2322
J. Phys. Chem. 1996, 100, 2322-2329
Mechanism of Carbon Monoxide Electrooxidation on Monocrystalline Gold Surfaces: Identification of a Hydroxycarbonyl Intermediate Gregory J. Edens, Antoinette Hamelin,† and Michael J. Weaver* Department of Chemistry, Purdue UniVersity, West Lafayette, Indiana 47907-1393 ReceiVed: August 30, 1995X
Kinetic data are presented for the electrooxidation of aqueous solution carbon monoxide to carbon dioxide on two monocrystalline gold surfaces, Au(210) and (110), with the objective of elucidating the reaction mechanism, especially regarding the nature of adsorbed intermediate(s). Tafel plots (i.e., log rate versus electrode potential) were obtained by means of linear sweep voltammetry, particularly as a function of the solution reactant concentration and over a wide range (0-13.5) of the electrolyte pH. Under most conditions, the reaction order in CO was found to be near unity, as anticipated from the low coverages of adsorbed CO ascertained from infrared spectroscopy. Interestingly, the log rate-pH dependence observed on both surfaces display three distinct regions. At low (e2) and higher (g4) pH values, essentially unit slopes were obtained (i.e., a unity reaction order in [OH-]), these regions being separated by one displaying apparently pH-independent kinetics. The potential region over which conveniently measurable electrooxidation kinetics occur lies substantially (ca. 0.8 V) below the onset of gold surface oxidation throughout the entire pH range. The pH-dependent kinetic behavior is consistent with a reaction pathway featuring the involvement of an adsorbed hydroxycarbonyl intermediate. While such intermediates have been identified in a number of metal complex-catalyzed CO oxidations in homogeneous solution, they apparently have not been considered previously for such electrocatalytic processes. The observed unity hydroxide reaction order at higher pH values is indicative of a rate-determining step (rds) involving OH- discharge onto adsorbed CO sites to form the hydroxycarbonyl species, while the apparent transition to zero-order kinetics at lower pH is consistent with water rather than OH- becoming the preferred reactant. This picture is supported by solvent isotope measurements which display the onset of a substantial H/D isotope effect below pH 4, signaling the occurrence of proton transfer within the rds. The emergence of another pH-dependent reaction pathway at the lowest pH values is attributed to a rds involving hydroxycarbonyl decomposition to form CO2. The mechanistic opportunities provided by the analysis of electrocatalytic rate-potential data over wide pH ranges are pointed out, along with the possibility that the proposed hydroxycarbonyl pathway occurs for a wide range of related processes on transition-metal surfaces.
Introduction The catalytic oxidation of carbon monoxide to form carbon dioxide, in both heterogeneous (electrochemical, metal-gas) and homogeneous solution-phase environments, is a process having enormous technological as well as fundamental importance. Of particular interest in the latter two environments is the so-called “water-gas shift” reaction (WGSR), whereby CO and water react to yield CO2 and hydrogen.1,2 Since electrochemical systems provide for separate oxidation and reduction half-reactions, the significant analogous process in aqueous media involves the electrooxidative formation of CO2 and hydronium ions. This reaction is of broad-based importance in its own right, since adsorbed CO is a common reaction intermediate and/or poison in the electrooxidation of methanol and other organic feedstocks of significance in fuel-cell technology.3 There have been numerous studies aimed at designing homogeneous-phase WGSR catalysts.1 These consist of transition-metal carbonyl complexes operative in aqueous alkaline and (to a lesser extent) in acidic media. While an overall emphasis has been placed on synthetic aspects rather than detailed mechanistic studies, an instructive facet of these WGSR catalysts has been the identification and even isolation in some cases of so-called hydroxycarbonyl intermediates formed by † Permanent address: Laboratoire d’Electrochimie, des Interfaces du CNRS, 1, Place A. Briand, 92195 Meudon, France. X Abstract published in AdVance ACS Abstracts, January 1, 1996.
