability. The reference electrode is maintained at ambient temperature in a compartment that is separated from the thermostated glass chamber by a wetted stopcock and connects to the chamber through glass tubing. The working electrode material was polycrystalline platinum (10 mm diam × 1.0 mm thick). The edges of the disk were sealed in the end of a soft glass tube. The glass extended onto the backside of the platinum disk by a few millimeters, leaving the center portion of the metal uncovered. A thermocouple (Omega Engineering, Stamford, CT) was secured to the exposed metal with silver solder. As described previously,32,33 a thermofoil (Mincoproducts, Minneapolis, MN) placed inside the glass tube of the working electrode provided the thermal energy for heating the cell. With the thermocouple and thermofoil in place, the tube was filled with fine silica for thermal mass and closed with a plug of silicone sealant. Cell temperatures were maintained with a PID temperature controller (LFI-3551, Wavelength Electronics, Bozeman, MT). With the cell assembled and filled with 0.1 M HClO4, about 5 and 15 min were required for thermal equilibrium to be attained at 25 and 70 °C, respectively. IR spectra were recorded with a Mattson Instruments R/S-1 FTIR spectrometer system (Madison, WI). The instrumentation used for reflectance measurements and spectral acquisition has been previously described.34 All single-beam IR spectra were computed from the average of 512 interferograms collected at 4 cm-1 resolution. A triangular function was used for apodization. (21) Chrzanowski, W.; Wieckowski, A. Langmuir 1998, 14, 1967. (22) Kunimatsu, K. J. Electroanal. Chem. 1986, 213, 149. (23) Kunimatsu, K.; Kita, H. J. Electroanal. Chem. 1987, 218, 155. (24) Furuya, N.; Motoo, S.; Kunimatsu, K. J. Electroanal. Chem. 1988, 239, 347. (25) Corrigan, D. S.; Weaver, M. J. J. Electroanal. Chem. 1988, 241, 143. (26) Leung, L. W. H.; Weaver, M. J. Langmuir 1990, 6, 323. (27) Iwasita, T.; Nart, F. C. J. Electroanal. Chem. 1991, 317, 291. (28) Iwasita, T.; Nart, F. C.; Lopez, B.; Vielstich, W. Electrochim. Acta 1992, 37, 2361. (29) Christensen, P. A.; Hamnett, A.; Troughton, G. L. J. Electroanal. Chem. 1993, 362, 207. (30) Markovic, N. M.; Gasteiger, H. A.; Ross, P. N., Jr.; Jiang, X.; Villegas, I.; Weaver, M. J. Electrochim. Acta 1994, 40, 91. (31) Shin, J.; Korzeniewski, C. J. Phys. Chem. 1995, 99, 3419. (32) Kardash, D.; Huang, J.; Korzeniewski, C. J. Electroanal. Chem. 1999, 476, 95. (33) Huang, J.; Korzeniewski, C. J. Electroanal. Chem. 1999, 471, 146. (34) Korzeniewski, C.; Huang, J. Anal. Chim. Acta 1999, 397, 53.

Figure 1. Cyclic voltammograms of polycrystalline platinum in 0.1 M HClO4 containing 1.5 × 10-2 M methanol recorded in a cell that was thermostated at the indicated temperatures. The electrode was held at 0.0 V while methanol was mixed into solution and maintained at 0.0 V for a total of 7 min prior to the start of the scan. The scan rate was 50 mV/s. A platinum counter electrode and a KCl-saturated Ag/AgCl reference electrode were used for all electrochemical measurements. Potentials reported in the paper are referenced to the reversible hydrogen electrode (RHE) scale. Cyclic voltammetry studies were performed with a Princeton Applied Research (PAR) model 273 potentiostat/galvanostat. A PAR model 173 potentiostat controlled the electrode potential during IR spectroscopy measurements. Perchloric acid solutions were prepared from distilled water that was further processed with a four-cartridge Nanopure II system. Perchloric acid (redistilled, 99.999% purity) was obtained from Aldrich (Milwaukee, WI). Methanol (Burdick & Jackson, GC grade) was washed over alumina, filtered, distilled, and stored refrigerated. For experiments with methyl-13C alcohol (Aldrich, Milwaukee, WI), 5.0 M stock solutions of 13CH3OH were prepared and stored refrigerated. The platinum working electrode was prepared for each experiment by cycling between 0.0 and 1.6 V in 0.1 M HClO4 until characteristic defined features appeared in the hydrogen adsorption and oxide potential regions. Afterward, the electrode was transferred to the experimental cell that contained freshly prepared 0.1 M HClO4 solution. The electrode was held at 0.0 V and methanol was added in sufficient volume to bring the concentration to the desired level (1.5 × 10-2, 1.0 × 10-1 or 1.0 M). The solution was mixed for 1 min by bubbling with either high purity argon or nitrogen, and then the cell was allowed to stand for an additional 6 min before the start of spectral or electrochemical measurements. Oxygen was removed from the solutions by bubbling for 10 min with high purity argon or nitrogen.

