Catalytic Methanol Decomposition Pathways on a Platinum Electrode

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J. Phys. Chem. B 2000, 104, 5566-5572

Catalytic Methanol Decomposition Pathways on a Platinum Electrode G.-Q. Lu, W. Chrzanowski,† and A. Wieckowski* Department of Chemistry, UniVersity of Illinois at Urbana-Champaign, Urbana, Illinois 61801 ReceiVed: January 13, 2000; In Final Form: April 6, 2000

Through a combination of potential steps (chronoamperometry) and fast potential sweeps (cyclic voltammetry), temporal advancements involved in methanol decompositionsindicated by the decomposition chargesswere referred to the extent of CO poisoning. In a broad potential range, especially at short times, there is an excess of charge above that needed for simple methanol dehydrogenation to surface CO. This indicates that there exist some efficient methanol decomposition pathways that lead to the formation of products other than CO, namely carbon dioxide, but also formic acid and/or formaldehyde. Because the formation of CO2 was not observed by others below the electrode potential of 0.35 V vs RHE (cf. infrared studies reported in J. Phys. Chem. B 1997, 101, 7542), we believe that, at low potentials, methanol dehydrogenation to formic acid and/ or formaldehyde accounts for the excess oxidation charge. The formation of the dissolved products along with surface CO confirms the applicability of a dual-path mechanism for methanol oxidation on platinum. We also successfully modeled the mechanism of methanol decomposition using a three-term rate equation. The modeling allowed us to estimate the rates of the elementary reaction channels involved in methanol decomposition without the complication of CO poisoning effects. We believe that the results reported in this paper lead to an advanced understanding of the methanol decomposition processes on a polycrystalline platinum electrode, especially at short reaction times. Information about such short-time methanol decomposition events was missing from the previous studies of this important electrocatalytic system.

1. Introduction Three organic molecules, from the series of CH3OH, HCHO, and HCOOH, catalytically decompose on platinum electrodes, producing carbon dioxide and interfacial electrons that are of interest fuel cell applications. Although all three molecules can, in principle, be considered for fuel cells, methanol is the best realistic option, as it has an acceptable toxicity level for largescale operations, a broad delivery base, and a high energy density. Paradoxically, from among those three molecules, methanol is the least reactive at platinum electrodes and generates surface-bound intermediates (catalytic poison) easily.1-8 Surface CO seems the most tenacious and most probable intermediate leading to the deactivation of platinum at room temperature or at moderate temperatures. Adding ruthenium or/ and other elements to platinum partially circumvents the CO poisoning quandary,9-22 but methanol decomposition rates are still too low to guarantee success in future development of direct methanol fuel cells.23-25 Fundamental research on the methanol decomposition process at catalytic electrodes is needed to capitalize on the momentum of public interest in alternative energy sources, especially for clean public transportation technologies.26-29 A variety of surface reactions combined with interfacial electron transfer occur when methanol decomposes on platinum in acidic media. They include the formation of CO as the main surface chemisorbed product, CO2 as the main bulk product (together with the release of six electrons per one methanol molecule), and side products, i.e., surface adsorbates other than * Author to whom correspondence should be addressed. E-mail: andrzej@ scs.uiuc.edu. † Permanent address: Technical University of Gdansk, 80-952 Gdansk, Poland.

CO or bulk products other than CO2, such as formic acid and formaldehyde.30-41 The variety of interfacial events has traditionally been labeled as a dual-path mechanism for methanol oxidation on platinum.42-47 This mechanism captures the idea that the surface bound products (1st path) and the released solution products (2nd path) can be formed simultaneously, at least in some potential range. Research by numerous electrochemical and surface science groups focused on the catalytic decomposition of methanol on platinum has led to several major conclusions. Namely, (i) the decomposition of methanol, following its catalytic activation, is complete and irreversible;1-3 (ii) the decomposition process produces a current transient that either decays to zero at low potentials (below 0.3 V vs NHE) or attains a steady-state value at higher potentials;14,16,44,45,48-51 (iii) above a critical potential threshold, surface CO is oxidized to CO2; (iv) the methanol decomposition rate is electrodepotential-dependent, whereas the nature of the chemisorbed products is apparently electrode-potential-independent; (v) the rate of methanol decomposition is surface-structure-sensitive and is strongly affected by the type of anions interacting with the electrode surface;44,49-54 and (vi) formic acid and/or formaldehyde can be formed as methanol oxidation intermediates.30,35,36,39 Nonetheless, several issues still remain unclear, among those are the following: (i) How does the decomposition reaction proceed at short times when methanol interacts with a CO-free (pristine) platinum electrode? (ii) How does the ratio of the surface to dissolved product depend on the reaction time and the electrode potential? (iii) What is the nature of the source of oxidant transforming surface CO to CO2?50,55 From these several outstanding issues, our focus is to investigate the distribution of methanol decomposition channels, leading to the formation of different product, as it depends on the time of the methanol/platinum interaction and on the