0022-3654/96/20100-2322$12.00/0
reacting the carbonyl ligand with hydroxyl.1a,b In contrast, mechanistic information for the corresponding electrooxidation processes of carbon monoxide at metal-aqueous interfaces is sketchy at best, particularly with regard to the nature of the adsorbed intermediates.3,4 The latter situation results not only from the relative dearth of careful kinetic studies on structurally defined monocrystalline surfaces5 but also from the strong CO chemisorption commonly encountered on transition metals. This leads typically to high and even saturated CO coverages under most catalytic operating conditions. Consequently, adsorbed CO electrooxidation on such catalysts tends to proceed via “nucleation-growth” processes where the reaction with the hydroxyl (or water) coreactant takes place chiefly at island “holes” created within the initially close-packed CO adlayer.5,6 Such kinetic circumstances, while certainly of catalytic interest, largely preclude the deduction of reaction mechanisms from observed rate laws since the pHdependent CO electrooxidation rates tend to be dominated by the availability (or lack thereof) of surface sites. A markedly different situation, however, commonly applies to CO electrooxidation on monocrystalline gold surfaces.7 While the CO oxidation kinetics are remarkably facile on gold even in acidic aqueous media, the reaction proceeding at potentials substantially (0.5-1 V) below the initiation of gold surface oxidation, CO is only weakly adsorbed (fractional coverage θCO e 0.1).7 The overall reaction is therefore not affected significantly by the availability of free surface sites. © 1996 American Chemical Society
CO Electrooxidation on Au Surfaces Indeed, the reaction order in solution CO is found to be essentially unity (vide infra); moreover, the reaction kinetics are extremely sensitive to the surface crystallographic orientation, the CO electrooxidation rates at a given potential and pH differing by as much as 100-fold between different gold surface planes.7a These circumstances as deduced from earlier electrochemical kinetic-infrared spectral studies in this laboratory have led us to undertake a further, more quantitative, examination of the CO electrooxidation kinetics on selected gold surfaces, with the objective of deducing details of the reaction pathway in aqueous media. We have chosen to examine in greatest detail the electrooxidation of CO on Au(210) and (110), since these surfaces yield the quantitatively reproducible kinetic responses necessary for the reliable deduction of rate laws. The results of this study are described herein. Of particular significance is the pH-dependent kinetic behavior, examined systematically throughout the range from strongly acidic to alkaline conditions. Three distinct pH-dependent regimes are evident, which are consistent with the occurrence of different rate-determining steps, yet all with the apparent involvement of a hydroxycarbonyl reaction intermediate. Further information on the nature of the rate-determining steps has been obtained from deuterium solvent isotope measurements. Besides yielding mechanistic insight likely to be of more general significance for CO oxidation catalysis, the findings highlight the previously untouted virtues of electrode kinetics for unraveling rate laws for aqueous-based oxygen-transfer catalytic processes over much wider pH ranges than are accessible for analogous processes in homogeneous solution. Experimental Section The single-crystal electrodes Au(111), (110), (100), (533), and (210) examined here were prepared in LEI-CNRS.8,9 The condition of each face was checked in 10 mM perchloric acid solutions with cyclic voltammetry: the current profiles for the formation of a monolayer of oxide and the potential of the capacitive current minimum in the double-layer region were in good agreement with literature voltammograms. The electrodes were annealed in a hydrogen-air flame to red heat and cooled for 1 min in ambient air and finally in ultrapure water before being transferred to the electrochemical cell. A two-compartment cell was used to isolate the test solution from the saturated sodium calomel reference electrode (SCE). The hanging meniscus method9 was used in some cases, allowing only the well-defined crystal face to contact the solution. For experiments in which immersion of the electrode was necessary, however (vide infra), the crystal walls were wrapped in Teflon tape to insulate them from solution. A gold wire served as the counter electrode. Most experiments utilized 99.3% CO from Airco. However, essentially identical electrode kinetic results were obtained when using a higher purity (99.99%) grade stored in an aluminum cylinder to prevent iron carbonyl formation. Supporting electrolytes were prepared from HClO4, NaClO4, LiClO4 (GFS Chemicals), or NaOH (Alfa Ultrapure). Perchlorate salts were twice recrystallized from water. Buffers were prepared from Na2BO4, Na2HPO4, KH2PO4, and NaC2H3O2 (all ACS Reagents) according to literature methods and the ionic strength brought to 0.1 by addition of NaClO4 (vide infra). Water was purified by means of a Milli-Q Plus system (Millipore). Deuterium oxide (Cambridge Isotopes) was used as received. The purity was deemed to be sufficient by comparing the potentialdependent double-layer capacitance of Au(210) in D2O media with that in corresponding aqueous solutions.
J. Phys. Chem., Vol. 100, No. 6, 1996 2323 Electrode kinetic measurements were made by cyclic voltammetry as described below. Measurements employed a PAR 173/175 potentiostat and a Soltec VP 6424S XYY recorder. The differential capacitance measurements utilized additionally a HP 3314A function generator and a PAR 5204 lock-in amplifier. Activation parameters were measured under nonisothermal cell conditions10,11 by the additional use of a jacketed cell and a circulating water thermostat (Braun Melsungen). Unless otherwise noted, measurements were made at 23 ( 1 °C. Electrode potentials are reported versus the SCE. Results and Data Analysis The basic tactic employed here involves the extraction of electrooxidation rates for carbon monoxide solute as a function of the electrode potential, E, with particular interest in the dependence of these rate-potential data on the solution pH and also on solvent deuteration and temperature. Linear sweep voltammetry (LSV) was employed here to obtain the electrode kinetics data for several reasons. While potential-step techniques offer some advantages for quantitative kinetic analysis, LSV has the major virtue of enabling the desired rates as a function of E to be obtained from a single voltammetric measurement. The ability to ascertain the data “quality” by inspecting the current-potential (i-E) morphology is also highly advantageous when studying electrocatalytic reactions on solid surfaces, as considered here. These reasons led us to select LSV in our earlier study of CO electrooxidation kinetics on monocrystalline gold surfaces.7a However, the more quantitative nature of the present study led us to undertake a slightly different data analysis procedure from that in ref 7a, which is outlined below along with a description of the kinetic results. Representative LSV data obtained in this work are shown in Figure 1A-D. Parts A and B are LSV’s obtained at 0.1 V s-1 for Au(210) and (110), respectively, in CO-saturated aqueous 0.1 M HClO4 (solid traces) and in the supporting electrolyte alone (dashed traces). While the influence of gold oxide formation, seen by comparison with the dashed traces, is clearly evident at high positive electrode potentials (g0.9 V vs SCE), the kinetically determinable region for CO electrooxidation, leading up to the first anodic peak (solid trace), occurs within the so-called “double-layer” region at markedly lower potentials. A similar situation is also encountered in alkaline media, as exemplified in Figure 1, C and D, which refer to a CO-saturated solution in 0.09 M NaClO4 + 0.01 M Na2BO4, again on Au(210) and (110), respectively. The occurrence of an essentially complete two-electron electrooxidation to form CO2, as expected, was confirmed in these alkaline as well as acidic media by means of thin-layer infrared spectral sequences, essentially as outlined in ref 7. Specifically, exhaustive thin-layer electrooxidation of the solution CO yielded quantitatively the same amount of CO2 in both media, as deduced from the intensity of the characteristic asymmetric O-C-O stretch at 2343 cm-1. There was no evidence on this basis for significant production of formate anions in the alkaline media. However, the CO2 was seen by infrared spectroscopy to form carbonate anions by reaction with hydroxide at pH values above 10, as discerned from the appearance of a characteristically broad CO32- feature at about 1300 cm-1. In each case, the most reproducible voltammetric behavior was obtained by rapidly immersing the electrode, wrapped with Teflon to expose only the working face, into the desired electrolyte. It was found necessary to minimize contact with any gas-phase CO present above the cell solution since, surprisingly, extensive and rapid CO adsorption is seen to occur upon exposure of gold surfaces to wet gas-phase CO (vide
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Figure 2. Anodic voltammograms obtained at 0.1 V s-1 in COsaturated 0.1 M HClO4 on unreconstructed (dashed trace) and largely reconstructed Au(100) (solid trace). See text for details.
Figure 1. Anodic voltammograms obtained at 0.1 V s-1 (solid traces) for the electrooxidation of saturated aqueous solution in CO in (A, B) 0.1 M HClO4 and (C, D) 0.09 M NaClO4 + 0.01 M Na2BO4 on (A, C) Au(210) and (B, D) Au(110). Dashed traces are corresponding anodiccathodic cyclic voltammograms obtained in supporting electrolyte alone.
infra).12 (This difficulty typically precluded using the otherwisepreferred “hanging meniscus” method of contacting the electrode face with the electrolyte.9) Removal of any such irreversibly adsorbed CO could be achieved by an anodic cycle prior to recorded the LSV data. As already mentioned, Au(210) and (110) were selected for detailed study here in view of the high degree of reproducibility of the kinetic data on these surfaces, desired for the present pH-dependent and other analyses. Thus, the i-E traces were typically independent of the initial potential, provided this was within the nonfaradaic double-layer region, and hold time (often within the width of the recorder pen trace!). Holding the potential at substantial negative potentials, especially in alkaline electrolytes, however, yielded a slow buildup of a chemisorbed film, which can inhibit substantially the electrooxidation of solution CO. This film has been identified as a “nonreactive” form of chemisorbed CO by infrared spectroscopy.7b,12 Nevertheless, this complication could usually by avoided entirely by choosing less negative potentials or shorter hold times. The highly reproducible kinetic behavior recalls the complete absence of hysteresis for capacitance-potential data on Au(210) obtained with either voltammetric or ac impedance methods,13 consistent with the lack of potential-induced reconstruction (and even long-range atomic order) as observed by in-situ atomic-resolution scanning tunneling microscopy (STM).14 Admittedly, a “missing-row” (1 × 2) reconstruction is observed by STM to occur on Au(110) at potentials below ca. 0 V vs
SCE in acidic media.14,15 However, the (1 × 2) / (1 × 1) surface phase interconversion is seen to be facile on Au(110) even in “nonadsorbing” electrolytes such as perchloric acid.14,15 In contrast, the other low-index faces Au(111) and (100) display significantly less reproducible kinetics for CO electrooxidation, which are attributable partly to sluggish potentialdependent reconstruction. This is exemplified for Au(100) in the anodic voltammograms shown in Figure 2, again obtained at 0.1 V s-1 in CO-saturated 0.1 M HClO4. The dashed trace refers to a surface prepared in such a manner to yield largely unreconstructed (1 × 1) planes;16 this was engendered by sweeping the potential to high values, g0.5 V, and subsequently holding above ca. 0 V. The solid trace, on the other hand, was obtained after holding the potential at -0.3 V for 10 min, a condition known to generate extensive so-called “hexagonal” [or more properly (5 × 27)] reconstructed domains.16 The latter condition is seen to yield significantly slower CO electrooxidation kinetics, as evident in the ca. 0.05 V positive shift in the voltammogram. (Note that the voltammetric “postpeak” seen during the solid trace at ca. 0.55 V signals the potential-induced removal of the reconstruction.16) Experiments involving the use of thermally reconstructed Au(100) surfaces, prepared by cooling entirely in inert gas following flame annealing,16a were only partly successful, since it was difficult to avoid entirely the occurrence of gas-phase CO chemisorption during electrode transfer. We now consider the extraction of kinetic data from voltammograms such as in Figure 1A-D. The procedure used here is a modification of that used in our earlier studies.7a The critical step involves correcting the measured current densities for CO electrooxidation, i, for the occurrence of diffusion polarization, whereby the solution reactant concentration immediately adjacent to the interfacial region, C0, falls below the bulk concentration, Cb. The desired “diffusion-corrected” current density, icorr, can then be obtained with the knowledge of these concentrations simply from
icorr ) i(Cb/C0)
(1)
The current densities read from the experimental voltammograms were also corrected for residual nonfaradaic currents in the usual way, from current-potential data recorded in the absence of the reactant. For voltammograms displaying a welldefined “diffusion-controlled” current peak (such as those in Figure 1A-D), the required Cb/C0 ratios were obtained by means of the so-called “dimensionless current function”, employing explicit relations derived from those summarized in ref 17 to relate the current density at any potential on the rising
CO Electrooxidation on Au Surfaces
J. Phys. Chem., Vol. 100, No. 6, 1996 2325 TABLE 1: Kinetic Parameters for the Electrooxidation of Carbon Monoxide on Selected Monocrystalline Gold Surfaces in 0.1 M HClO4 surface
Rapp
103kEb/cm s-1
(110) (100) (1 × 1) (100) hex (111)
0.73 0.65 0.55 0.53
7.0 0.85 ∼0.4 0.15
surface
Rapp
103kEb/cm s-1
(210) (533) (221)
0.67 0.45 0.45
4.5 1.3 1.6
a Apparent anodic transfer coefficient, determined from Tafel plots as outlined in the text. b Rate constant for CO electrooxidation measured at (or extrapolated to) 0.3 V vs SCE, determined by means of eq 3.
