Surface Oxidation of Platinum-Group Transition Metals in Ambient

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J. Phys. Chem. B 2000, 104, 8250-8258

Surface Oxidation of Platinum-Group Transition Metals in Ambient Gaseous Environments: Role of Electrochemical versus Chemical Pathways Hai Luo, Sungho Park, Ho Yeung H. Chan, and Michael J. Weaver* Department of Chemistry, Purdue UniVersity, West Lafayette, Indiana 47907-1393 ReceiVed: April 3, 2000; In Final Form: June 26, 2000

The effect of water vapor on the temperature-dependent surface oxidation of Pt-group metals in ambientpressure gaseous oxygen environments is explored by means of surface-enhanced Raman spectroscopy (SERS). This exploits the ability of SERS to monitor monolayer-level oxide formation on thin Pt-group films on gold substrates in ambient gaseous as well as solution environments from the characteristic lattice vibrational (phonon) spectra. In contrast to the markedly elevated temperatures (g200 °C) required to initiate surface oxidation on rhodium and ruthenium in dry oxygen, the presence of water vapor triggers monolayer-level oxidation of rhodium and ruthenium surfaces even at room temperature. Exposure of initially reduced rhodium surfaces to wet O2 at different temperatures showed that this catalytic influence of water vapor is limited to ca. 50 °C or below, where water forms a liquid surface film. Rhodium surface oxidation is also observed upon rinsing with aerated water. Related measurements undertaken for rhodium in aqueous electrochemical environments reveal that the electrode potential-dependent formation of metal oxide from water accounts for the water-catalyzed surface oxidation observed in both gaseous and solution-phase oxygen. This follows from the observed ability of O2 electroreduction (to water) to shift the surface potential to sufficiently high values so to trigger water electrooxidation to surface oxide under the open-circuit conditions necessarily pertaining in the gaseous system. This “electrochemical half-reaction” pathway is markedly more facile than the alternative “thermal chemical” route necessarily followed in dry O2. Only slight (submonolayer) surface oxidation of palladium is induced at near-ambient temperatures in gaseous wet O2 , extensive oxide production only occurring above 200 °C, as is the case in dry oxygen. This behavior can also be understood in terms of an “electrochemical” pathway in wet gaseous O2 , the occurrence of O2 electroreduction shifting the potential to insufficiently positive values to induce extensive water electrooxidation to oxide on palladium, due primarily to the lower thermodynamic stability of PdO compared to rhodium and ruthenium oxides. Furthermore, the inability of water to catalyze extensive palladium surface oxidation in gaseous oxygen suggests that oxide formation via a concerted metal-oxygen “place-exchange” mechanism occurs only in conjunction with the “electrochemical half-reaction” pathway.

Introduction Elucidating the factors controlling the formation and removal of oxide films on metal surfaces in contact with gaseous or liquid phases is a topic having broad-based scientific significance as well as enormous technological importance. The Group VIII (platinum-group) metals are of particular interest in this regard given their potency as heterogeneous catalysts for oxidation reactions in the gas phase1 and in aqueous electrochemical environments.2 While Pt-group surfaces readily dissociate and bind oxygen and their metal oxides are thermodynamically stable, oxide film formation often does not occur until higher temperatures or more positive electrode potentials.3 To the extent to which surface oxide formation can inhibit heterogeneous reaction rates,1,2 then, the surfaces can remain highly catalytic even under strongly oxidizing conditions. Establishing the conditions where surface oxide is absent or present on such surfaces in oxidizing environments is therefore of substantial importance. One interesting, yet relatively unexplored, question concerns the conditions necessary for surface oxide formation in oxygencontaining gaseous environments when water is also present, * Corresponding author. E-mail: [email protected].

such as in ambient air. The practical significance of this issue arises in part from the common exposure of catalytic materials to laboratory air at some point prior to, or even during, their chemical utilization. While the potential role of water vapor in influencing surface metal oxide formation in the gas phase may not be obvious at first sight, this question was prompted by recent experiments in our laboratory concerned with examining the oxidation of Pt-group metal surfaces in ambient-pressure gaseous and electrochemical systems using a common in-situ vibrational probe, surface-enhanced Raman spectroscopy (SERS).3,4 As demonstrated previously,5-8 SERS provides a sensitive means of monitoring surface oxide film formation on Pt-group metals down to the monolayer level, utilizing ultrathin (ca. 1 nm) films electrodeposited on gold substrates.5,9 The oxides yield characteristic lattice vibrational bands in the 300800 cm-1 region that are sensitive to the film structure and composition.6,7 A major virtue of this tactic is that the formation of surface oxide, as well as the adsorption of numerous atomic and molecular species, can be monitored on such “overlayermodified” SERS-active surfaces with unique sensitivity. This characteristic, along with freedom from bulk-phase interferences, enables the vibrational properties of chemical interfaces even