10.1021/jp000193c CCC: $19.00 © 2000 American Chemical Society Published on Web 05/18/2000

Catalytic Methanol Decomposition Pathways

Figure 1. Cyclic voltammograms for a polycrystalline Pt electrode (real surface area is 0.0453 cm2) at a sweep rate of 35 V s-1. Solid lines represent the curves obtained in 0.1 M H2SO4 containing 0.6 M (A) CH3OH or (B) CD3OD; dashed lines are background curves taken in clean supporting electrolyte. Insets show the activity in the doublelayer potential range.

electrode potential. Fulfilling this goal required the collection of methanol oxidation rate data in the time window from a few milliseconds to a few seconds using a combination of chronoamperometry and fast cyclic voltammetry techniques.57 We found that, in the first tens of milliseconds, even at potentials that border the hydrogen underpotential deposition region, methanol decomposed to soluble products at high yield. 2. Experimental Section A polycrystalline platinum electrode, 0.2 cm in diameter, was used in the measurements. The working surface was chemically and electrochemically cleaned via standard procedures. The real surface area was determined using hydrogen adsorptiondesorption charges.56 A two-compartment electrochemical cell equipped with a Luggin capillary and a PAR 273 potentiostat interfaced to a PC were used. The iR drop was measured and well compensated for by using a positive feedback option of the potentiostat. All potentials were measured against a reversible hydrogen electrode (RHE). The potential pulse-sweep sequence was obtained and data were collected by utilizing PAR HEADSTART software with our own setups at sampling rates of up to 5000 points per second (200 µs per point). Chemicals were Millipore water (18 MΩ cm), high purity sulfuric and perchloric acids (double distilled from Vycor, GFS Chemicals), ACS-certified methanol (Fischer Scientific), and 99% D-methanol (Aldrich). The experiments were carried out at a temperature of 20 ( 1 °C, controlled by a Haake model WC10/C1 thermostating bath. To measure the current transients of methanol decomposition (chronoamperometry), the following program was implemented. First, three activating/cleaning potential steps between the onsets of hydrogen evolution and the oxide region were applied to the electrode. A 10-ms pre-step at the electrode potential (EP), 250 mV more positive than the lower potential bias, was next applied48-50 (see also initial part of potential waveform in the inset to Figure 7). The final potential step was to a measuring potential (EM) at which methanol decomposition was investigated. 3. Results and Discussion 3.1. Information from Independent Measurements by Fast Cyclic Voltammetry and Chronoamperometry. Limited

J. Phys. Chem. B, Vol. 104, No. 23, 2000 5567

Figure 2. Chronoamperometric curves for the Pt electrode in 0.6 M CH3OH + 0.1 M H2SO4 at selected potentials: (A) 0.47 and (B) 0.67 V. Solid line, background current; dashed line, total current.

mechanistic insight into the initial stage of the methanol decomposition reaction can be gained from the slow (traditional) cyclic voltammetry measurements (below 1 V s-1). Deeper understanding of the nature of the methanol oxidation process can be obtained when a sufficiently high sweep rate is used.30,48,54 For instance, at 35 V s-1, the voltammogram in a 0.6 M CH3OH solution in sulfuric acid (0.1 M) displays many features characteristic of the CO-free platinum surface (Figure 1A). Clearly, the modification to the hydrogen adsorption/ desorption is negligible, and the oxidation peak for surface CO (generated from methanol or formic acid decomposition) is absent (see also Figure 7 and ref 57). There exists, however, an oxidative current in the potential range from 0.4 to 0.8 V, peaking at around 0.6 V, obviously due to some form of a methanol decomposition process.32 Because no noticeable surface CO is formed,48 the decomposition products in this potential range are mainly soluble.30 As shown by the data in Figure 1A vs B, the voltammetric peak current depends on the isotope composition of methanol, CH3OH vs CD3OH, as the H-substituted methanol is 3 times more reactive than the D-methanol. This indicates that a C-H scission is involved in the rate-determining step, confirming our previous conclusions obtained from instantaneous current measurements under chronoamperometric conditions.49,50,57 Figure 2 shows typical chronoamperometric curves for methanol electro-oxidation in sulfuric acid solution, including blank transients taken in the clean supporting electrolyte. The curves are given for two electrode potentials, 0.47 V (Figure 2A) and 0.67 V (Figure 2B). The net transient current was obtained after the blank current was subtracted from the total oxidation current.49,50 The decay of the background-corrected current depends on the electrode potential, surface crystallography, and bulk concentration of methanol.49,50 At low potentials (E < 0.52 V), the current decays quickly to a very small value (essentially zero). At more positive potentials, a steady-state current was measured. The highest current obtained in this study is ca. 3 orders of magnitude lower than the diffusion-controlled current, as determined using the Cottrell equation.58 That is, the diffusion delivery of methanol for our measurements is practically infinite. In Figure 3, the logarithm

