Formic Acid Decomposition on Polycrystalline Platinum and

9700

J. Phys. Chem. B 1999, 103, 9700-9711

Formic Acid Decomposition on Polycrystalline Platinum and Palladized Platinum Electrodes Guo-Qiang Lu, Alechia Crown, and Andrzej Wieckowski* Department of Chemistry, UniVersity of Illinois at Urbana-Champaign, Urbana, Illinois 61801 ReceiVed: July 7, 1999; In Final Form: August 31, 1999

This is a comprehensive study in which a formic acid decomposition reaction is examined as a probe of catalytic properties of polycrystalline platinum and palladized platinum electrodes. The electrode potential varies in a broad range, and the reaction is carried out in perchloric acid and sulfuric acid solutions containing different concentrations of HCOOH. Analytical methods used to access the decomposition reaction are chronoamperometry and cyclic voltammetry. At very short times, we prove that only a negligible amount of surface CO is formed, and the CO unaffected decomposition reaction, leading to CO2 formation, can be interrogated. Surprisingly, the decomposition reaction displays Tafel behavior only in a very narrow potential range. This observation, made with both clean Pt and Pt/Pd electrodes, suggests that water-surface interactions, and/or (bi)sulfate-surface interactions, increase with increasing electrode potential and create a steric/electronic barrier for the decomposition of formic acid (and methanol, J. Phys. Chem. 1994, 98, 5074). We therefore offer a pessimistic view about platinum as a universal material for heterogeneous catalysis applications involving rearrangements of organic molecules. Such rearrangements may only be fulfilled with a low electrochemical driving force, at least at room temperature, but at higher potentials, the electrode becomes deactivated due to the unique attributes of the double layer structure on the platinum electrode. We have also found that the deceleration of formic acid oxidation (to CO2) is primarily due to CO chemisorption only at potentials overlapping with those from the hydrogen adsorption range, or not too positive from this range. At more positive potentials, the decay in formic acid decomposition is neither due to CO formation nor to solution mass transfer limitations. The presence of interfacial CO2 (J. Electroanal. Chem. 1994, 376, 151) or adsorption of formic acid and/or formate anion, could account for the decay. Finally, a detailed analysis of kinetic isotherms involved in the two pathways, CO2 formation and CO chemisorption, is made and the mechanism of formic acid decomposition on platinum is discussed. The electrolyte anion effects involved in formic acid oxidation in HClO4 and in H2SO4 solutions are also presented.

I. Introduction Catalytic decomposition of formic acid at platinum is a subject of numerous studies that have been reviewed periodically1-5 (where ref 4 deals with HCOOH decomposition in UHV). The reaction has been perceived as a clear manifestation of the socalled dual path mechanism; see the review articles from above and the work done by Parsons and others.6-8 Namely, at double layer potentials, formic acid decomposes to CO2 and, in parallel, to chemisorbed CO. The CO molecule occupies the Pt sites and suppresses the overall decomposition reaction. However, in contrast to methanol decomposition,9-11 the CO2 formation pathway is much more efficient than that of CO formation. The fact that the CO poisoning channel is inefficient, or that the CO2 formation process occurs at high yield, gives a scientist an opportunity of examining the decomposition of an organic molecule with very little interference by chemisorption of molecular organic fragments. Avoiding chemisorption while examining the HCOOH decomposition reaction (to CO2) is the opportunity that we want to explore in this project. The decomposition involves an interfacial electron transfer as well as a chemical bond splitting on Pt sites, and both processes depend, or may depend, on the chemical state of the electrode surface. * To whom correspondence should be addressed. E-mail: andrzej@ scs.uiuc.edu.

We focus on the decomposition reaction at short times where the accumulation of CO chemisorption can be considered negligible. However, the determination of rate parameters that are indifferent to chemisorption requires a complete understanding of the full process that, eventually, i.e., at longer times, is strongly affected by CO chemisorption. Therefore, in this paper, we present results of a systematic study of the formic acid reaction, including modeling, with a polycrystalline platinum electrode in a broad range of potentials and reaction times. We want likewise to be able to discriminate between the two (traditionally12) controlling rate-limiting steps: kinetic (involving the interfacial electron transfer and the chemical bond breaking) and mass transport (diffusion of formic acid molecules to the electrode surface). We also explore the rates obtained on clean platinum in the context of formic acid decomposition on palladized platinum electrodes.13,14 As shown in the quoted papers, modifying platinum by palladium causes a significant acceleration of the decomposition kinetics, which we confirm. In this study, the rates at the palladized electrode give us a contrasting case to which the kinetics obtained with clean platinum are referred. The combination of chronoamperometry and fast cyclic voltammetry allow us to control the extent of CO chemisorption as the formic acid decomposition reaction unfolds.15 That is, at any time, we may correlate the reduction in the oxidation current to information about the surface CO uptake. On clean platinum

10.1021/jp992297x CCC: $18.00 © 1999 American Chemical Society Published on Web 10/15/1999