10.1021/jp001289+ CCC: $19.00 © 2000 American Chemical Society Published on Web 08/09/2000

Surface Oxidation of Platinum-Group Metals in ambient gaseous or electrochemical environments to readily be examined by means of SERS. Significantly, we found that the propensity for metal oxide formation in dry ambient-pressure dioxygen and aqueous electrochemical environments is markedly different, the former typically demanding temperatures above ca. 200 °C to trigger surface oxidation, while the latter requires only mild overpotentials (i.e., electrochemical driving forces) to induce oxide formation via water electrooxidation.3 Furthermore, both the chemical (free-energy) and electrostatic (interfacial field) driving forces for metal oxide formation are comparable or greater in the gaseous oxygen environment, suggesting further that the electrochemical oxidations proceed via a distinct, inherently more facile, pathway.3 The mechanistic difference was surmised to involve stabilization of the incipient metal cations formed in the metal-oxygen place-exchange process by interfacial water.3 These findings also suggest that interfacial water films, as anticipated for example in wet air, may exert a major influence on the occurrence of Pt-group surface oxidation. The presence of a water film in an oxidizing gaseous medium can also trigger electrochemical “half reactions”, such as oxygen electroreduction coupled with metal electrooxidation, which will not occur in the absence of interfacial solvent. This type of mechanism, akin to well-known processes in metal-corrosion phenomena,10 may furnish a more facile surface oxidation pathway than is available in dry oxygen. Given this situation, we decided to extend our recent SERSbased explorations of the temperature-dependent surface oxidation of Pt-group metals in ambient-pressure oxygen to encompass humidified gases, along with accompanying measurements in aqueous electrochemical environments. The salient findings are present herein. Significantly, the presence of water vapor along with gaseous oxygen is seen to catalyze oxide formation on rhodium and ruthenium at near-ambient temperatures, a process that requires substantially higher temperatures (e200 °C) in dry oxygen. The former observation is consistent with the occurrence of a coupled electrochemical half-reaction oxidative pathway, which is unavailable in the absence of interfacial water. The smaller propensity for surface oxide formation in wet oxygen observed for other Pt-group metals by SERS is also discussed on the basis of this mechanism. Experimental Section Most details of the gas-phase and electrochemical SERS measurements can be found in refs 7d, 11, and 12. A Spectra Physics Model 2017 Kr+ laser provided the Raman excitation at 647.1 nm with ca. 30 mW incident power on the surface. Scattered light was collected into a SPEX Triplemate spectrometer equipped with a Photometrics CCD detector. The gold templates were either a 0.4 cm diameter rod sheathed in Teflon (for exclusively electrochemical experiments), or a 0.6 cm diameter gold foil (99.95%, Alfa-Aesar, for variable-temperature gas-phase or combined gaseous-electrochemical measurements). Electrochemical experiments were performed on the foils by mounting them on an electrode holder with an O-ring seal. The gold surface was roughened so to create optimal and (stable) SERS activity by means of 25 oxidation-reduction cycles in 0.1M KC1, entailing 1 V s-1 sweeps from -0.3 to 1.2 V vs saturated calomel electrode (SCE) and return, holding at the negative limit for 30 s each time.13 The iridium and ruthenium films were prepared by constant-current electrodeposition as described in refs 14 and 15, respectively, and the other three metals as outlined in ref 9. Conditions were arranged so to