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Figure 3. Initial current density of methanol decomposition as a function of methanol bulk concentration in 0.1 M H2SO4 on the Pt electrode at five selected potentials (as indicated in the graph legend). Average slope is 0.49 ( 0.05 at 95% probability level.

Figure 4. QCA (total charge density, or charge density due to the sum of all methanol decomposition reactions, shaded area) determined from the integration of the background-corrected current densities recorded for the Pt electrode in a 0.6 M CH3OH + 0.1 M HClO4 solution at 0.52 V. The initial current density (j0) and the steady-state current density (js-s) are also indicated.

of the initial (instantaneous) background-corrected current density is plotted vs the logarithm of methanol bulk concentration. The dependence is linear for potentials varying from 0.32 to 0.52 V with a slope of ca. 0.5. Finally, the total charge (QCA) is obtained by integrating the background-corrected currents over the reaction time, as illustrated in Figure 4, and the plots of QCA as a function of reaction time for several electrode potentials are shown in Figure 5. The plot of QCA obtained for a time of 1.5 s as a function of the electrode potential is shown in Figure 6. 3.2. Combined Chronoamperometry with Fast Cyclic Voltammetry. As mentioned in the initial remarks of subsection 3.1, the separate voltammetric and chronoamperometric data do not reveal information about the surface/bulk product distribution along the progress of methanol decomposition, and the remedy is to combine chronoamperometry and fast cyclic voltammetry.57 To do so, the experiments were conducted from 10 ms to a few seconds, which essentially covers the full time range for the methanol decomposition reaction. After the completion of the current measurements at the reaction time t, cyclic voltammetric sweeps at 50 V s-1 were applied, starting from the measuring potential (EM for chronoamperometry) and going in the positive direction. The potential pulse-sweep

Lu et al.

Figure 5. Plot of QCA as a function of decomposition time (Pt electrode in 0.6 M CH3OH + 0.1 M HClO4) at five selected potentials: 0.32 (O), 0.37 (0), 0.42 (4), 0.47 (3) and 0.52 V (]).

Figure 6. Plot of QCA (measured at 1.5 s) as a function of the decomposition potential (Pt electrode in 0.6 M CH3OH + 0.1 M HClO4).

scheme is as represented in the inset to Figure 7, with the upper potential limit selected to be such a value that the complete oxidative removal of surface CO was achieved, while the negative potential bias nearly coincided with the hydrogen evolution edge. The outcome of the combined chronoamperometry (CA) and cyclic voltammetry (CV) experiment is demonstrated in Figure 7, where the vertical dotted lines result either from the double layer charging (the pre-step at 0.25 V) or from the currents due to methanol decomposition (the measurement at 0.32 V). The sweeps offer information about the extent of surface CO generation from methanol decomposition at EM, as shown by the CO oxidation peak (1st sweep) between 0.85 and 1.35 V, superimposed on the Pt oxidation current (following sweeps). A complete correction for platinum surface oxidation is assumed. The plot of QCV as a function of time is shown in Figure 8. Although, overall, the data demonstrate the expected trend, it is surprising to observe a kinetic delay (an induction period) at the beginning of the reaction, below 50 ms (see inset in Figure 8). This behavior demonstrates sluggishness in CO uptake from methanol decomposition on platinum at the beginning of the reaction (see below). Surface CO and CO2 dissolved in solution are the broadly recognized products of methanol decomposition on platinum.