Figure 3. Logarithmic plots of current density (corrected for diffusion polarization, see text) versus the electrode potential for electrooxidation on Au(110) of various concentrations of solution CO in 0.1 M HClO4.
part of the voltammogram, iE (ratioed to the peak current density, iE/ip), to the corresponding Cb/C0 values. Full details of this procedure are given in ref 12. The above analysis, however, presumes implicitly that the reaction order in solution CO is unity, i.e., that the anodic current at a given electrode potential is proportional to Cb. While this assumption appears reasonable given that the adsorbed reactant coverage is typically low, θCO e 0.1,7 approximate reactionorder estimates were nonetheless extracted as follows. Plots of log icorr versus E were constructed using the above procedure not only for saturated CO solutions but also for lower concentrations, the latter being obtained by sparging the COcontaining solution with argon for short (e30 s) periods. The ensuing reactant concentrations were determined from the voltammetric peak current densities, ip, by means of the usual relationship,18 taking the reactant diffusion coefficient, D, as 2 × 10-5 cm2 s-1 19 and also the apparent anodic transfer coefficient, Rapp, as 0.7.20 Reactant concentrations were varied in this manner by about 10-20-fold below the saturated value, ca. 0.9 mM. Independent estimates of Rapp [d(RT/F)(d ln i/dE)] were also obtained by inspecting the potential shift at a fixed point in the voltammetric wave (such as the current halfway to ip) induced by altering the sweep rate, ν, utilizing the relation17
Rapp ) (RT/2F)(d ln ν/dE′)
(2)
where E′ refers to a given “point” on the voltammetric wave, most conveniently the half-peak potential or the peak potential itself. The sweep rate could conveniently be varied over the range 20-500 mV s-1 for this purpose and also to expand the range of icorr values accessed prior to the occurrence of substantial diffusion polarization at potentials corresponding to ip. Typical Tafel (i.e., log icorr - E) plots obtained for varying reactant concentrations in this fashion on Au(110) in 0.1 M HClO4 are shown in Figure 3. The vertical displacement between the lines for each C0 value, specifically the (∆ log icorr/∆ log C0)E slope, yields the reaction order, m. Admittedly, the extraction of the highest potential segments of the Tafel plots, where C0 falls most significantly below Cb, strictly requires that m ∼ 1, even though the lower potential portions are insensitive
to m. The average reaction order, however, is determined to be persistently close to unity (1.0 ( 0.1) from these and other data on Au(110) in both acidic and alkaline media. Corresponding data obtained on Au(210) tended to yield slightly subunit (0.85 ( 0.1) m values. This result, however, is probably partly artifactual, arising from the mild yet significant chemisorption of the CO reactant which occurs on this surface,7 distorting slightly the log i-E traces especially at potentials close to the bottom of the voltammetric wave. A partial summary of kinetic data determined in this fashion for six gold faces, including (1 × 1) and reconstructed Au(100), is given in Table 1. For convenience, values of Rapp extracted from Tafel slopes are given along with rate constants, kE, determined at a fixed electrode potential of 0.3 V vs SCE, obtained from the usual first-order relationship
kE ) icorr/nFCb
(3)
As emphasized in our earlier study,7a the kE values are markedly dependent on the crystallographic orientation. Nonetheless, the rate parameters were found to be essentially independent of ionic strength in both acidic and alkaline media varied over the range 0.1-1 M by the addition of LiClO4. Of central interest here is the dependence of the electrooxidation rates upon the solution pH. This stems from the mechanistic insight that thereby may be obtained into the role of hydroxide ions, especially when the pH is varied over a wide range. An inevitable stumbling block is the provision of electrolytes with reasonably buffered pH. For solutions having pH values below about 4 and above 10, this could be satisfactorily achieved by the use of appropriate concentrations of HClO4 or NaOH, respectively, with NaClO4 added to yield an ionic strength of 0.1. (An exception was pH 0, where concentrated HClO4 was employed instead.) For the middle pH range, however, specific buffer solutions were required. These were as follows: pH 5, 0.05 M NaC2H3O2 + 0.018 M HClO4 + 0.032 M NaClO4; pH 6, 0.048 M KH2PO4 + 0.003 M Na2HPO4 + 0.05 M NaClO4; pH 6.9, 0.025 M KH2PO4 + 0.025 M Na2HPO4 + 0.05 M NaClO4; pH 9.2, 0.01 M Na2B4O7 + 0.09 M NaClO4. Shown in Figure 4A are Tafel plots for the electrooxidation of 1 mM CO on Au(210) in 12 electrolytes varying in pH by over 13 units. In almost all cases, the E-log icorr slopes are closely similar, 95 ( 10 mV. The sole exception is the pH 5 acetate buffer which yields a higher Tafel slope, 125 mV. This deviation is most readily ascribed to the influence of acetate adsorption decelerating the reaction increasingly toward higher potentials. A similar, although more marked, effect is seen on Au(210) at higher potentials in sulfuric acid electrolyte:7a in both cases the reaction kinetics encompass a region above the potential of zero charge for Au(210), Epzc ≈ -0.1 V vs SCE,13 where oxyanion adsorption should become more prevalent. Figure 4B shows corresponding Tafel plots over a similar range of pH for CO electrooxidation on Au(110). Closely comparable
2326 J. Phys. Chem., Vol. 100, No. 6, 1996
Edens et al.