J. Phys. Chem. B, Vol. 104, No. 34, 2000 8251 produce transition-metal thicknesses of ca. 3-10 monolayers, thereby yielding near-optimal SERS as well as electrochemical properties.16 The chamber used for the controlled-environment gaseous SERS experiments consisted of a 100 cm3 six-way cross, with pumping and gas-flow arrangements allowing operation in the pressure range 10-3 to 760 Torr. Ambient-pressure gas flow was usually at 100 cm3 min-1. The “wet O2” gas flow (vide infra) involved near-saturation of the O2 stream with water vapor by sparging through liquid water at room temperature. This gas was usually diluted 2-fold with dry nitrogen so to avoid water droplet formation in the SERS chamber, yielding a water partial pressure of ca. 12 Torr. In-situ electrochemical SERS experiments were performed in a conventional three-electrode cell with a quartz optical window. Rotating disk voltammetry was performed with the Teflon-shrouded electrodes, using a Model ASR2 rotator (Pine Instruments). All potentials were measured and are quoted versus the SCE. Results and Discussion Raman Spectral Behavior. In our earlier study,3 we examined SER spectra for five Pt-group metal films - platinum, palladium, iridium, rhodium, and ruthenium - at initially reduced surfaces for (i) progressively higher temperatures in dry O2, and (ii) as a function of electrode potential in aqueous 0.1 M HClO4. As alluded to above, the onset of surface oxide formation in (i) requires elevated temperatures, 200 °C or above for Pt, Pd, Rh, and Ru, although some oxidation on Ir was detected at ca. 100 °C.3 Given that oxide formation on all of these metals, especially Rh and Ru, is thermodynamically favorable in ambient oxygen pressures at room temperature, and extensive O2 dissociation to form chemisorbed atomic oxygen occurs under these conditions, the production of surface metal oxide from adsorbed oxygen clearly involves substantial kinetic barriers.3 Nevertheless, we have observed that significant or even substantial oxide SERS features are commonly evident on ruthenium, rhodium, and iridium following film electrodeposition in 0.1 M HClO4 at negative electrode potentials, and upon subsequent rinsing with water in air. Since it is known from electrochemical and in-situ SERS measurements that oxide films are absent under the former conditions (vide infra), the “wet air” environment is evidently capable of inducing surface oxide formation at room temperature, contrasting the situation with dry oxygen. A systematic series of SERS-based experiments was therefore performed to examine the occurrence of surface oxide formation in water-containing as compared to dry gaseous O2 as a function of the exposure temperature. Some pertinent spectral findings for rhodium surfaces are summarized in Figure 1. The left-hand column (A) shows a series of SER spectra in the 200-800 cm-1 region in aqueous and gaseous environments. The bottom spectrum (a) refers to a freshly electrodeposited film at open circuit in deaerated 0.1 M HClO4, and spectrum (b) in laboratory air following rinsing with water. Spectra similar to (b) were also obtained upon rinsing with acidified aerated water. The broad vibrational feature peaked at ca. 520 cm-1 in (b) characterizes the formation of a surface rhodium oxide film, chiefly Rh2O3 on the basis of X-ray photoelectron spectroscopic (XPS) and other data.3,7b,17 This oxide can readily be removed under mild reducing conditions in either electrochemical or gaseous environments. The latter is illustrated in spectrum (c), which shows no trace of the oxide SER band following exposure to flowing dry hydrogen for 5 min at room temperature (cf., ref 7e). The ability to reform

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Figure 1. (A) Comparison of typical surface-enhanced Raman (SER) spectra obtained for rhodium films in aqueous and gaseous environments. From bottom: (a) freshly deposited film in deaerated 0.1 M HClO4; (b) in laboratory air after rinsing with water, showing formation of oxide; (c) in gas-phase cell after room-temperature exposure to flowing dry H2, showing reductive removal of oxide; (d) after subsequent exposure to wet oxygen, showing similar re-formation of oxide; (e) after subsequent H2 reduction and heating to 250 °C in dry O2, showing thermally produced oxide. (B) Set of SER spectra obtained for rhodium in gas-phase cell for initially reduced surface (a), exposed to flowing wet O2 sequentially at (b) 150, (c) 100, (d) 50, and (e) 25 °C. Surface was reduced each time in H2 prior to wet O2 exposure, before (b)-(e).

spectrally identical rhodium oxide upon subsequent exposure of this reduced surface to flowing wet oxygen (50:50 O2/N2; H2O vapor pressure 12 Torr) at room temperature, is illustrated in spectrum (d). The close similarity between spectra (b) and (d) is clearly evident, indicating that exposure of the rhodium to liquid and vapor-phase water in the presence of oxygen yields a similar oxide film. For comparison, the top spectrum (e) in Figure 1A was obtained for the same rhodium surface after H2 reduction, followed by heating in dry O2 up to 250 °C. A similar oxide feature is obtained, with the peak slightly blueshifted to 540 cm-1. (Identical spectra were also obtained by heating to 250 °C in wet O2.) Comparably blue-shifted rhodium oxide bands, however, are also obtained in aqueous electrochemical environments by increasing the electrode potential to higher values, whereupon thicker (2-3 monolayer) oxide films are produced.4 Having established the potency of water vapor at room temperature to trigger surface oxide formation on rhodium in gaseous oxygen, it is clearly of interest to ascertain the corresponding behavior at elevated temperatures, below that (ca. 200 °C) required for oxide production in dry O2. The righthand column (B) in Figure 1 shows a SER spectral set on rhodium obtained in the gas-phase cell for a reduced surface