Results Solutions of 1.5 × 10-2 M Methanol. Figure 1 shows cyclic voltammograms recorded with a polycrystalline platinum electrode in N2-purged 0.1 M HClO4 containing 1.5 × 10-2 M (15 mM) methanol at 25 and 70 °C. In each case, the electrode was held at 0.0 V in the perchloric acid solution as methanol was added to the cell and the electrode was left exposed to methanol at 0.0 V for a total of 7 min prior to the start of the sweeps. The voltammograms display waves that are characteristic of methanol oxidation on platinum in aqueous acid solutions.1-3,9,25,35,36 On the forward scan, methanol oxidation is rapid between 0.4 and 0.8 V. The decline in current above 0.8 V reflects the inhibition of methanol oxidation by surface oxides.

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Figure 3. Potential difference IR spectra showing the 13CO2 formed from the oxidation of 1.5 × 10-2 M 13CH3OH in 0.1 M HClO4 at polycrystalline platinum. The electrode was held at 0.6 V, and the cell was thermostated at the indicated temperatures. A single beam spectrum recorded at the start of the experiments at 0.0 V was used as the background.

Figure 2. Linear sweep voltammograms of polycrystalline platinum in 0.1 M HClO4 containing (A) 1.5 × 10-2 M methanol and (B) 1.0 × 10-1 M methanol. The cell was thermostated at the indicated temperatures. Experimental conditions were the same as those described in Figure 1. The potential region between 200 and 600 mV is shown. The sweep rate was 50 mV/s.

The peak current on the forward scan is approximately four times larger at 70 °C than at 25 °C. Similar temperature-dependent enhancements have been observed for methanol oxidation and ascribed to improved reaction kinetics at the higher temperatures.19,20 At 70 °C, the current on the forward scan rises earlier and falls at more positive potentials than at 25 °C. The anodic peak on the reverse scan at 70 °C also reaches a higher current relative to the wave on the forward scan. The wider potential range and greater current for methanol oxidation at 70 °C compared to ambient is consistent with the idea that water dissociative chemisorption (eq 2) is thermally activated over the range of temperatures studied.19,20 The faster rate of reactive oxide production at the higher temperature is expected to reduce the steady-state coverage of adsorbed CO (eq 3) and other methanol intermediates, and thereby expose more surface sites for methanol adsorption and dissociation (eq 1). Figure 2a shows an expanded view of the rising portion of the methanol oxidation wave for platinum in 1.5 × 10-2 M methanol at 25, 50, and 70 °C. As the temperature increases from ambient to 70 °C, the foot of the methanol oxidation wave shifts negative by about 50 mV. These shifts in the methanol oxidation potential are in good agreement with earlier temperature-dependent studies20 and further demonstrate the improvement in reaction kinetics that occurs at above ambient temperatures and potentials associated with water activation. In 1.5 × 10-2 M methanol solutions, methanol adsorption (eq 1) was slow at the initial contact potential of 0.0 V, consistent with earlier measurements in acid electrolyte solutions containing methanol at low millimolar concentrations,19,22,23 and adsorbed CO was barely detectable in IR spectroelectrochemical experiments. However, as the potential was stepped positive, CO2 evolution was readily observed. Infrared spectral bands of CO2 entrapped in the thin layer cavity of the spectroelectrochemical cell during methanol oxidation at 0.6 V are displayed in Figure 3. The spectra show the different quantities of CO2 formed

Figure 4. Potential difference IR spectra of polycrystalline platinum in 0.1 M HClO4 containing 1.0 × 10-1 M methanol in a cell thermostated at the specified temperatures. The electrode was held at 0.0 V while methanol was mixed into solution and maintained at 0.0 V for a total of 7 min prior to spectral acquisition. Spectral acquisition at each potential required approximately 2.7 min. The spectral region for atop-coordinated adsorbed CO is shown. The single beam spectrum recorded at the start of the experiments at 0.0 V was used as the background.