Figure 4. (A) Logarithmic plots of current density (corrected for diffusion polarization) versus electrode potential for electrooxidation on Au(210) of 1 mM CO in electrolytes of differing pH, as indicated. See text for details and identity of electrolytes. (B) As for (A), but on Au(110).
Tafel slopes were obtained, around 95 mV. Data are not shown for Au(110) in pH 5 or 6 media (acetate, phosphate buffers) since “distorted elongated” voltammograms were obtained, yielding curved log i-E plots with anomalously small Rapp values toward higher potentials. This finding is suggestive of a tendency of anion adsorption to commence at lower electrode potentials on Au(110) than on Au(210). Of chief significance are the marked shifts to lower electrode potentials of the log i-E segments as the pH is progressively increased (Figure 4A,B). Strictly speaking, examining the reaction order in hydroxide (or, equivalently, in hydronium ions) requires constructing log i-pH plots at constant electrode potential. Given the sensitivity of the electrooxidation rates to the potential, this procedure necessarily incurs substantial log i-E extrapolations in order to encompass even a fraction of the entire pH range. Nevertheless, the extent of the extrapolations is minimized by selecting a suitably intermediate electrode potential. Parts A and B of Figure 5 show typical examples of such log icorr-pH plots for Au(210) and (110), respectively, both referring to an electrode potential of 0 V. Evident in both these plots is the presence of two separate log icorr-pH regions for pH values below about 2 and above 5, within which roughly unity log icorr-pH slopes are obtained, being separated by a short “arrest” pH region that exhibits a markedly smaller slope. Essentially similar behavior is obtained if other electrode potentials are selected for data extrapolation. Nonetheless, given the inevitable uncertainties in the validity of such extrapolations, it is instructive also to examine plots of
Figure 5. (A) Plot of logarithm of current density (corrected for diffusion polarization) on Au(210) for 1 mM CO at 0 V vs SCE versus electrolyte pH. Obtained by extrapolation of log icorr-E data in Figure 4A. (B) As for (A), but on Au(110), extracted from data in Figure 4B.
the electrode potentials at a fixed reaction rate, Ei, as a function of pH. While perhaps of less fundamental significance than a true pH-reaction order analysis, such plots avoid the need for lengthy data extrapolation since the measured range of electrode potential necessarily encompasses that required to adjust the electrochemical reaction rates to the appropriate range accessed by LSV. Parts A and B of Figure 6 show such plots (filled squares), again for Au(210) and (110), respectively, all referring to a fixed icorr value of 250 µA cm-2. As indeed can be anticipated from the appearance of the pH-dependent Tafel plots themselves (Figure 4A,B), the Ei-pH plots again show three regions, with higher slopes found for pH values below about 2 and above 4 (labeled regions III and I, respectively), with noticeably smaller slopes obtained within the intervening pH range (labeled II). Also included in Figure 6A,B are the corresponding pH-dependent potentials for the onset of anodic oxide formation, Ea, also obtained from voltammetric data (filled circles). The roughly constant, albeit substantial (ca. 0.8 V), difference in electrode potential between these points and the Ei values is discussed below. Given the implication from the above data that distinct ratedetermining hydroxylation steps are involved, depending on the pH range, it is of particular interest to examine the effects of
CO Electrooxidation on Au Surfaces
J. Phys. Chem., Vol. 100, No. 6, 1996 2327
Figure 7. Ratios of diffusion-corrected current densities for CO electrooxidation in aqueous to deuterated solvent as a function of electrolyte pH on Au(210) (filled circles) and Au(110) (open circles).