Catalytic Methanol Decomposition Pathways

J. Phys. Chem. B, Vol. 104, No. 23, 2000 5569 CO (become part of reaction 1, especially at lower potentials) or CO2 (become part of reaction 2, especially at higher potentials) or diffuse to the bulk of solution (and escape chemisorption). Chances that such an escape will happen are much higher on smooth than on rough surfaces.32,39,59 We now introduce a time-dependent parameter that will characterize the relationship between the methanol decomposition and surface CO formation processes, which is defined as

r(t) )

Figure 7. Typical current vs potential plot obtained in a combined chronoamperometry/cyclic voltammetry experiment (the potential pulsesweep scheme is shown in the inset). 1 refers to the last cleaning/ activating step (10 ms at 50 mV); 2 to the prestep (10 ms at 250 mV); 3 to the methanol decomposition step (adjustable time at adjustable potential, here 900 ms at 320 mV); 4 to the methanol oxidation peak in the double layer that does not produce surface CO; 5 to the CO stripping charge (QCV, in this case is 66 µC cm-2). The Pt electrode was in a 0.6 M CH3OH + 0.1 M H2SO4 solution. The sweep rate in the cyclic voltammetric measurements was 50 V s-1.

Figure 8. Plot of QCV (CO stripping charge obtained from fast cyclic voltammetry) as a function of time for the Pt electrode at 0.37 V in 0.6 M CH3OH +0.1 M H2SO4. Inset shows the initial stage of CO formation.

However, other products, such as formic acid and/or formaldehyde, have been found30,35,36,39 and must be considered here. Therefore, the interfacial reactions to cover all realistic possibilities are as follows:

CH3OH f CO + 4H+ + 4e-

(1)

CH3OH + H2O f CO2 + 6H+ + 6e-

(2)

CH3OH + H2O f HCOOH + 4H+ + 4e-

(3)

CH3OH f HCHO + 2H+ + 2e-

(4)

CO + H2O f CO2 + 2H+ + 2e-

(5)

As written, reactions 2-4 occur without adsorbed product formation. However, the HCOOH and HCHO formed at the platinum/solution interface can either decompose to form surface

QCA(t) QCV(t)

(6)

where QCA is the total decomposition charge in chronoamperometry and QCV is the CO oxidation charge in fast cyclic voltammetry (reaction 5) (see also subsection 3.1). In essence, this is an operation of normalizing the overall methanol decomposition reaction to the extent of CO poisoning. If the charge delivery in a chronoamperometric measurement is due to surface CO formation only (a complete methanol dehydrogenation, reaction 1), the r(t) value should be equal to 2 and should be time-independent. Statistically significant r values greater than 2 are indicative of soluble product formation in parallel with surface CO formation. Obviously, if a constant current is recorded, the reaction is at steady state, and the r(t) value is time-dependent and increases faster or slower depending on the CO coverage. In other words, the r(t) value offers a unique opportunity for analyzing the applicability of the dualpath mechanism in methanol decomposition on platinum. Because the critical relationship interrogated here concerns the effects of the electrode potential and the reaction time simultaneously, three-dimensional plots of the r values were constructed for a fixed bulk concentration of methanol in perchloric and sulfuric acid solutions, respectively, as shown in Figures 9A and B. Apparently, r values greater than 2 are observed in the full potential range investigated. That is, the values much greater than 2 are found not only at high potentials (E > 0.52 V), but also at low potentials. It is obvious that, at high potentials, soluble products (mainly CO2) are formed in high yields, while CO accumulation is small or even negligible, as it rearranges to CO2 quickly. This yields the large r(t) ratios. However, because CO is not oxidized at low potentials (E < 0.52 V), the observation of the high r(t) ratios at short times leads to the conclusion that soluble products also are formed at such low potentials, in parallel to surface CO formation (dual path). It should be noted that QCV charges obtained at short times are very small, and as a consequence, the experimental uncertainty in determining the r(t) ) QCA/QCV ratio is quite high, especially at the shortest times. This is demonstrated in Figure 10, in which the data scattering is shown (five independent measurements were made). Although the scatter is large, the reproducibility is good enough to conclude that the key effect observed in this study, namely, the high r values at short times and low potentials, is statistically significant. The effect is not related to uncompensated resistance because, as we found in numerous control experiments, the results were not affected by the electrode surface area, i.e., by the overall magnitude of the oxidation current. The three-dimensional curves obtained in sulfuric acid and in perchloric acid (Figure 9A and B) show a strong anion effect in the methanol decomposition reaction. Clearly, the electrode remains in an active state longer in perchloric acid than in sulfuric acid media. This observation is in accordance with previous results50 showing a reduction in the magnitude of the

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Lu et al.