Formic Acid Decomposition on Platinum Electrodes and at electrode potentials located in the outmost negative part of the double layer potential range, we have found that the deceleration of formic acid oxidation to CO2 is primarily due to CO chemisorption. At more positive potentials, however, despite a negligible CO uptake, the decay in formic acid oxidation current still persists, and yet the rate decay is not due to the solution mass transfer limitations. On the palladized electrodes, the connection between the CO chemisorption and CO uptake is even less apparent. We have previously reported major successes with platinum surface modifications by some noble metal additives such as ruthenium15-20 and osmium21 using an electroless (spontaneous) deposition process. For instance, we have prepared highly active platinum/ruthenium single-crystal surfaces, and used them as electrocatalysts for methanol oxidation. In principle,16 and in practice,18,22 such surfaces are ideal substrates to simulate behavior of real-world methanol oxidation fuel cell catalysts,23 and the experimental data obtained from such surfaces will be used broadly, we believe, for catalyst modeling and theory. Means for using the spontaneous deposition of palladium on platinum were reported by Attard and Bannister24 and Llorca et al.,13,25 and we utilized their method to produce the palladized polycrystalline platinum surfaces for this project. However, only very small amounts of palladium can be added to platinum using this method in a straightforward manner. Our formula for obtaining higher palladium coverage of spontaneously deposited palladium is reported in the Experimental Section of this paper. II. Experimental Section 1. Platinum Electrode. A platinum rod of 2 mm in diameter was cut and polished as a polycrystalline Pt working electrode, a meniscus contact was employed to allow a fixed geometric working area, and the actual surface area was determined from the hydrogen adsorption-desorption charge.26 2. Combined Chronoamperometry and Fast Cyclic Voltammetry Method. Combined chronoamperometry (CA) and fast cyclic voltammetry (CV) measurements15 were carried out using a PAR 273 potentiostat interfaced to PC, and data were acquired using our own setups in a Headstart program. 3. Preparation of Pt/Pd Surface. It is well-known that palladium adlayers on platinum surfaces (Pt/Pd) can be obtained by spontaneous adsorption,24 forced deposition,13 or electrolytic deposition.14 In this paper, we have developed a method of a repeated spontaneous deposition. The procedure was as follows: The Pt electrode was (1) electrochemically cleaned in 0.1 M H2SO4 solution by cycling between 0 and 1.5 V vs RHE, ending at a potential in the double layer range; (2) immersed in the palladium nitrate solution (5 mM Pd2+ + 0.1 M H2SO4) for 5 min; (3) rinsed sufficiently with Milli-Q water; (4) electrochemical annealed by sweeping between 0 and 0.95 V for five scans in 0.1 M H2SO4. The above procedure was repeated four times such that the Pt/Pd surface was prepared by five cycles of spontaneous adsorption. Figure 1 compares the voltammograms of Pt/Pd (dot-dot-dashed line) and clean Pt (solid line) surfaces, at the palladium coverage of 0.65 ML. (See Appendix for description of the method of palladium coverage determination using X-ray photoelectron spectroscopy.) Voltammograms for some of the intermediate Pt/Pd surfaces are also shown, i.e., after one cycle (dotted line) and three cycles (dashed line) of palladium adsorption. It was observed that with palladium deposited on the surface, the two pairs of hydrogen adsorption-desorption peaks distinct of polycrystalline Pt merged into one pair at around 0.12 V. It was also noted that the oxide formation occurred at a more negative potential and

J. Phys. Chem. B, Vol. 103, No. 44, 1999 9701

Figure 1. Voltammograms recorded at various stages upon palladium deposition. Clean Pt (solid line); Pt/Pd surface obtained by one cycle (dotted line), three cycles (dashed line), and five cycles (dot-dot-dashed line) of immersion in a 5 mM Pd2+ solution in 0.1 M H2SO4. Scan rate was 200 mV s-1.

the double layer charges were compressed. Although palladium is stable in H2SO4 solutions up to 0.95 V, it can be removed by sweeping to higher potentials.24 4. Chemicals and Other Experimental Conditions. The chemicals used were sulfuric, perchloric, and formic acid (GFS, double distilled from Vycor), palladium nitrate (JohnsonMatthey), and Milli-Q water. All experiments were carried out at an ambient temperature of 25 ( 2 °C. The potentials were measured against a silver/silver chloride electrode but are reported versus RHE in the present paper. III. Results and Discussion 1. Cyclic Voltammetry. A cyclic voltammogram (0.2 V s-1) for Pt in 0.1 M HCOOH + 0.1 M H2SO4 is shown in Figure 2A (solid line). It is in good agreement with those previously reported in the literature.6,27-29 Three anodic peaks located near 0.6, 0.9, and 1.4 V and one cathodic peak near 0.6 V are observed. Since it is well-known that the electrooxidation of HCOOH on a Pt electrode yields CO2 (and H+) as the final product30-32 and CO is the overwhelmingly predominant stable intermediate,3,5,8,33,34 it is apparent that all the noticeable current concerning HCOOH decomposition can be attributed to the oxidation of HCOOH and/or CO to CO2 (note that there is no faradaic current involved in the CO formation pathway HCOOH f CO + H2O). The hydrogen adsorption-desorption charges are significantly suppressed, indicating that surface active sites have been noticeably blocked. The first anodic peak (near 0.6 V) is due to the oxidation of HCOOH to CO2 on surface sites that remain unblocked by CO.6 As for the second anodic peak (near 0.9 V), the oxidation of surface CO is obviously involved. However, the possible value of CO oxidation charge alone is far less than the charge observed in this peak, suggesting contribution from HCOOH oxidation which increases significantly as a consequence of the release of surface sites by CO removal. After the formation of surface oxides, the electrode becomes quite inactive. At even higher potentials, some catalytically active surface oxides are formed, accounting for the third anodic peak near 1.4 V. On the negative-going sweep, the surface stays inactive until partial reduction of the irrevers-

9702 J. Phys. Chem. B, Vol. 103, No. 44, 1999

Figure 2. (A) Cyclic voltammograms of Pt(poly) in 0.1 M HCOOH with 0.1 M H2SO4 (solid line) and with 0.1 M HClO4 (dotted line). Scan rate was 200 mV s-1. (B) Cyclic voltammograms of Pt(poly) in 0.01 M HCOOH with 0.1 M H2SO4 (solid line) and with 0.1 M HClO4 (dotted line). Scan rate was 200 mV s-1.

ibly formed surface oxides. The cathodic peak demonstrates the real catalytic activity of the Pt surface, since neither CO nor oxide exists noticeably on the surface. It is interesting to compare the voltammogram obtained in perchloric acid with that in sulfuric acid, as illustrated in Figure 2A. The main features of the two voltammograms are very similar. However, anion influence is clearly demonstrated. The second anodic peak and the cathodic peak are higher and extend to more positive potentials in perchloric acid (dotted line) than that in sulfuric acid, indicating that anion effects seems to occur at least at potentials above 0.35 V. Since there is no noticeable CO adsorbed on the surface during the cathodic scan, the relative height of the two peaks implies that Pt is more active in HClO4 than in H2SO4 toward HCOOH oxidation as far as there exists the same amount of free surface sites. Thus, the fact that the first anodic peaks are the same in height may suggest that more surface sites are blocked (by CO) in HClO4 than in H2SO4. In this sense, anions interfere with the CO uptake process (see also Figure 7), which occurs in the low potential region. In fact,