Discussion
Figure 6. (A) Electrode potential for CO electrooxidation at a fixed rate (250 µA cm-1) on Au(210) plotted against the electrolyte pH (squares). Obtained by extrapolation of log icorr-E data in Figure 4A. Corresponding circles refer to onset of gold surface oxidation in the absence of solution CO. (B) As for (A), but on Au(110), extracted from data in Figure 4B.
solvent deuteration. To this end, kinetic data for CO electrooxidation on both Au(210) and (110) were obtained for otherwise identical conditions over a wide pH range. Ratios of the electrooxidation rates, icorr, at a convenient electrode potential within the rising portion of the voltammetric wave, (iH/iD)E, are plotted as a function of pH in Figure 7, the filled and open circles referring to Au(210) and (110), respectively. While some scatter is inevitably present, immediately evident is a marked dependence of the solvent isotope effect on the pH, the (iH/iD)E values increasing sharply from essentially unity at pH values above ca. 5 to current ratios close to 6 ( 1 for pH values below 3. Interestingly, these two regions correspond closely to the distinct pH regions I and II/III, respectively, identified from the above pH reaction-order analysis. Some measurements of electrochemical activation parameters were also undertaken to ascertain whether such mechanistic differences are manifested clearly in either enthalpic or entropic contributions to the reaction kinetics. So-called “ideal” activation enthalpies, ∆Hi*, were evaluated from the temperature dependence of icorr at a fixed nonisothermal cell potential, i.e., essentially at a constant Galvani surface potential (cf. ref 10). Values of ∆Hi* varying from about 30 to 45 kJ mol-1 were typically obtained in both acidic and alkaline media. However, no systematic marked pH dependence of ∆Hi* could be discerned.
As already mentioned, the foregoing pH-dependent rate data for CO electrooxidation indicate the occurrence of three distinct kinetic regions. For two of these, at low (e2) and high (g5) pH values (regions III and I, respectively), the kinetics display essentially a first-order dependence on the hydroxide concentration. While the narrowness of the intervening pH region (ca. pH 2-4) prevents a quantitative analysis, the appearance of the log rate-pH data (Figure 5A,B) at least suggests that a [OH-]independent pathway may predominate here. A simple mechanistic scheme, commonly considered for the homogeneous-phase catalytic oxidation of CO in aqueous (or water-containing) media by metal complexes, involves the formation of a hydroxycarbonyl intermediate, M(CO2H), i.e., involving OH binding to the coordinated CO moiety via the C atom, followed by decomposition to yield CO2.1 The relevant steps can be written as either
M-CO + OH- / M(CO2H) + e-
(4a)
M-CO + H2O / M(CO2H) + e- + H+
(4b)
or
where M is a metal coordination site, followed by
M(CO2H) f CO2 + e- + H+
(5)
Note that these steps (eqs 4a,b and 5) are written here as oneelectron components of the overall 2e- process. This situation can be readily envisaged in electrochemical catalytic systems such as here since the metal electrode can act as a flexible “electron sink”. For metal complexes, on the other hand, the stability of particular oxidation states can clearly influence the charge state of intermediate species. Moreover, the occurrence of true homogeneous-phase catalytic cycles for the WGSR requires that the protons generated by water oxidation of CO are reduced to hydrogen by the same metal complexes that catalyzes the oxidative steps, thus restricting the choice of metal catalyst systems and conditions (pH, etc.) for which the WGSR can operate.1c This limitation is probably partly responsible for
2328 J. Phys. Chem., Vol. 100, No. 6, 1996 the dearth of mechanistic kinetic studies, such as pH-reaction order analysis, for metal complex-catalyzed CO oxidations. Nevertheless, clear evidence for the formation of hydroxycarbonyl complexes has been gathered in some cases not only from kinetic studies but also from their analytical detection by spectroscopic means, or even isolation.1a Consequently, then, there is strong precedent to select eqs 4 and 5 as a starting point for examining analogous CO electrooxidations, even though such a proposition has not, to our knowledge, been made explicitly in previous discussions appearing in the electrochemical literature. Considering first the electrode kinetic behavior in alkaline or neutral media (regions I and II), the occurrence of the simple first-order behavior in [OH-] can be rationalized simply as signaling hydroxide discharge onto a CO-containing (or adjacent) site to yield hydroxycarbonyl (eq 4a) as the ratedetermining step (rds). This notion is supported by the observed absence of a deuterium isotope effect under these pH conditions (Figure 7), indicating that deprotonation (or protonation) does not occur during the rds. The observed transition to an apparently pH-independent rds for pH values below ca. 4 (Figure 5A,B) suggests strongly that the oxidant involved is water rather than hydroxide. The coincident appearance of a significant deuterium isotope effect under these conditions (Figure 7) indicates that the rds involves deprotonation, as indeed would be required to form the hydroxycarbonyl intermediate from water discharge onto a COcontaining site as in eq 4b. Since both eqs 4a and 4b at least formally involve 1etransfer, provided that distortions due to double-layer environmental effects are fairly small one would expect the occurrence of Tafel slopes close to 120 mV, i.e., that Rapp ≈ 0.5. The observation of somewhat larger Rapp values (0.6-0.7) on both Au(210) and (110) in both weakly acidic and alkaline media, however, is not inconsistent with the proposed rate-controlling pathways (eqs 4a and 4b) given the likely presence of doublelayer effects associated with the potential-dependent stability of the interfacial reactants (adsorbed CO, OH-, water) and/or the intermediate M(CO2H) species. It remains to consider the mechanistic origin of the kinetics observed at low pH values (e2) (region III), which are also characterized by a rate law which is first order in [OH-] (Figure 5A,B). Even though this rate law is the same as that observed at intermediate and high pH values, the rds is clearly different. This is apparent not only from the clear offset in the log i-pH and (Ei-pH) plots corresponding to regions I and III but also from the presence of a marked deuterium isotope effect within the latter but not the former pH region. A plausible, if not the sole, explanation for the kinetic behavior at low pH (region III) invokes the occurrence of hydroxycarbonyl decomposition to form CO2 (or carbonate) as the rds (eq 5), with the preceding formation of this species from water (eq 4b) being in preequilibrium. This situation will occur at sufficiently low pH (i.e., at high proton concentration) so that the reverse of eq 4b is faster than the decomposition of M(CO2H) to yield products (eq 5), whereupon the latter will become rate determining. Since the concentration of the hydroxycarbonyl intermediate will thereby be proportional to [OH-], the observed reaction rate will become first order with respect to hydroxide, in accordance with the results. A similar situation is apparently encountered for several metal ioncatalyzed CO oxidations in acidic media.1c This mechanism is also consistent with the large deuterium isotope effect observed at low pH (Figure 6), since proton loss will necessarily be involved in hydroxycarbonyl decomposition.