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Figure 2. (A) Comparison of typical SER spectra obtained for palladium films in aqueous and gaseous environments. From bottom: (a) freshly deposited film in dearated 0.1 M HClO4; (b) in laboratory air after rinsing with water; (e) after formation of oxide layer at 1.0 V vs SCE in 0.1 M HClO4 (intensity scale compressed by 10× for clarity); (d) oxide layer formed upon heating reduced surface in dry O2 to 250 °C. (B) Temperature-dependent SER spectra for initially reduced palladium film in gas-phase cell (a), after exposure to wet O2 at 25, 100, 150, 200, 250, and 300 °C [spectra (b)-(g), respectively].

exposed to flowing wet oxygen at different temperatures. Spectra (a)-(e) refer to the rhodium surface after H2 reduction (a), and after exposure to gaseous wet O2 sequentially at (b) 150, (c) 100, (d) 50, and (e) 25 °C. The metal surface was reduced by H2 after each step, (b)-(d), prior to further wet O2 exposure. These spectra show clearly that the production of surface oxide in wet oxygen is attenuated severely for gas-exposure temperatures aboVe 50 °C. Thus only a very weak SERS feature is evident at 100 °C and especially 150 °C, whereas a fully developed 520 cm-1 band is obtained upon gas exposure of the reduced surface at 25 °C and 50 °C (Figure 1B). The mechanistic implications of this interesting “anti-Arrhenius” behavior are considered below. In comparison to rhodium, palladium and platinum form oxides having a lower thermodynamic stability (i.e., less negative free energies of formation).3 This difference is reflected in the ca. 0.3 V higher electrode potentials at which the surface oxides are formed and removed on the latter two metals in aqueous electrochemical environments, even though the onset of thermal oxidation in ambient-pressure dry O2 occurs at similarly elevated temperatures, ca. 200 °C, on all three metals.3 The nature of palladium oxide generated in aqueous electrochemical and gaseous oxygen environments is of interest; while an oxide (rather than a hydroxide) is formed in both cases, the SER spectral signatures are noticeably different.3 Column (A) of Figure 2 shows representative SER spectra in the 200-800 cm-1 region obtained in aqueous and gaseous environments.

Surface Oxidation of Platinum-Group Metals The bottom spectrum (a) again refers to a freshly electrodeposited film at open circuit in deaerated 0.1 M HClO4. The lack of detectable oxide SERS bands is consistent with electrochemical information (vide infra). Nonetheless, rinsing with water in laboratory air yields a pair of weak yet discernible bands at 500 and 835 cm-1 [(spectrum (b)]. These features are still substantially (at least ca. 20-fold) weaker than the band centered at 530 cm-1 produced upon forming a layer of palladium oxide at 1.0 V vs SCE in 0.1 M HClO4, shown in spectrum (c). [Note that the Raman intensity scales between (b) and (c) have been adjusted for clarity.] The top spectrum, (d) in Figure 2A, refers to palladium oxide (probably PdO) formed upon heating the surface in dry O2 to 250 °C (cf., ref 3). The marked spectral difference between the oxide formed in the aqueous electrochemical and dry oxygen environments is clearly evident. Similarly to the data for rhodium considered above, SER spectra were also obtained upon exposure of a reduced palladium surface to flowing wet oxygen at various temperatures. (Reduction of residual oxide was accomplished, similarly to rhodium, by flowing dry H2 , usually at 50-100 °C.) A marked difference to rhodium, however, is that only very weak SER features attributable to surface palladium oxide are formed in this fashion even at room temperature, as expected from (b) in Figure 2A. The spectral sequence (b)-(g) in Figure 2B shows the effects of heating an initially reduced palladium surface [spectrum (a)] in wet O2 to progressively higher temperatures, specifically 25, 100, 150, 200, 250, and 300 °C, respectively. (No surface reduction is undertaken between each step in this case.) A weak SERS feature peaked at ca. 600 cm-1 becomes discernible at 100 °C, growing into a substantial band with a partly resolved feature at lower wavenumbers only above 200 °C. While the resulting SER spectrum at 250 °C obtained in wet O2 [(d) in Figure 2B] is somewhat distinct from the example referring to heating in dry O2 [(d) in Figure 2A], the extent of these differences is attributable to variations in sample preparation, etc., rather than signaling an intrinsic dissimilarity in PdO structure. For palladium, then, the presence of water vapor in gaseous oxygen exerts little influence on the nature of the oxide or the thermal conditions necessary for its formation. Qualitatively similar findings also apply to the formation of surface oxide on platinum. Ruthenium, given its 4d7 configuration, forms more stable oxides than either of its second-row neighbors rhodium and especially palladium. Indeed, it is difficult to reduce surface ruthenium oxides by heating in H2, complicating the preparation of reduced SERS-active ruthenium for examination in ambient gaseous environments. Following ref 15, this was achieved by forming a “protective” chemisorbed carbon monoxide layer by sparging CO into 0.1 M HClO4 with the freshly deposited ruthenium film held at -0.3 V vs SCE, where oxides are electroreduced. This surface was then transferred to the gasphase SERS chamber and the CO desorbed thermally by heating to 200 °C in a vacuum. This yielded relatively featureless SER spectra, as exemplified by (a) in Figure 3A. Exposure of this reduced ruthenium surface to wet gaseous O2 at 25 °C, or rinsing in water, yielded broad SERS features centered at ca. 470 and 670 cm-1, consistent with the formation of RuO2.15 Exposure instead at elevated temperatures yielded related yet discernibly different spectra. As an example, spectra (b) and (c) in Figure 3A were obtained for the same surface as in (a) (shown on a common Raman intensity scale) upon exposure to wet O2 at 100 °C, and after cooling to 25 °C, respectively. While the spectral bands are rather ill-defined, there is clear evidence