Figure 11. Tafel plots of the steady-state currents for the Pt electrode in 0.1 M HClO4 (O, solid trace) and in 0.1 M H2SO4 (0, dashed trace) containing 0.6 M CH3OH. Open symbols represent points deviating from linearity and not taken into account in regression calculations. Slopes are 130 and 140 mV dec-1, and regression coefficients are 0.997 and 0.993 for HClO4 and H2SO4, respectively.

Figure 9. Three-dimensional mesh plots of the QCA/QCV ratio as a function of potential and time for the Pt electrode in (A) 0.1 M HClO4 and (B) 0.1 M H2SO4 containing 0.6 M CH3OH (see text).

Figure 10. Plot of the QCA/QCV ratio as a function of time for the Pt electrode at 0.42 V in 0.6 M CH3OH + 0.1 M H2SO4 (confidence intervals were estimated from five independent measurements at the 95% confidence level).

instantaneous methanol oxidation current upon a change in supporting electrolyte from perchloric acid to sulfuric acid. The high r(t) value also demonstrates that the anion adsorption is more inhibitive toward decomposition of methanol to CO2 than to CO. 3.3. Origin of the High QCA/QCV Ratio at Short Times. The high r(t) ) QCA(t)/QCV(t) ratio at short times (Figures 9 and 10) is due to the relative sluggishness of methanol decomposition to surface CO on platinum at the beginning of the reaction, while the soluble product formation proceeds. However, the type of the decay with the high r(t) ratio shown in Figure 2 requires an assumption that, at short times, the

soluble product formation channel is suppressed by surface processes other than CO poisoning, which otherwise accounts for a majority of the overall decay data. An explanation of these findings is already available in the literature. Specifically, Burke55 has suggested that a reactive form of surface oxide (Otype reagents) exists on platinum at very low potentials and is responsible for the initiation of oxidation of a variety of organic molecules on platinum electrodes. Our data may tentatively be envisioned in favor of this concept under the proviso that such an O-type reagent can rapidly (between 10 and 40 ms) be scavenged from the surface by reacting methanol and cannot be replenished to sustain a soluble product formation process. Notably, at E g 0.52 V, the current transient does not decay to zero but, instead, yields a steady-state current, is-s. A typical curve recorded for E ) 0.52 V is shown in Figure 4 (right part). The js-s vs E plots are clearly Tafelian (Figure 11), with the nearly 120 mV/dec slope independent of both the type of anion and the methanol bulk concentration. 3.4. Model of the Decomposition of Methanol on Polycrystalline Platinum. The model that we propose below is an extension of our relatively simple, two-term model formulated previously on the basis of the chronoamperometric data only.44,49,50 The new model, in view of the present data obtained by the combination of chronoamperometry and fast cyclic voltammetry, must accommodate more than two terms. Apparently, to fit all of the data, both CA and CV, yet another reaction channel must be considered, in addition to the previous two branches. Specifically, one must consider the channel leading to soluble species formation within the first 40 ms of the experiment (iex) to account for the rapidly decaying current component. It is assumed here that this process can be approximated by an exponential current decay.

iex ) (iex)0 e-bt

(7)

where (iex)0 is the instantaneous current of the exponential decay process and b is the rate constant of this process. With these assumptions in mind, the current model was constructed using the sequence of rate equations shown below.

Catalytic Methanol Decomposition Pathways

J. Phys. Chem. B, Vol. 104, No. 23, 2000 5571

iΣ ) inet + ibackg

(8)

inet ) iox + iads + iex

(9)

iox ) iox,0 (1 - mθ)2

(10)

iads ) Qmax/θmax dθ/dt ) iads,0(1 + θmaxkadt)-2

(11)

inet,0 ) iox,0 + iads,0 + iex,0

(12)

Qmax ) NPtθmax4e

(13)

where iΣ is the total current; inet, the net current (resulting from the correction of iΣ for the background current, ibackg); iox, the methanol oxidation current leading to soluble product formation; iads, the methanol dehydrogenation current leading to CO formation; inet,0, iox,0, and iads,0, the corresponding instantaneous currents; θmax, the maximum coverage of CO; Qmax, the electric charge needed to establish θmax; NPt, the total number of surface Pt sites per 1 cm2; m, the number of surface sites occupied by one CO molecule; 4, the number of electrons produced in the CO formation process (reaction 1); and e, the elementary charge. For m ) 2,49,50 the current-time profile has the following form:

(

i ) iex,0 e-bt + iox,0 1 -

) (

Figure 12. Decays of the total net current density (solid trace) and its three components (according to eq 14) for the Pt electrode at 0.52 V in 0.6 M CH3OH + 0.1 M HClO4: the exponential decay component (dashed trace, first term in equation), the soluble product formation component (dot-dot-dashed trace, second term), and the methanolic CO formation component (dotted trace, third term).