Edens et al. A disquieting feature of this mechanism, however, concerns the log i-E dependence. If the preceding formation of the hydroxycarbonyl species is in quasi-equilibrium, the Nerstian nature of this step should yield a sharply increased stability of M(CO2H) toward higher potentials. At least in the absence of adsorbate-adsorbate interactions or other double-layer environmental effects, higher values of Rapp, ca. 1.5, are predicted on this basis than are observed.7a One possibility is that the separate steps (4a) and (5) are actually concerted or that the adsorbed moiety formed in strongly acidic media involves hydroxyl and carbonyl species in adjacent sites rather than being a true hydroxycarbonyl adduct. Some evidence for the latter possibility on gold in acidic media has indeed been obtained by using surface-enhanced Raman spectroscopy (SERS), specifically from the observation of an apparently surface-oxygen stretching vibration along with the characteristic surface-CO vibration at potentials close to the commencement of CO electrooxidation.22 While the likelihood of these and other alternatives cannot be judged with confidence, the present results nonetheless point to the occurrence of proton transfer during the rds, involving a partly deprotonated carbonyl-oxygen adduct. An additional enlightening view of the energetic role of the proposed hydroxcarbonyl intermediate can be gleaned by comparing the difference between the onset of significant oxidation of carbon monoxide, Ei, and that for the gold surfaces themselves, Ea, plotted in Figure 6A,B. Although some details (including competitive anion effects) remain unclear, the initial stage of gold surface oxidation involves hydroxide or water discharge to form Au-OH species.23 Given the mechanisms discussed above, the roughly constant observed difference between Ea and Ei, 0.8 V, suggests that OH- (or water) discharge onto a M-CO site to form adsorbed hydroxycarbonyl rather than hydroxyl engenders a substantial energetic advantage, by about 75 kJ mol-1 if a 1e- step is involved. Such a stabilization can be understood on thermodynamic grounds from the incipient C-O bond formation. The proposed mechanism therefore provides a satisfying explanation of the intriuging ability of gold surfaces to catalyze the electrooxidation of carbon monoxide at potentials markedly below those corresponding to the initial formation of adsorbed hydroxyl. While the well-known “reactant pair” mechanism, involving reaction between adsorbed hydroxyl (or water) and CO in adjacent sites,24 clearly requires both species to be independently stable under conditions where electrocatalysis occurs, the hydroxycarbonyl mechanism inherently involves a chemical coupling and hence stabilization of the oxidant. One remaining issure involves the possibility of directly detecting by vibrational spectroscopy (or otherwise) the hydroxycarbonyl intermediate itself. Infrared reflection-absorption spectroscopic (IRAS) measurements on Au(210) and other monocrystalline gold electrodes detect typically low fractional coverages (θCO e 0.1) of carbon monoxide in acidic and neutral media from the characteristic C-O stretching (νCO) band at ca. 2100-2115 cm-1 that is removed by the onset of electrooxidation.7 This finding suggests that the adsorbed CO reactant is typically more stable than the ensuing hydroxycarbonyl intermediate: the latter is anticipated to feature a νCO band at distinctly lower frequencies (ca. 1600 cm-1) on the basis of infrared spectra for stable transition-metal hydroxycarbonyl complexes.1a Nonetheless, as already mentioned, moderate or even high coverages of another, less reactive, form of adsorbed CO are detected by IRAS under certain conditions, especially at high hydroxide concentrations.7,12 While this form of adsorbed CO is only prevalent after long (several minutes)