during methanol oxidation for fixed time periods at 25 and 60 °C. The larger quantity of CO2 produced at 60 °C is consistent with the higher methanol oxidation currents in Figures 1 and 2. In these initial experiments, C-13labeled methanol was used to avoid potential spectral interferences near 2343 cm-1 from atmospheric CO2 and CO2 arising from the oxidation of any organic contaminants present in the reagents. Since the CO2 from these sources was well below the CO2 quantity formed during methanol oxidation, C-13-labeled methanol was not needed for subsequent measurements. Solutions of 1.0 × 10-1 M Methanol. When the methanol concentration was increased to 1.0 × 10-1 M, the shift in equilibrium toward adsorbed methanol (eq 1) resulted in higher adsorbed CO coverages. Figure 4 shows the spectral region for the C-O stretching mode of atopcoordinated CO. At 25 °C, adsorbed CO accumulated on the electrode at potentials between 0.2 and 0.5 V over the

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Figure 5. Same as Figure 4 except the CO2 spectral region is shown.

period of the IR spectroelectrochemical measurements (Figure 4a). The weak atop CO band present at 0.2 V in Figure 4a suggests that the CO coverage is low at this potential. It is likely that the methanol-surface interactions are inhibited at 0.2 V by coadsorbed hydrogen.2,19,23,31,36,37 The atop CO band is strongest between 0.3 and 0.6 V. The spectral bands across this potential region indicate that the CO coverages are moderate to high as the integrated intensities of the bands reach 6575% of the values attained in CO-saturated 0.1 M HClO4. The band intensity starts to decrease positive of 0.5 V, which is a signal that the coverage of CO is becoming lower as surface oxide formation begins to compete more favorably with methanol dissociative chemisorption. This idea is supported by the linear sweep voltammograms in Figure 2b and the CO2 spectra in Figure 5a. The spectra show CO2 starts to form by 0.4 V at 25 °C (Figure 5a) and continues as the potential is stepped positive (see also Figure 9a, vide infra). The decrease in adsorbed CO band intensity positive of 0.5 V accompanied by CO2 evolution and a rise in the voltammetric current indicate that there is diminished poisoning by adsorbed CO at these potentials. At 70 °C, the adsorbed CO bands are near the limit detectable with the IR technique (Figure 4b). CO2 is present at 0.3 V (Figure 5b). This is just positive of the hydrogen adsorption region and indicates that water activation becomes enabled at lower potentials as the temperature increases. Below 0.3 V, methanol adsorption may be inhibited by coadsorbed hydrogen, and, above 0.3 V, the thermal activation of surface oxides opens CO2producing pathways and likely helps maintain the CO coverage as low. The dependence of the atop CO spectral band position on electrode potential in Figure 4 contains information about the structure of the adsorbed CO layer. At 25 °C, the band shifts up in energy as potential is stepped positive between 0.2 and 0.5 V and down in energy above 0.5 V. At 70 °C, in contrast, the atop CO band position and band intensity appear insensitive to potential between 0.2 and 0.6 V. The different spectral characteristics at 25 versus 70 °C likely reflect thermal effects on the CO coverage and CO packing density and possibly the CO surface diffusion rate. The bands for the C-O stretching modes