)

2 θ2maxkadt 1 + iads,0 1 + θmaxkadt 1 + θmaxkadt (14)

amperometry and QCV is the CO oxidation charge from fast cyclic voltammetry, and shown how this ratio depends on the electrode potential, the methanol-platinum interaction time, and the composition of supporting electrolyte. From our analysis, we conclude that the methanol decomposition reaction that occurs at short times (below 40 ms) does not lead to an effective surface CO formation process, even at relatively low potentials. In contrast, the formation of soluble products proceeds actively. The measurements provided us with a basis for confirming, but also reformulating, the dual-path mechanism for methanol oxidative decomposition on platinum. The reformulation consists of the notion that formic acid and/or formaldehyde are formed in addition to the products that previously were believed to be associated with methanol decomposition on platinum, namely, surface CO and CO2. Apparently, the formation of formic acid and/or formaldehyde accounts for the excessive charge beyond that of simple methanol dehydrogenation to surface CO, especially below the potential for the CO2 formation threshold, 0.35 V vs RHE. From our modeling study, we conclude that three channels are needed to account for the net oxidation current as it decays in time, two leading to soluble species formation and one to CO formation. Possible pathways leading to the rapidly decaying current component are discussed. Still, the drop in methanol oxidation current under potential step conditions is mainly accounted for by the formation of chemisorbed CO, except for the initial stage within the first 40 ms of reaction.

This model gives a satisfying fit for the experimental chronoamperometric data in the potential range 0.32-0.52 V. An example of the fit is shown in Figure 12. The fitting results and parameters are given in Table 1, and some of them are compared with the experimental results, showing very good agreement. To obtain the experimental jads, the QCV vs t data (in Figure 8) were treated by a cubic spline procedure, and the first derivative dQ/dt was found, which served as a pseudocurrent for CO formation. Notably, there was an increase in the derivative at the very beginning, because of the “sigmoid” feature in the QCV vs t profile. In Table 1, the values of the parameters depend on the electrode potential, with the notable exception of θmax, which is potential-independent in the studied range. Recently, Sriramulu et al.46,47 proposed a model based on Langmuir-Hinshelwood kinetics for methanol oxidation on a Pt(111) electrode, using chronoamperometric and linear sweep voltammetric measurements. We believe that the limitation of this model, as of our previous models, is in the assumption that the net oxidation current is constructed from the contributions of the CO and CO2 formation processes (and, eventually, of the CO oxidation process). However, our data no longer support such an assumption; rather, a third current component, iex ) (iex)0 e-bt, accounting for the short time events is needed. Our present model is free of this previous limitation.

Acknowledgment. This work was supported by the National Science Foundation under Grant CHE 97-000963 and by the Department of Energy under Grant DEFG02-96ER45439, administered by the Frederick Seitz Materials Research Laboratory at the University of Illinois.

4. Conclusions We have defined the ratio r(t) ) QCA(t)/QCV(t), where QCA is the methanol decomposition charge obtained from chrono-

TABLE 1: Selected Experimental and Fittinga Results for the Chronoamperometric Data in 0.6 M CH3OH + 0.1 M H2SO4 E vs RHE V

jnet,0 (exp)b mA cm-2

jnet,0 (fit) mA cm-2

jads,0 (exp)c mA cm-2

jads,0 (fit) mA cm-2

jex,0 mA cm-2

jox,0 mA cm-2

b s-1

kad s-1

θmax ML

0.32 0.37 0.42 0.47 0.52

0.93 1.61 3.00 4.97 7.06

0.92 1.64 3.09 5.11 7.25

0.56 0.57 0.98 2.45 3.46

0.32 0.62 1.07 2.54 3.54

0.45 0.85 1.72 1.86 2.55

0.15 0.17 0.30 0.71 1.16

8.0 12.0 21.9 45.1 65.2

2.1 2.9 4.0 11.9 14.9

0.36 0.40 0.43 0.40 0.42

a Using the three-term eq 14. b Read by extrapolation of the initial stage of the chronoamperometric curve to t ) 0 (see also Figure 4). c Read by extrapolation of the result of differentiation of an experimental QCV vs t plot to t ) 0 (see also Figure 8).