Lu et al. anion influence in the low potential region can also be seen in the hydrogen adsorption-desorption features. Figure 2B compares the cyclic voltammograms for the Pt surface in the two supporting electrolytes, but containing only 0.01 M HCOOH. We can see that the oxidation current densities are much smaller, and as a consequence, the features due to surface processes are exhibited more clearly. Nevertheless, the results confirm the conclusions that have been reached above from Figure 2A. 2. Combined Chronoamperometry and Fast Cyclic Voltammetry. The combination of chronoamperometry and fast cyclic voltammetry allows us to control the extent of CO chemisorption as the formic acid decomposition reaction unfolds.15 A typical CA/CV potential program is shown in the inset to Figure 3A, which includes presteps, CA and fast CV. Cleaning presteps were applied between the preset higher and lower potentials. The higher potential was selected at 0.95 V where surface CO poison was completely removed, the oxidation of HCOOH ceased to a negligible level (see below), while surface oxide was less significantly formed (∼0.25 ML). The lower potential was set at 0.01 V, which precedes the hydrogen evolution. The waiting time at each low or high potential was 1 s, and four cleaning steps were applied to achieve good reproducibility of the data. The chronoamperometric measurement (CA) immediately followed the cleaning presteps, commencing with a potential step from the higher potential limit, 0.95 V, to a selected potential. Thus, just before recording a current transient, (1) the electrode surface is free of CO poison; (2) there are only small quantities of surface oxide formed, which is indifferent after ca. 0.2 s or so during the CA measurement (see below); and (3) the local species concentrations near the surface can be considered very close to bulk concentrations since the surface was practically inactive at 0.95 V. The corresponding background current transients were also recorded under the same conditions in clean supporting electrolyte and, wherever needed, they were subtracted from those recorded in formic acid containing solutions. Figure 3A shows an example of the corrected CA plot for Pt at 0.3 V in 0.1 M HCOOH + 0.1 M H2SO4. To examine the formation of surface CO, a fast CV (50 V s-1) plot was recorded immediately following the chronoamperometric measurement, as shown in Figure 3B. The first positive going sweep (solid line) exhibited a pronounced oxidation peak around 1.2 V, which can be easily ascribed to the oxidative stripping of surface CO formed during the CA section (10 s at 0.3 V). The existence of surface CO also accounts for the significant suppression of current in the hydrogen and double layer potential range. Successive sweeps (dotted line, after CO removal) show no evidence of surface CO, as can be seen by the well-developed hydrogen adsorptiondesorption features. With a high sweep rate of 50 V s-1, the total time at low potential regions (below 0.45 V, where CO can build up on the Pt electrode5,7,8,35) is only 10-20 ms, so no noticeable surface CO accumulates. Therefore, the integrated charge obtained by subtracting the background (dotted line) current from the initial sweep current after the cross (at around 0.9 V) may be reasonably regarded as the CO oxidation charge (represented as the shaded area in Figure 3B). This technique is very convenient for acquiring and calculating the oxidation charge, and provides data with extremely good reproducibility. Furthermore, its accuracy has been strictly evaluated by an alternative calculation, wherein the CO oxidation charge was obtained by subtracting the oxide formation charge (obtained

Formic Acid Decomposition on Platinum Electrodes

Figure 3. (A) Typical CA curve for 0.1 M HCOOH in 0.1 M H2SO4 obtained with a PAR 273 potentiostat at 0.30 V. Inset shows the potential program of four cleaning presteps followed by a potential step down from 0.95 V to a given potential (in this case, 0.30 V) where the current density-time transient is recorded. Immediately following the CA step, the potential is stepped down to 0.01 V, at which point the fast (50 V s-1) CV is commenced. (B) Typical CV curve obtained using this program. The solid line shows the first sweep. The high oxidation current is attributed to the stripping of surface CO (in this case, accumulated at 0.30 V for 10 s). The dotted line shows the subsequent sweeps, upon removal of surface CO. The QCO value is the integrated charge obtained by subtracting the background (dotted line) current from the initial sweep current after the cross at around 0.9 V (see text for detail).

either from fast CV in clean supporting electrolyte, or from the dotted line in Figure 3B by choosing a proper baseline to eliminate the small contribution from HCOOH oxidation) from the charge under the peak at 1.2 V (again, by finding a proper baseline, the small contribution from HCOOH oxidation can be eliminated), yielding basically the same charge data (within ca. 3% of difference). So far, our technique provides a better approach for the on-line exploration of CO uptake on a Pt

J. Phys. Chem. B, Vol. 103, No. 44, 1999 9703 surface from HCOOH solutions, especially with relatively high concentrations (say, >0.01 M), in comparison to the dipping and isolating technique13,36 or to the potential stepping technique.35,37 It should be mentioned that we started all fast CV sweeps at 0.01 V (by a potential step to this value immediately following the CA measurement) in order to monitor the extent of the hydrogen adsorption suppression and for better comparability of data (all beginning with the common starting point). Such a negative potential step to 0.01 V did not give rise to any uncertainty and has been confirmed by systematic comparisons. This is obvious since CO is a very stable adsorbate at low potential regions (at least for short time). 3. Sampled Current Densities. To evaluate the activation parameters, instantaneous current densities may be the best data to utilize. However, they are not directly available from the experiment (see below) since they can only be obtained via modeling the experimental data (will be documented in the next subsection). Obviously, sampled current densities12 are the most straightforward alternatives and we want to examine them first. The current densities recorded in the beginning of the decay (say within 0.2 s) are not reliable because of the double-layer charging at the beginning of the step. On the other hand, in the catalytically significant potential region (i.e., 0.2-0.4 V), CO accumulates so fast that there exists a considerable amount of CO poison after a short period of time (e.g., after 2 s). Therefore, only data within a very limited time window (0.2-1 s) are possible candidates. We have found that using current densities sampled at 0.5 s is a good observable (and will be demonstrated in the next subsection). This is not surprising since there is neither uncertainty concerning double-layer charging needed for erasing the memory from the potential step nor noticeable CO on the surface. Figure 4A compares the sampled current data at 0.5 s for 0.1 M HCOOH in the two supporting electrolytes. It shows the real surface activities under the present experimental conditions. In each case, a well-shaped volcanic plot was observed, with a maximum near 0.57 V and with the current density decreasing successively at both sides of this potential. In 0.1 M HCOOH, therefore, it is at a potential around 0.57 V that the Pt surface exhibits its highest activity. However, in terms of the Tafel behavior of the HCOOH molecule (see Figure 5, the slope values are 140 mV dec-1 in sulfuric and 133 mV dec-1 in perchloric acid and are discussed in a later subsection), the deceleration of the oxidation reaction begins at a potential much lower than 0.57 V. This electrode potential is as low as 0.4 V in H2SO4 solutions for Pt and is as low as 0.15 V for Pt/Pd (see subsection III.7). For clean platinum, the drop of the catalytic performance at even lower potentials than 0.4 V may be ascribed to the decrease in the electrochemical driving force that favors the oxidation reaction (the Tafel behavior) and also to hydrogen adsorption. On the other hand, the deactivation at potentials higher than the platinum oxidation threshold should be accounted for by oxide formation. However, noticeable surface oxidation does not commence until around 0.75 V, demonstrating that other processes exist which are capable of changing surface properties in the potential region of 0.4∼0.75 V. In fact, from the cyclic voltammogram for Pt in HClO4 without HCOOH (not shown), within the potential range from 0.4 to 0.75 V, a small but clearly measurable current is superimposed on the double-layer current, which may be attributed to adsorption and partial discharge (polarization) of interfacial water. Our observations suggest that the strength of water-surface interactions tends to increase with increasing electrode potential, aiding the