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of adsorbed CO are expected to shift higher in energy with increasing positive potential on the basis of Stark tuning effects.38,39 However, similar upshifts can accompany an increase in CO surface coverage or CO packing density, because of vibrational coupling between adsorbates.40-42 In Figure 4a, growth in the CO coverage between 0.2 and 0.4 V can contribute to the band upshifts along with the Stark effect. Between 0.5 and 0.7 V, the shift in the atop CO band back toward lower energy is consistent with a decrease in the CO coverage in response to a faster rate of CO oxidation as the potential is stepped positive.23 It is not clear why the atop CO band position appears insensitive to potential at 70 °C (Figure 4b). However, the broadness of the peaks and the low signalto-noise ratio make it difficult to identify the band positions. Weak vibrational coupling is probably one factor that keeps the atop CO band positions at low energy. In addition to the improved oxidation kinetics and the possibility for CO thermal desorption at 70 °C,43-45 both of which can maintain the CO coverages as low, rapid CO surface diffusion16,17,46 can disrupt CO island formation and further diminish intermolecular interactions. Experiments analogous to those in Figures 4 and 5 were performed at intermediate temperatures, and the results are summarized in Figure 6. With increasing temperature, the integrated absorbance of the CO2 and the adsorbed CO spectral bands follow similar potential-dependent trends. Adsorbed CO appears between 0.2 and 0.6 V in all experiments, but the quantity of adsorbed CO in this potential range decreases above 25 °C. For CO2, in general, the integrated absorbance becomes progressively larger as the temperature is raised. Comparing parts A and B in Figure 6 shows that CO2 evolves from a surface that contains a partial monolayer of adsorbed CO. The CO that is detected can be an oxidation resistant poison at these potentials, or can reflect the steady-state coverage that results from the parallel reactions of methanol dissociative chemisorption and water activation leading to CO oxidation (vide infra).3,25,26 It is notable that the plots in Figure 6 correspond well with related in situ IR spectroscopy studies of methanol oxidation.23,25 In Figure 6, the integrated intensity of the adsorbed CO bands and the quantity of CO2 measured depend on the time for data acquisition at each potential. Advancing the potential at a faster rate reduces the amount of CO2 produced in an experiment and extends the range of potentials over which adsorbed CO is detected. The ambient temperature data in Figure 6 match the sweep rate dependence for methanol oxidation reported in reference 25;25 the speed for acquisition of data in Figure 6 was just below that for the 1 mV/s sweep rate tested in reference 25. Solutions of 1.0 M Methanol. In 1.0 M methanol solutions, the adsorbed CO vibrational bands observed at (35) Adzic, R., In Modern Aspects of Electrochemistry; White, R. E., Bockris, J. O. M. Conway, B. E., Eds.; Plenum Press: New York, 1990; Vol. 21, p 163. (36) Franaszczuk, K.; Herrero, E.; Zelenay, P.; Wieckowski, A.; Wang, J.; Masel, R. I. J. Phys. Chem. 1992, 96, 8509. (37) Chang, S. C.; Ho, Y.; Weaver, M. J. Surf. Sci. 1992, 265, 81. (38) Weaver, M. J. Appl. Surf. Sci. 1993, 67, 147. (39) Chang, S. C.; Weaver, M. J. J. Phys. Chem. 1991, 95, 5391. (40) Severson, M. W.; Stuhlmann, C.; Villegas, I.; Weaver, M. J. J. Chem. Phys. 1995, 103, 9832. (41) Persson, B. N. J.; Ryberg, R. Phys. Rev. B: Solid State 1981, 24, 6954. (42) Chang, S. C.; Weaver, M. J. J. Chem. Phys. 1990, 92, 4582. (43) Collins, D. M.; Spicer, W. E. Surf. Sci. 1977, 69, 85. (44) McCabe, F. W.; Schmidt, L. D. Surf. Sci. 1977, 66, 101. (45) Borup, R. L.; Sauer, D. E.; Stuve, E. M. Surf. Sci. 1997, 374, 142. (46) Petukhov, A. V.; Akemann, W.; Friedrich, K. A.; Stimming, U. Surf. Sci. 1998, 402-404, 182.

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Figure 8. Same as Figure 5 except the solution contained 1.0 M methanol in 0.1 M HClO4.

Figure 6. Plot of the integrated vibrational band intensities for atop-coordinated adsorbed CO (A) and CO2 (B) derived from 1.0 × 10-1 M methanol as a function of potential at the indicated temperatures. Spectral acquisition parameters were the same as those described in Figures 4 and 5 for CO and CO2, respectively. Experimental temperatures are specified in the inset of part B.

Figure 7. Same as Figure 4 except the solution contained 1.0 M methanol in 0.1 M HClO4.