5572 J. Phys. Chem. B, Vol. 104, No. 23, 2000 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) Adzic, R. R. In Modern Aspects of Electrochemistry; Bockris, J. O’M., Conway, B. E., White, R. E., Eds.; Plenum Press: New York, 1990; Vol. 21, p 163. (4) Horanyi, G.; Wieckowski, A. Proc.-Electrochem. Soc. 1992, 92, 70. (5) Jarvi, T. D.; Stuve, E. M. In Electrocatalysis; Lipkowski, J., Ross, P. N., Eds.; Wiley-VCH: New York, 1998; p 75. (6) Hamnett, A. In Interfacial Electrochemistry; Wieckowski, A., Ed.; Marcel Dekker: New York, 1999; p 843. (7) Chrzanowski, W.; Wieckowski, A. In Interfacial Electrochemistry; Wieckowski, A., Ed.; Marcel Dekker: New York, 1999; p 937. (8) Wasmus, S.; Kuver, A. J. Electroanal. Chem. 1999, 461, 14. (9) Watanabe, M.; Motoo, S. J. Electroanal. Chem. 1975, 60, 267. (10) Watanabe, M.; Uchida, M.; Motoo, S. J. Electroanal. Chem. 1987, 229, 395. (11) Hable, C. T.; Wrighton, M. S. Langmuir 1993, 9, 3284. (12) Frelink, T.; Visscher, W.; van Veen, J. A. R. Surf. Sci. 1995, 335, 353. (13) Frelink, T.; Visscher, W.; van Veen, J. A. R. Langmuir 1996, 12, 3702. (14) Gasteiger, H. A.; Markovic, N.; Ross, P. N., Jr.; Cairns, E. J. J. Phys. Chem. 1993, 97, 12020. (15) Markovic, N. M.; Gasteiger, H. A.; Ross, P. N.; Jiang, X. D.; Villegas, I.; Weaver, M. J. Electrochim. Acta 1995, 40, 91. (16) Wang, K.; Gasteiger, H. A.; Markovic, N. M.; Ross, P. N., Jr. Electrochim. Acta 1996, 41, 2587. (17) Kabbabi, A.; Faure, R.; Durand, R.; Beden, B.; Hahn, F.; Leger, J.-M.; Lamy, C. J. Electroanal. Chem. 1998, 444, 41. (18) Ley, K. L.; Liu, R.; Pu, C.; Fan, Q.; Leyarovska, N.; Segre, C.; Smotkin, E. S. J. Electrochem. Soc. 1997, 144, 1543. (19) Anderson, A. B.; Grantscharova, E.; Seong, S. J. Electrochem. Soc. 1996, 143, 2075. (20) Gurau, B.; Viswanathan, R.; Liu, R. X.; Lafrenz, T. J.; Ley, K. L.; Smotkin, E. S.; Reddington, E.; Sapienza, A.; Chan, B. C.; Mallouk, T. E.; Sarangapani, S. J. Phys. Chem. B 1998, 102, 9997. (21) Chrzanowski, W.; Kim, H.; Wieckowski, A. Catal. Lett. 1998, 50, 69. (22) Tremiliosi-Filho, G.; Kim, H.; Chrzanowski, W.; Wieckowski, A.; Grzybowska, B.; Kulesza, P. J. Electroanal. Chem. 1999, 467, 143. (23) Acres, G. J. K.; Hards, G. A. Philos. Trans. R. Soc. London Ser. A 1996, 354, 1671. (24) Hogarth, M. P.; Hards, G. A. Platinum Met. ReV. 1996, 40, 150. (25) Acres, G. J. K.; Frost, J. C.; Hards, G. A.; Potter, R. J.; Ralph, T. R.; Thompsett, D.; Burstein, G. T.; Hutchings, G. J. Catal. Today 1997, 38, 393. (26) Schmidt, V. M.; Bro¨ckerhoff, P.; Ho¨hlein, B.; Menzer, R.; Stimming, U. J. Power Sources 1994, 49, 299. (27) Hards, G. A. Int. J. Hydrogen Energy 1996, 21, 777. (28) Reddington, E.; Sapienza, A.; Gurau, B.; Viswanathan, R.; Sarangapani, S.; Smotkin, E. S.; Mallouk, T. E. Science 1998, 280, 1735. (29) Patil, P. G. J. Power Sources 1992, 37, 171. (30) Ota, K.-I.; Nakagawa, Y.; Takahashi, M. J. Electroanal. Chem. 1984, 179, 179.