Figure 4. (A) Current densities sampled at 0.5 s as a function of electrode potential for HCOOH oxidation (0.1 M) on a Pt electrode in 0.1 M H2SO4 (filled symbols) and in 0.1 M HClO4 (open symbols). (B) Current densities sampled at 0.5 s as a function of electrode potential for HCOOH oxidation (0.01 M) on a Pt electrode in 0.1 M H2SO4 (filled symbols) and in 0.1 M HClO4 (open symbols).

deactivation of the electrode toward HCOOH oxidation. In the case of H2SO4 solution, the water partial discharging process was suppressed, implying a stronger (bi)sulfate-surface interaction than that of the water-surface interaction in the doublelayer potential range. Despite the similar overall shape, noticeable differences can be found between the plots from the two solutions. That is, the current densities are significantly higher in HClO4 than that in H2SO4, especially at potentials above 0.4 V (Figure 4, open symbols). These results confirm again the specific anion effects. It is surprising, however, that the plot for 0.01 M HCOOH differs noticeably in shape from that for 0.1 M HCOOH, as displayed in Figure 4B. Unlike that in Figure 4A, wherein each plot shows a well-defined smooth peak, the profiles in Figure 4B exhibit more fine structures. Typically, the maximum shifts to a lower potential around 0.4 V and the current density values decrease very rapidly at potentials below 0.4 V and above 0.75 V, while they decrease much more slowly at those from 0.4 to 0.75 V. It is obvious that the deactivation below 0.4 V and above

Lu et al.

Figure 5. (A) Tafel plots for HCOOH oxidation (0.1 M) on a Pt electrode in 0.1 M H2SO4 (filled symbols and solid line) and in 0.1 M HClO4 (open symbols and dotted line). The obtained Tafel slope values are 140 and 133 mV dec-1, respecitvely. (B) Tafel plots for HCOOH oxidation (0.01 M) on a Pt electrode in 0.1 M H2SO4 (filled symbols and solid line) and in 0.1 M HClO4 (open symbols and dotted line). The obtained Tafel slope values are 176 and 164 mV dec-1, respecitvely.

0.75 V is accounted for by the electric field effects together with the increase of hydrogen adsorption and the surface oxidation, respectively, as also interpreted for data from 0.1 M HCOOH. In HClO4, the very distinct slow drop in surface activity from 0.4 to 0.75 V confirmed the gradual increase of water-surface interaction in this potential range, while the relatively faster deactivation observed in H2SO4 demonstrated the anion adsorption. It is necessary to point out that, in the case of a lower HCOOH concentration, the changes in the surface state in terms of surface-water and surface-anion interactions are amplified. Finally, results obtained here are in good agreement with those from voltammetric measurements (see subsection III.1 and Figure 2). 4. Formic Acid Oxidation Current Decay: Model of HCOOH Decomposition. To obtain the instantaneous current densities, and also to have a better understanding of the HCOOH decomposition process, it is essential to model the oxidation



Figure 6. Surface CO coverage profiles as a function of time (CO uptakes on a Pt electrode at various deposition times) in 0.1 M HCOOH with 0.1 M H2SO4 (filled symbols) and with 0.1 M HClO4 (open symbols). Plot shows data at two selected potentials: 0.24 V (circles) and 0.33 V (squares). It can be seen that the maximum coverage for both electrolytes nears 0.57.

current decay, especially in the Tafel potential region (0.240.36 V; see subsection III). In our initial attempt, as generally expected, we assumed that, in this potential region, the decay of HCOOH oxidation is accounted for by surface poisoning as a consequence of CO chemisorption,9 as required by the “dualpath” mechanism; see Introduction. We have also confirmed that the decay is not mass transport limited in this region since the calculated Cottrell current decay for the given conditions is much higher than the experimentally determined current decay, and as such, the current-time transients can be fitted using a site-blocking model. That is, the oxidation current decays as a function of surface CO coverage (since the CO blocking process involves no faradaic current formation, HCOOH f CO + H2O, as has been mentioned in subsection III.1),

j ) j0(1 - θCO)

m

(1)

where j0 is the instantaneous current density, θCO is the fraction of electrode surface covered by CO, and m is treated as an adjustable parameter in considering that the site-blocking effect may not be simply linear. The use of this model requires detailed monitoring of the CO accumulating process; that is, CO coverage profiles as a function of time (typical examples are as in Figure 6). The application of the kinetic isotherm treatment to the CO formation reaction,7 and to the data in Figure 6, suggests that the CO uptake process undergoes a second-order rate mechanism, 2 max dθCO/dt ) kCO ad (θCO - θCO)

(2A)

CO max 1/(θmax CO - θCO) ) kad t + 1/θCO

(2B)

where kCO ad is the rate constant for CO formation process, and is the maximum surface coverage of CO. Fittings accordθmax CO ing to eq 2B are shown in parts A and B of Figure 7 and they demonstrate that the rate of CO formation is reduced with the increase in the electrode potential, as reported previously,7 yielding the Tafel slope of ca. 110 mV/dec (Figure 8).

Figure 7. (A) Plots of 1/(θmax - θ) against time for surface CO uptake on Pt in 0.1 M HCOOH + 0.1 M H2SO4 at selected electrode potentials: 0.24 V (open circles), 0.27 V (filled triangles), 0.30 V (open triangles), 0.33 V (filled squares), 0.36 V (open squares). (B) Plots of 1/(θmax - θ) against time for surface CO uptake on Pt in 0.1 M HCOOH + 0.1 M HClO4 at selected electrode potentials: 0.24 V (open circles), 0.27 V (filled triangles), 0.30 V (open triangles), 0.33 V (filled squares), 0.36 V (open squares).