25 and 70 °C (Figure 7) are stronger relative to those measured using 1.0 × 10-1 M methanol at corresponding temperatures. At 25 °C, the integrated band intensities between 0.3 and 0.7 V range from 70-90% of the values for adsorbed CO measured in CO-saturated 0.1 M HClO4. The high CO coverages, which the intense spectral bands suggest, appear to inhibit water activation and methanol dissociative chemisorption to some extent. Despite the 10-fold increase in methanol concentration, the linear

sweep voltammetry of 1.0 M methanol between 0.0 and 0.5 V (data not shown) was nearly identical to the scans reported in Figure 2b for 1.0 × 10-1 M methanol. Related disparities in the concentration dependence of methanol voltammetry on platinum have been reported.19 The spectra in Figure 7a show that, at 25 °C in 1.0 M methanol, the equilibrium in eq 1 is shifted strongly toward methanol adsorption and subsequent dissociation. Furthermore, the intense spectral bands above ∼0.5 V indicate that high coverages of CO persist at potentials usually associated with rapid CO oxidation. We have observed similar effects at 25 and 50 °C in experiments with 0.3 M methanol.47 Also, the CO2 produced from the reaction of 1.0 M methanol between 0.3 and 0.5 V (Figure 8a) is less than the quantities observed under the same conditions with 1.0 × 10-1 M methanol (Figure 5a). Thus, at least at 25 °C in 1.0 M methanol, the surface appears strongly poisoned by high coverages of adsorbed CO. Comparing the spectra in Figure 7a and b indicates that the methanolic CO coverages are lower at 70 °C than near ambient. Between 0.2 and 0.5 V, the integrated intensities of the adsorbed CO bands in Figure 7b range from 40-50% of the values attained at this temperature in CO-saturated 0.1 M HClO4. At 70 °C, improved water activation kinetics and rapid CO desorption from terrace structures on platinum43-45 can combine to lower CO coverages. As expected, the rate of CO2 evolution is faster at 70 °C (Figure 8b) than at 25 °C (Figure 8a) at all potentials studied. It is also interesting to compare the quantities of CO2 formed from 1.0 M methanol and 1.0 × 10-1 M methanol at 70 °C (Figures 8b and 5b). CO2 evolution is slower in 1.0 M methanol, consistent with the more intense adsorbed CO bands (Figures 4b and 7b) and lower than expected voltammetric currents. Discussion This study examined pathways in the electrochemical oxidation of methanol on polycrystalline platinum at 2570 °C. Methanol oxidation was probed by following the vibrational bands of the complete oxidation product (CO2) and the adsorbed intermediate (CO) with changes in the (47) Kardash, D.; Huang, J.; Korzeniewski, C. Langmuir 2000, 16, 2019.

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Figure 9. Plots of the integrated intensity of the CO2 vibrational band as a function of potential for the indicated methanol concentrations. (A) The potential region between 0.3 and 1.0 V is shown. Spectra were recorded with the cell at 25 °C. The concentrations specified in the inset are for methanol. (B) The potential region between 0.3 and 0.6 V is shown. The inset specifies the experimental temperatures and methanol concentrations.

experimental variables of temperature, potential, and methanol concentration. The spectroscopic results are consistent with the conclusion of Gasteiger and co-workers on the basis of voltammetry measurements showing that the thermal activation of water dissociation on platinum can decrease the steady-state coverage of surface poisons and thereby increase the rate of methanol oxidation.20 For methanol concentrations between 1.5 × 10-2 and 1.0 M, increasing the temperature from 25 to 70 °C lowered the quantities of adsorbed CO detected and shifted the onset potential for methanol oxidation negative. The results also confirm the notion that the rate of methanol oxidation is controlled by a competition between methanol adsorption from solution and the removal of incomplete reaction products that poison the surface.7,8,19 For the lowest methanol concentration studied, 1.5 × 10-2 M, methanol adsorption appears to limit the reaction, at least between 25 and 60 °C. Raising the methanol concentration to 1.0 × 10-1 M led to corresponding increases in CO coverage, consistent with the shift in the equilibrium in eq 1. The CO2 formation rate also became greater (Figures 3 and 5), showing that adsorbed CO did not completely eliminate gains in the CO2-producing pathways allowed by faster methanol adsorption. However, in 1.0 M methanol, CO2 formation was slower than in 1.0 × 10-1 M methanol below 0.9 V at the temperatures studied (Figure 9). Figure 9b shows the inhibiting effects at low potentials. At both 25 and 70 °C, CO2 formation is fastest in 1.0 × 10-1 M methanol solutions. Since the CO surface coverages appear to be higher in 1.0 M than in 1.0