Lu et al. (31) Papoutsis, A.; Leger, J.-M.; Lamy, C. J. Electroanal. Chem. 1987, 234, 315. (32) Lopes, M. I. S.; Beden, B.; Hahn, F.; Leger, J.-M.; Lamy, C. J. Electroanal. Chem. 1991, 313, 323. (33) Laborde, H.; Leger, J.-M.; Lamy, C.; Garnier, F.; Yassar, A. J. Appl. Electrochem. 1990, 20, 524. (34) Peremans, A.; Tajeddine, A. Chem. Phys. Lett. 1994, 220, 481. (35) Korzeniewski, C.; Childers, C. L. J. Phys. Chem. B 1998, 102, 489. (36) Childers, C. L.; Huang, H.; Korzeniewski, C. Langmuir 1999, 15, 786. (37) Xia, X. H.; Iwasita, T.; Ge, F.; Vielstich, W. Electrochim. Acta 1996, 41, 711. (38) Iwasita, T.; Nart, F. C.; Lopez, B.; Vielstich, W. Electrochim. Acta 1992, 37, 2361. (39) Lin, W.-F.; Wang J.-T.; Savinell, R. F. J. Electrochem. Soc. 1997, 144, 1917. (40) Christensen, P. A.; Hamnett, A.; Munk, J.; Troughton, G. L. J. Electroanal. Chem. 1994, 370, 251. (41) Kunimatsu, K. Ber. Bunsen-Ges. 1990, 94, 1025. (42) Breiter, M. W. J. Electroanal. Chem. 1967, 14, 407. (43) Breiter, M. W. In Modern Aspects of Electrochemistry; Bockris, J. O’M., Conway, B. E., Eds.; Plenum Press: New York, 1975; Vol. 10, p 161. (44) Herrero, E.; Chrzanowski, W.; Wieckowski, A. J. Phys. Chem. 1995, 99, 10423. (45) Jarvi, T. D.; Sriramulu, S.; Stuve, E. M. J. Phys. Chem. B 1997, 101, 3649. (46) Sriramulu, S.; Jarvi, T. D.; Stuve, E. M. Electrochim. Acta 1998, 44, 1127. (47) Sriramulu, S.; Jarvi, T. D.; Stuve, E. M. In Interfacial Electrochemistry; Wieckowski, A., Ed.; Marcel Dekker: New York, 1999; p 793. (48) Wieckowski, A.; Chrzanowski, W.; Herrero, E. In New Materials for Fuel Cell Systems I; Savadogo, O., Roberge, P. R., Veziroglu, T. N., Eds.; EÄ ditions de l’EÄ cole Polytechnique de Montre´al: Montreal, Canada, 1995; p 326. (49) Franaszczuk, K.; Herrero, E.; Zelenay, P.; Wieckowski, A.; Wang, J.; Masel, R. I. J. Phys. Chem. 1992, 96, 8509. (50) Herrero, E.; Franaszczuk, K.; Wieckowski, A. J. Phys. Chem. 1994, 98, 5074. (51) Ren, B.; Li, X. Q.; She, C. X.; Tian, Z. Q. Electrochim. Acta, in press. (52) Adzic, R.; Tripkovic, A. V.; O’Grady, E. Nature 1982, 296, 137. (53) Markovic, N.; Ross, P. N. J. Electroanal. Chem. 1992, 330, 499. (54) Herrero, E.; Franaszczuk, K.; Wieckowski, A. J. Electroanal. Chem. 1994, 361, 269. (55) Burke, L. D.; Casey, J. K. Electrochim. Acta 1992, 37, 1817. (56) Conway, B. E.; Angerstein-Kozlowska, H. Acc. Chem. Res. 1981, 14, 49. (57) Lu, G.-Q.; Crown, A.; Wieckowski, A. J. Phys. Chem. B 1999, 103, 9700. (58) Bard, A. J.; Faulkner, L. R. Electrochemical Methods; John Wiley & Sons: New York, 1980. (59) One of us (A.W.) appreciates fruitful discussions of this issue with H. Baltruschat from the University of Bonn.