Inserting eq 2B back into eq 1 yields the expression

(

j ) j0 1 -

2 CO (θmax CO ) kad t CO 1 + θmax CO kad t

)

m

(3A)

Since θmax CO (∼0.57 in both sulfuric acid and perchloric acid supporting electrolytes for all the potentials within 0.24-0.36 V) and kCO ad (different slightly in sulfuric acid from those in perchloric acid solutions, but both in a reducing manner with increasing electrode potentials) have been obtained from the fittings for the CO formation processes, only j0 and m are treated as adjustable parameters in eq 3A. Still, the fitting of the chronoamperometric curves turned out to be very poor. That is, CO formation and coverage cannot account for the total current decay (or the current decay is not solely due to surface


Lu et al.

Figure 8. Tafel plots for surface CO uptake on a Pt electrode from 0.1 M HCOOH with 0.1 M H2SO4 (filled symbols and solid line) and with 0.1 M HClO4 (open symbols and dotted line). The obtained Tafel slope values are -105 and -112 mV dec-1, respectively.

CO formation) and the model should be revised. A straightforward modification is to postulate an additional surface siteblocking process by another adsorbate and also to assume that it follows the second-order mechanism,

dθ′/dt ) k′ad(θ′max - θ′)2

(2C)

where θ′ is the fraction of electrode surface covered by another adsorbates and θ′max is the maximum coverage, k′ad is the rate constant of its site blocking process. Accordingly, eq 3A then becomes

(

j ) j0 1 -

2 CO (θmax CO ) kad t CO 1 + θmax CO kad t

-

(θ′max)2k′adt 1 + θ′maxk′adt

)

m

(3B)

Now j0, θ′max, k′ad, and m are treated as adjustable parameters. This model gives a satisfying fit for the experimental chronoamperometric curves at potentials from 0.24 to 0.36 V in both sulfuric acid and perchloric acid supporting electrolytes. An example of the fit is shown in Figure 9 for 0.36 V in sulfuric acid, the R2 fit value for this curve was 0.99. All the instantaneous current densities (j0), maximum coverages (θ′max), rate constants (k′ad), power indexes (m), and the regression coefficients (R2) are given in Table 1. Figure 10 shows the Tafel plots using the instantaneous current densities obtained from this model. The Tafel slopes of 127 and 132 mV/dec in sulfuric acid and perchloric acid, respectively, were obtained. These values are in good agreement with those obtained using the sampled current density values (see above), suggesting that our model is satisfactory and also confirming that the current densities sampled at 0.5 s are good observables in this kinetic study. It should be stressed that the main value of the model developed above is in providing a reliable means for the instantaneous current determination. From the mechanistic point of view, the model is an oversimplification since it does not consider the possible surface heterogeneity. However, and quite surprisingly, the heterogeneity effects are relatively weak, as demonstrated by the data in Figure 7 of this section. We

Figure 9. Measured current density decay at 0.36 V and subsequent fitting results for 0.1 M HCOOH + 0.1 M H2SO4. Fitted equation uses a second-order two-site blocking mechanism (see text).

therefore believe that the model is a good approximation for the potential range studied in the sense that the current decay is predominately due to surface poisoning or a decrease in surface sites available for oxidation of formic acid into carbon dioxide. Also, while CO chemisorption undoubtedly occurs, our model suggests that, at these potentials, other surface adsorbates, interfacial CO2,13 and intermediate forms of formic acid (e.g., formate anion formed from a surface induced dissociation of formic acid)38,39 can also be attached to the surface or be present at the interface, and thus aid the current decay. The reaction is very sensitive to its local environment as pointed out by the electrode potential dependent exponent m, which may be as high as 3.6 (see Table 1). 5. Potentials above 0.4 V. Using the CA/CV procedure, we have found that the deceleration of formic acid oxidation (to CO2) is closely connected to CO chemisorption only at potentials practically overlapping with those from the hydrogen adsorption range (or not too positive from such potentials, i.e., below 0.4 V). At more positive potentials, the decay in formic acid is neither due to CO formation (since the CO surface coverage is negligible) nor to solution mass-transfer limitations. The current may decay mainly because of the interference of a species whose adsorption does not give a clear signature in the electrode potential range where surface CO is electro-oxidized, nor in the oxide range of the Pt electrode. Llorca et al. have already proposed that the CO2 formation process may deactivate the interface in this potential range, thus preventing the catalytic decomposition of the HCOOH molecule from further progress.13 Another possibility, however, is that either formic acid adsorption40 and/or the adsorption of a formate anion released from the forced dissociation of HCOOH at the platinum/solution interface,38,39 could be alternative candidates accounting for the decay. Direct spectroscopic data are needed to substantiate this latter assumption. 6. Mechanisms versus Activation Parameters. We have pointed out the Tafel slopes obtained from modeling are 127 mV/dec and 132 mV/dec in sulfuric acid and perchloric acid, respectively. This shows that the Tafel slopes are not far from 120 mV/dec, indicating simple interfacial kinetics involving a single electron transfer reaction (in the rate-determining step (rds)). While the range is too narrow to make a more advanced



TABLE 1: Instantaneous Current Densities, Maximum Coverage and Rate Constants for Anion Adsorption, Power Index, and Regression Coefficient Involved in HCOOH Oxidation; See Eq 3B 0.1 M HCOOH in 0.1 M H2SO4

0.1 M HCOOH in 0.1 M HClO4

E/V

j0/mA cm-2

θ′max

k′ad/s-1

m

R2

j0/mA cm-2

θ′max

k′ad/s-1

m

R2

0.24 0.27 0.30 0.33 0.36

0.46 0.68 1.2 2.4 3.6

0.25 0.33 0.42 0.43 0.43

15 9.9 7.9 7.2 4.9

2.5 2.3 2.3 3.1 3.5

0.98 0.97 0.96 0.98 0.99

0.53 0.83 1.3 2.5 4.1

0.25 0.35 0.39 0.43 0.43

17 10 8.6 6.3 4.4

2.8 2.6 2.5 3.0 3.6

0.99 0.97 0.96 0.97 0.98

Mechanisms explaining the inverted reaction rate-electrode potential relationship for the CO formation process (see Figure 7 and subsection III.4) were proposed by Capon and Parsons27 and, independently, by Wieckowski and Sobkowski.7 While these mechanisms need to be revised, since species other than CO stable intermediates3,5,8 were taken into account,7,8,27 their general principles may remain the same. Namely, the CO formation channel is associated with a reduction reaction of either the COOH radical8,27 or the adsorbed HCOOH molecule by hydrogen.7 In the latter case, rds,k2