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× 10-1 M methanol solutions (Figures 7 and 4), poisoning by CO is believed to be at least partially responsible for limiting CO2 formation below 0.9 V. Between 0.9 and 1.0 V, the quantity of CO2 produced from 1.0 M methanol exceeds that from 1.0 × 10-1 M methanol (Figure 9a). It appears that the coverage of surface poisons is low at these potentials, probably because the electrochemical driving force for water activation at 0.9 -1.0 V allows for a fast rate of reactive oxide formation and subsequent CO oxidation (eqs 2-3). Thus, CO2 evolution from 1.0 M methanol becomes greater than that from 1.0 × 10-1 M methanol at high positive potentials as methanol adsorption becomes rate limiting. In Figure 9b, the effect of temperature on the water activation kinetics is also demonstrated. Comparing the same methanol concentrations, CO2 appears at lower potentials and is produced in greater quantities at 70 °C than at 25 °C. The present experiments do not allow the nature of the activated oxygen-containing species (i.e., water, hydroxide, or oxide) to be identified. Important efforts in the study of methanol electrochemical oxidation are addressing this issue.7,9 Additional findings relate to the role of CO as a surface poison versus a reactive intermediate on platinum below about 0.6 V.9,19,20,26 At the potentials studied (0.2-0.6 V), the adsorbed CO characteristics depend on the methanol concentration and cell temperature. In general, large gains in current and CO2 formation coincide with a loss of detectable adsorbed CO. This behavior is consistent with the surface-poisoning effects of CO. However, Figures 4-9 demonstrate how the initial stages of CO2 production take place at potentials where CO coverages remain high. Similar findings have been reported for methanol adsorption on platinum electrodes at ambient temperatures.23,25 It is possible that, on the highly CO covered Pt surfaces, water activation starts at defects in the CO adlayer. This type of response was predicted in recent experimental46 and theoretical48 studies of CO oxidation. Under these conditions, the adsorbed CO coverage and quantity of CO2 produced will depend on the balance between the oxidation of CO near sites of water activation and the rate at which CO is replenished by subsequent methanol adsorption. The CO involved in this cycle likely can be regarded as a reactive intermediate. In relation to this idea, diffusion of CO to sites of water activation may also be an important mechanistic step below about 0.6 V.16,17,48,49 It has been suggested that, at low potentials on platinum, water activation occurs preferentially at low coordination steps and kinks.46 The rate of diffusion to these sites is expected to increase with temperature, and indeed CO2 formation begins at less positive potentials and is faster at 70 °C than at 25 °C (Figures 5 and 8). Another important issue in methanol surface electrochemistry concerns the significance of non-CO partial oxidation products as intermediates and surface poisons (cf. references 1, 2, 9, 23, 27, 28, 50). Pathways to CO2 involving the reaction of activated oxides with incomplete methanol dehydrogenation products, such as COH and CHO,9,27,28,50 are believed to be important, particularly for avoiding CO formation and the deleterious effects of adsorbed CO. Vibrational bands of these non-CO adsorbates were not detected in the present work. However, the intensities of the adsorbed CO bands in Figures 4 and (48) Koper, M. T. M.; Jansen, A. P. J.; van Santen, R. A.; Lukkien, J. J.; Hilbers, P. A. J. J. Phys. Chem. 1998, 109, 6051. (49) Koper, M. T. M.; Lukkien, J. J.; Jansen, A. P. J.; van Santen, R. A. J. Phys. Chem. B 1999, 103, 5522. (50) Xia, X. H.; Iwasita, T.; Ge, F.; Vielstich, W. Electrochim. Acta 1996, 41, 711.

Temperature Effects at Platinum Electrodes

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7 indicate that the CO surface coverages do not reach complete saturation, especially at 70 °C. Electrodes that have low methanolic CO coverages but appear strongly poisoned toward methanol oxidation are believed to be affected by non-CO, methanol-derived adsorbates.50,51 Furthermore, it is thought that methanol oxidation to CO2 through partially dehydrogenated intermediates, especially formyl (CHO), may be more energetically favorable than through adsorbed CO.1,2,4,8,52 The in situ

measurements presented here leave open possibilities for electrode poisoning by non-CO adsorbates and CO2 formation, in part, through pathways that do not involve adsorbed CO.

(51) Chang, S. C.; Leung, L. W. H.; Weaver, M. J. J. Phys. Chem. 1990, 94, 6013.

(52) Jarvi, T. D.; Sriramulu, S.; Stuve, E. M. J. Phys. Chem. 1997, 101, 3649.

Acknowledgment. The Office of Naval Research is gratefully acknowledged for support of this work. LA000544E