CO Electrooxidation on Au Surfaces solution exposure times at electrode potentials below CO electrooxidation, and at high coverages can act as a reaction inhibitorsdelaying the occurrence of overall reaction toward higher potentials7sit is characterized by a νCO band at lower frequencies, typically 1900-2000 cm-1. It is therefore tempting to suggest that this surprisingly stable form of “adsorbed CO” also involves an intimate interaction between CO and OH, i.e., a hydroxycarbonyl species. Most significantly, the present kinetic results suggest a mechanistic framework for rationalizing (and eventually understanding) the nature of hydroxide/water-induced oxidations of CO, and possibly other feedstocks, on conventional metal electrocatalysts. While largely unaddressed here, the notion of a hydroxycarbonyl-funneled pathway along with its anticipated surface coordination structure sensitivity may also provide a useful basis for considering the observed marked dependence of the CO electrooxidation rate on the gold crystallographic orientation (vide supra). More generally, the aforementioned inorganic catalytic literature together with the present electrode kinetic results provides persuasive reasons to envision a major role of hydroxycarbonyl and related intermediates for oxidations on transition-metal electrocatalysts over a wider perspective. Acknowledgment. This work was supported by the National Science Foundation and the Office of Naval Research. References and Notes (1) For representative reviews of homogeneous-phase catalysis, see: (a) Ford, P. C.; Rokicki, A. AdV. Organomet. Chem. 1988, 28, 139. (b) Laine, R. M.; Crawford, E. J. J. Mol. Catal. 1988, 44, 357. (c) Halpern, J. Comments Inorg. Chem. 1981, 1, 3. (2) For a review of related gas-phase heterogeneous catalysis, see: Campbell, C. T. In Chemical Physics of Solid Surfaces; Elsevier: New York, 1993; Vol. 6, Chapter 9. (3) Parsons, R.; Vandernoot, T. J Electroanal. Chem. 1988, 257, 9. (4) (a) Adzic, R. R. In Modern Aspects of Electrochemistry; Bockris, J O’M., Conway, B. E., Eds.; Plenum: New York, 1990; Vol. 21, p 163.
J. Phys. Chem., Vol. 100, No. 6, 1996 2329 Beden, B.; Lamy, C. Electrochim. Acta 1990, 35, 691. (5) Love, B.; Lipkowski, J. ACS Symp. Ser. 1988, No. 378, 484. (6) Chang, S.-C.; Roth, J. D.; Weaver, M. J. Surf. Sci. 1991, 244, 113. (7) (a) Chang, S.-C.; Hamelin, A.; Weaver, M. J. J. Phys. Chem. 1991, 95, 5560. (b) Chang, S.-C.; Hamelin, A.; Weaver, M. J. Surf. Sci. 1990, 239, L543. (8) See Appendix in: Hamelin, A.; Morin, S.; Richer, J.; Lipkowski, J. J. Electroanal. Chem. 1990, 285, 249. (9) Hamelin, A. In Modern Aspects of Electrochemistry; Conway, B. E., White, R. E., Bockris, J. O’M., Eds.; Plenum Press: New York, 1986; Vol. 16, Chapter 1. (10) Weaver, M. J. J. Phys Chem. 1979, 83, 1748. (11) (a) Hamelin, A.; Stoicoviciu, L.; Chang, S.-C.; Weaver, M. J. J. Electroanal. Chem. 1991, 307, 183. (b) Hamelin, A.; Weaver, M. J. J. Electroanal. Chem. 1987, 223, 171. (12) Edens, G. J. Ph.D. Thesis, Purdue University, May 1995. (13) Hamelin, A.; Stoicoviciu, L.; Silva, F. J. Electroanal. Chem. 1987, 229, 107. (14) Gao, X.; Edens, G. J.; Hamelin, A.; Weaver, M. J. Surf. Sci. 1994, 318, 1. (15) (a) Gao, X.; Hamelin, A.; Weaver, M. J. Phys. ReV. B 1991, 44, 10983. (b) Magnussen, O. M.; Wiecher, J.; Behm, R. J. Surf. Sci. 1993, 289, 139. (16) (a) Gao, X.; Edens, G. J.; Hamelin, A.; Weaver, M. J. Surf. Sci. 1993, 296, 333. (b) Hamelin, A.; Stoicoviciu, L.; Edens, G. J.; Gao, X.; Weaver, M. J. J. Electroanal. Chem. 1994, 365, 47. (17) Bard, A. J.; Faulkner, L. R. Electrochemical Methods; Wiley: New York, 1980; p 223. (18) Reference 17, eq 6.3.8 (p 222). (19) Roberts, J. L.; Sawyer, D. T. Electrochim. Acta 1965, 10, 989. (20) While an estimate of Rapp (as well as D) is required to extract reactant concentrations from peak currents for irreversible voltammograms, the dependence i ∝ (Rapp)1/2 is sufficiently mild so to yield acceptably reliable ((10%) estimates of Cb in this manner. (21) Weaver, M. J. In ComprehensiVe Chemical Kinetics; Compton, R. G., Ed.; Elsevier: Amsterdam, 1987; Vol. 27, Chapter 1. (22) Zhang, Y.; Weaver, M. J. Langmuir 1993, 9, 1397. (23) (a) Angerstein-Kozlowska, H.; Conway, B. E.; Hamelin, A.; Stoicoviciu, L. J. Electroanal. Chem. 1987, 228, 429. (b) AngersteinKoslowska, H.; Conway, B. E.; Hamelin, A.; Stoicoviciu, L. Electrochim. Acta 1986, 31, 1051. (24) Gilman, S. J. Phys. Chem. 1964, 68, 70.
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