(HCOOH)ads + H+ + e 98 [CH(OH)2]ads

(7)

[CH2(OH)2]ads f (CHO)ads + H2O

(8)

(CHO)ads f COads + H+ + e

(9)

V ) k2aH+[(HCOOH)ads] exp(RnFE/RT)

(10)

and Figure 10. Tafel plots using tabulated j0 values (see Table 1) obtained by fitting experimental curves and extrapolating to zero. The Tafel slope values are 127 mV dec-1 in sulfuric acid (filled symbols and solid line) and 132 mV dec-1 in perchloric acid (open symbols and dotted line).

analysis of the activation parameters reliable, it is still safe to conclude that the rate-determining step, reaction 4, is a C-H break on the HCOOH molecule, involving an associated electron transfer reaction: rds,k1

HCOOH 98 (COOH)ads + H+ + e fast

COOH 98 CO2 + H+ + e

(4) (5)

Step 4 is followed by the decomposition of the unstable intermediate COOH to CO2 (reaction 5) or to CO (see below). The rate equation therefore is

V ) k1[HCOOH] exp(RnFE/RT)

(6)

where the bulk concentration is constant during the reaction (at least at 0.1 M of HCOOH), and R is close to 0.5 (see above). This relatively simple picture is made more complicated however in the 0.01 M solutions of HCOOH (Figure 5B) because the Tafel slopes are significantly higher than 120 mV, approaching 180 mV/dec. Obviously, how the formic acid molecule is predisposed for the catalytic attack by the Pt sites at a distance close to, but not yet at, the surface affects the overall electric field effect on the reaction rate. At this point, interrogating the origin of the increase in the Tafel slope with the decrease in formic acid concentration would amount to speculation. We may just note for the record that the simple decomposition reaction of the HCOOH molecule, even unaffected by CO chemisorption, is a quite complex surface reaction, especially at low concentrations of formic acid in solution.

(where aH+ ) constant since constant pH solutions were used throughout this project). Therefore, a simple Tafel behavior with the approximate 120 mV/dec slope would be expected, as observed experimentally. No assumption is made about hydrogen adsorption since a different Tafel slope than observed would result. On the other hand, carboxylic groups of organic acids are indeed capable of coordinating platinum electrodes, as first reported by Horanyi41 and confirmed by Sobkowski and collaborators.42 It can therefore be proposed that formic acid is reduced in a preadsorbed state (and the adsorption of formic acid is electrode potential independent in the studied range). Overall, these mechanisms would indicate that formic acid is “amphoteric” with respect to surface redox reactions (can be either reduced or oxidized) that are associated with its catalytic decomposition on platinum. In contrast to the COOH formation process that senses a highly heterogeneous surface, the CO formation process can be described by simple Langmuirian adsorption kinetics involving two adjacent surface sites (see subsection III.4 and Figure 7). However, the disregard to surface heterogeneity is not fully justified since the kinetic isotherms show some convex-type deviations from the straight lines, indicative of lateral interactions among the surface CO products. Because the deviations are relatively small, we believe that through-space, dipoledipole interactions are operative between the adsorbed form(s) that, upon more advanced treatment, would yield Roginsky/ Zeldowitch kinetic isotherms43-45 or the equivalent.46,47 7. Pt/Pd Surface. Figure 11 shows the cyclic voltammogram obtained on a Pt/Pd surface in 0.1 M HCOOH and 0.1 M H2SO4. With the presence of irreversibly adsorbed palladium on the surface, the voltammogram differs from that for a clean Pt surface. First, no clear inhibition of HCOOH oxidation can be observed, implying that the formation of CO poison does not occur or occurs at a very slow rate and to a much less extent


Lu et al.

Figure 11. Cyclic voltammogram of a Pt/Pd surface in 0.1 M HCOOH with 0.1 M H2SO4. With the presence of Pd, the CV differs remarkably from that of a clean Pt surface. There is no clear presence of CO adsorbate, as the HCOOH oxidation is not inhibited. Scan rate was 200 mV s-1.

than on a clean Pt surface. This is conceivable at this stage since no noticeable CO stripping current is visible after the doublelayer HCOOH decomposition and will be further examined. Second, there is little difference between the anodic and the cathodic peaks. Third, the current maximum occurs near 0.3 V, a potential considerably lower than that for a clean Pt surface (cf. with Figure 2A). These results are consistent with those of Llorca et al.13 and Baldauf and Kolb et al.14 Selected chronoamperometric and fast cyclic voltammetric profiles on the Pt/Pd surface are presented in parts A and B of Figure 12, respectively. The current density-time transients in Figure 12A show clearly the decay of the densities with time, especially in the beginning and at 0.21 V. Therefore, the deactivation occurs also on a palladium-modified Pt surface under the potentiostatic conditions. This observation suggests that CO poison may also build up on a Pt/Pd surface, however, apparently more slowly and to a lesser extent. Notably, at 0.39 V, the current stabilizes at 120 s, indicating that the reaction is at steady state. The fast CV results in Figure 12B clarify the uncertainties concerning CO poisoning on a Pt/Pd surface. The solid line and dashed line correspond to the first anodic sweep after holding the potential for 120 s at 0.21 and 0.39 V, respectively. The successive second sweeps in the two measurements are identical and displayed in this figure as a dotted line. We can see that after holding the electrode potential at 0.21 V for 120 s, there is a considerable amount of CO accumulated on the surface. The CO coverage (θCO) obtained was 0.26, which is close to a maximum coverage of CO poison for the prepared Pt/Pd surface. Note that for the case of a clean Pt surface, CO poison approaches its maximum coverage (θmax CO ) 0.57) in approximately 20 s. We can see that θmax (Pt/Pd) is much CO (Pt) and agrees with the findings of Llorca et lower than θmax CO al. for Pd/Pt(100).13 Figure 12B shows also that, at 0.39 V, only a very small amount of CO (θCO ) 0.05) is detected after 120 s of accumulation. Since the reaction has reached a steady state at this potential, no more uptake of CO molecules, beyond the 120 s reaction time, is expected.

Figure 12. (A) Chronoamperometric current density-time profiles for a Pt/Pd surface in 0.1 M HCOOH + 0.1 M H2SO4 at 0.21 V (solid curve) and at 0.39 V (broken curve). (B) First anodic sweeps in the fast CV measurements for a Pt/Pd surface after holding at 0.21 V (solid line) and 0.39 V (dashed line) for 120 s. The subsequent sweep (dotted line) is also displayed. The obtained θCO values are 0.26 and 0.05, respectively.

The plots of current densities sampled at 0.5 s versus potentials are compared in Figure 13A for Pt/Pd and clean Pt surface in 0.1 M HCOOH + 0.1 M H2SO4. The features for the Pt surface have been previously described. Unambiguously, the Pt/Pd surface shows higher activity than the clean Pt surface, especially at potentials below 0.5 V. On the Pt/Pd surface, the current density maximum is located near 0.27 V, 0.3 V lower than that on a clean Pt surface. Once again, the drop of current densities at lower potentials is due to the electric field effect and the increase of hydrogen adsorption. Meanwhile, the deactivation at higher potentials can be related to the change of surface properties, such as the increase in the strength of water-surface and anion-surface interactions. These are much more pronounced on Pt/Pd than on clean Pt. Furthermore, the high activity demonstrates the synergetic effect,13,14 since the current densities are much higher than the simple addition of the fractional contribution from Pt and Pd surface sites.


Figure 13. (A) Current densities sampled at 0.5 s as a function of electrode potential for HCOOH (0.1 M) oxidation on a Pt/Pd surface in 0.1 M H2SO4 (open symbols) compared to those obtained on a Pt electrode (filled symbols). Below 0.5 V, the palladized Pt surface is much more active than the clean Pt. (B) Tafel plot for Pt/Pd (open symbols) compared with that for clean Pt (filled symbols), 0.1 M HCOOH in 0.1 M H2SO4. Tafel values are 136 and 140 mV dec-1, respectively.

Interestingly, Tafel slopes for the two surfaces (Figure 13B) are comparable (140 mV dec-1 for clean Pt vs 136 mV dec-1 for Pt/Pd), though they are observed in different potential ranges. However, the meaning of the Tafel slope on Pt/Pd in the traditional hydrogen potential range, given the unique type of interactions of hydrogen with palladium, has yet to be established. IV. Conclusions We have demonstrated control, by electrochemical means, of the decomposition process of formic acid on a platinum electrode at room temperature and provided contrasting rate data for the Pt/Pd electrode. We have made several major conclusions concerning the decomposition reaction on platinum. The reaction involves, we believe, several elementary processes (reactions 4, 5, and 7-9): the oxidation of formic acid to an unstable

J. Phys. Chem. B, Vol. 103, No. 44, 1999 9709 surface intermediate COOH, further decomposition of COOH to CO2, and a reduction process of the HCOOH molecule leading to surface CO formation. Thus, formic acid is amphoteric with respect to surface redox associated with its decomposition on platinum. Other processes that are essential to the HCOOH decomposition mechanism include water and anion adsorption on the Pt electrode. From our work, the traditional double layer range of the Pt electrode does not appear a promising rate-promoting host for oxidation reactions of organic molecules. Namely, the oxidation appears efficient only at the beginning of the double layer potentials (up to 0.4 V). At higher potentials, the rate of the formic acid decomposition reaction decelerates, and the deceleration continues until the surface reactivity toward formic acid is essentially terminated by the build-up of platinum surface oxides. Notably, similar deceleration phenomena have previously been reported for methanol decomposition on several platinum single-crystal surfaces, in particular on Pt(111).11,48 Still, methanol on a platinum electrode is a molecule difficult to oxidize directly to CO2, in contrast to HCOOH, so an alternative, rather than identical, reactivity pattern in response to the electrode potential change was expected. The fact that the two molecules experience the same reactivity constraints at a very similar potential thresholds indicates that the reactivity limitations are due to the properties of the surface rather than to those of the reagents. As we have already alluded to above, chemical barriers due to adsorption of water (in perchloric acid media) and due to specific adsorption of (bi)sulfate (in sulfuric acid solution) are the most obvious candidates to account for the electrode deactivation, the former being less detrimental than the latter. The water and anion barriers are effective despite that both adsorbates are involved in a fast exchange with the bulk species,49,50 indicating that a relatively weak steric hindrance (since the H2O and anion species should exchange with HCOOH easily) must be augmented by some electronic (electrostatic) effects. For instance, due to the increasingly stronger dipole moment of the surface species along with the increase in the electrode potential, an electrostatic barrier may develop that will affect the orientation of the molecule at the surface with the reactive end (the HC group) directed outward from the Pt sites. Notably, we have pointed out that, at potentials higher than 0.45 V, the decay in formic acid oxidation is neither due to CO chemisorption nor to solution mass-transfer limitations. We conclude that the current decay results from the interference by interfacial CO2 and/or adsorption of formic acid (formate anion), which are easily displaced by the Pt oxide formation process and gives no voltammetric signature in the potential range where surface CO is electro-oxidized. However, direct spectroscopic data are needed to substantiate this assumption. Whereas our study of the Pt/Pd electrode was limited, the data were meant to provide a useful platform for comparison with the Pt-HCOOH data, we have made several new observations and conclusions and confirmed several major pieces of previous data in the field. For instance, the rate of decomposition of formic acid on Pt/Pd is indeed much higher than on Pt.13,14 However, this observation is subject to the proviso that the comparison is made under voltammetric conditions (especially, upon the positive going sweep) or that too positive potentials are not used for chronoamperometric current measurements. We also show that a high-reactivity steady-state pattern is observed with Pt/Pd at a potential not exceeding 0.4 V (the current density in the range of a few milliamperes per cm2). Also, in agreement with work by other authors, the high surface activity demon-


Lu et al.

strates a synergetic effect, since the current densities are much higher than the simple addition of the fractional contribution from Pt and Pd surface sites. Finally, our data strongly suggest that a Pt/Pd surface is more sensitive than a clean Pt surface to water and anion adsorption as factors affecting HCOOH electrocatalysis. Acknowledgment. The contribution by Dr. Rick Haash to the XPS part of this project is highly appreciated. This work is supported by the National Science Foundation under Grant CHE 97-000963, by the Department of Energy Grant DEFG0296ER45439 administered by the Frederick Seitz Materials Research Laboratory at the University of Illinois, and by the Army Research Office Grant DAAH04-94-G-0055. Appendix: XPS Calibration of the Pt/Pd Surface It has been reported that palladium coverage of Pt/Pd surface, in monolayer equivalents (ML), can be estimated either from electrolytic charge14 or from the distinct modification of hydrogen adsorption-desorption or oxide reduction features in the voltammograms.13,24,25 However, the former method works only for electrolytically deposited surfaces, and the latter only for single crystal Pt surfaces. Therefore, neither works for polycrystalline Pt with spontaneously deposited Pd. Nevertheless, UHV techniques may provide reliable estimates and XPS was employed in the present study for the calibration of the Pt/Pd surface. X-ray photoelectron spectra were collected with a PHI model 5400 X-ray photoelectron spectrometer (Physical Electronics, Inc., Eden Prairie MN) equipped with small-area electron extraction optics, a spherical capacitor electrostatic energy analyzer, and a dual channel plate position sensitive detector. The samples were excited using achromatic Mg KR X-radiation, 1253.5 eV (15 kV, 400 W). Spectra were collected from a 1 mm2 area of the surface at a 45° emission angle (relative to surface). Survey spectra were collected using a pass energy of 178.95 eV, 1 eV/step, and multiplex spectra were collected using a pass energy of 35.75 eV, 0.1 eV/step. The samples were sputter cleaned with 3 kV Ar+ for 2 min using a PHI model 04-303 differentially pumped ion gun. The survey and selected multiplex spectra for Pt/Pd surface before and after sputtering are displayed in Figures14 and 15, respectively. Spectrum a in Figure 14 shows clearly the existence of spontaneously deposited Pd. The presence of oxygen and carbon, as contaminants from the transfer through atmosphere, was also determined. After sputtering, all adsorbates were completely removed from the surface, as illustrated by spectrum b, wherein only the Pt signal is visible. The Pt 4f (both before and after sputtering) and Pd 3d peaks (as showed in Figure 15) were used to quantitatively analyze the Pt/Pd surface. The calculation is based primarily on the attenuation mechanism.51 For an adsorbed species (element A) on a substrate (element B),

IAB /I∞B IAB

) 1 - θA + θA exp(-RA/λ sin φ)

I∞B

(11)

Where and are the photoelectron intensities (after correction for relative elemental sensitivities) of the substrate with attenuation only by adsorbate A and without attenuation, respectively. The symbols θA and aA are coverage and atomic diameter of the adsorbed element, λ is the electron mean free path, and φ is the emission angle. In any case, I∞B corresponds to the intensity of the substrate after sputtering. For a single

Figure 14. XPS spectra for the Pt/Pd surface before (A) and after (B) sputter cleaning with 3 kV Ar+ for 2 min. The spectra are shifted vertically for clarity.

Figure 15. XPS spectra of the Pt/Pd surface for the Pd 3d and Pt 4d peaks.

adsorbate system, IAB corresponds to the intensities of the substrate before sputtering. However, for more complex submonolayer systems (thus no attenuation among adsorbates), according to the approximation, n

IAi ) I∞B - IB ∑ i)1

(12)

where IAi is the intensity of the ith element, and IB is the intensity of the substrate with full attenuation, then ∞ IAi B ) IB - IAi

(13)

Therefore, in the present study, θPd can be calculated by

(I∞Pt - IPd)/I∞Pt ) 1 - θPd + θPd exp(-aPd/λ sin φ) (14) In our experiments, Pd coverage (θPd) was determined to be 0.65 ML.

Formic Acid Decomposition on Platinum Electrodes References and Notes (1) Capon, A.; Parsons, R. J. Electroanal. Chem. 1973, 44, 1. (2) Parsons, R.; VanderNoot, T. J. Electroanal. Chem. 1988, 257, 9. (3) Beden, B.; Leger, J. M.; Lamy, C. In Modern Aspects of Electrochemistry; Bockris, J. O., Conway, B. E., White, R. E., Eds.; Plenum Press: New York, 1992; Vol. 22, p 97. (4) Columbia, M. R.; Thiel, P. A. J. Electroanal. Chem. 1994, 369, 1. (5) Sun, S.-G. In Electrocatalysis; Lipkowski, J., Ross, P. N., Eds.; Wiley-VCH: New York, 1998; p 243. (6) Capon, A.; Parsons, R. J. Electroanal. Chem. 1973, 44, 239. (7) Wieckowski, A.; Sobkowski, J. J. Electroanal. Chem. 1975, 63, 365. (8) Sun, S. G.; Clavilier, J.; Bewick, A. J. Electroanal. Chem. 1988, 240, 147. (9) Franaszczuk, k.; Herrero, E.; Zelenay, P.; Wieckowski, A.; Wang, J.; Masel, R. I. J. Phys. Chem. 1992, 96, 8509. (10) Markovic, N. M.; Gasteiger, H. A.; Ross, P. N.; Jiang, X. D.; Villegas, I.; Weaver, M. J. Electrochemica Acta. 1995, 40 (1), 91. (11) Herrero, E.; Franaszczuk, k.; Wieckowski, A. J. Phys. Chem. 1994, 98, 5074. (12) Bard, A. J.; Faulkner, L. R. Electrochemical Methods; John Wiley & Sons: New York, 1980. (13) Llorca, M. J.; Feliu, J. M.; Aldaz, A.; Clavilier, J. J. Electroanal. Chem. 1994, 376, 151. (14) Baldauf, M.; Kolb, D. M. J. Phys. Chem. 1996, 100, 11375. (15) Tremiliosi-Filho, G.; Kim, H.; Chrzanowski, W.; Wieckowski, A.; Grzybowska, B.; Kulesza, P. J. Electroanal. Chem. 1999, 467, 143. (16) Chrzanowski, W.; Wieckowski, A. Langmuir 1997, 13, 5974. (17) Chrzanowski, W.; Kim, H.; Wieckowski, A. Catal. Lett. 1998, 50, 69. (18) Chrzanowski, W.; Wieckowski, A. Langmuir 1998, 14, 1967. (19) Chrzanowski, W.; Kim, H.; Tremiliosi-Filho, G.; Wieckowski, A.; Grzybowska, B.; Kulesza, P. J. New Mater. Electrochem. Systems 1998, 1, 31. (20) Herrero-Rodriguez, E.; Feliu, J. M.; Wieckowski, A. Langmuir, in press. (21) Crown, A.; Kim, H.; Lu, G.-Q.; Moraes, I. R.; Rice, C.; Wieckowski, A. J. New Mater. Electrochem. Systems, submitted. (22) Demongeot, F. B.; Scherer, M.; Gleich, B.; Kopatzki, E.; Behm, R. J. Surf. Sci. 1998, 411 (3), 249. (23) Ren, X.; Wilson, M. S.; Gottesfeld, S. J. Electrochem. Soc. 1996, 143, L12.

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Formic Acid Decomposition on Polycrystalline Platinum and

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