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Figure 3. (A) Typical SER spectra obtained for initially reduced ruthenium films in gaseous environments. (a) After exposure to wet O2 at room temperature; (b) in wet O2 at 100 °C, and (c) after cooling to 25 °C. Spectra (d) and (e) obtained upon heating reduced surface in dry O2 sequentially to 150 and 300 °C, respectively. (B) Typical SER spectra obtained for initially reduced iridium film in gaseous environments: (a) in air after rinsing with acidified water; (b) after heating to 200 °C in dry O2; (c) after heating to 200 °C in wet O2; (d) in air after rinsing with water (pH 7).

for significant oxide production only upon cooling to 25 °C from the broad envelope developed in the 400-700 cm-1 region at this point [spectrum (c)]. Spectra (d) and (e) in Figure 3A, also shown on a common intensity scale, refer to heating an initially reduced ruthenium surface in dry O2 sequentially to 150 and 300 °C, respectively. While little oxide is apparently formed at 150 °C on the basis of the weak SERS features, pronounced bands at 500-600 cm-1 and ca. 840 cm-1 are evident at the highest temperatures. On the basis of XPS data as well as from Raman spectra of bulk-phase oxide, these features are attributable to RuO2 along with higher oxides, including RuO3 and RuO4.15 Significantly, ruthenium surfaces heated in wet O2 to temperatures above 200 °C yielded similar SER spectra as shown in dry O2 [spectrum (e)]. The last metal considered here, iridium, yielded less clearcut results. We have recently examined in detail surface oxidation of iridium in electrochemical and gaseous environments triggered by NO as well as water/O2.4 In addition to a broad oxide SERS feature centered at ca. 550 cm-1, a sharp band at 570 cm-1 is often observed upon iridium surface oxidation, especially in the presence of NO.4 Unlike the former, once formed the latter band is difficult to remove, even upon heating in H2 or at large negative potentials in 0.1 M HClO4 electrolytes.4 Similarly to rhodium, rinsing a freshly electrodeposited iridium film with laboratory water yields intense SER oxide bands in the 500-700 cm-1 region, as exemplified by the top spectrum (d) in Figure 3B. Surprisingly, however, further rinsing in acidified (0.1 M HClO4) water attenuated, and

8254 J. Phys. Chem. B, Vol. 104, No. 34, 2000 sometimes even removed entirely, these spectral features. This finding contrasts that mentioned for rhodium above, where rinsing with either acidified or neutral water triggered formation of intense SERS oxide bands. We briefly note below possible reasons for this differing behavior. Another surprising finding concerns the temperature-dependent propensity of iridium toward surface oxidation in wet and dry gaseous O2. Unlike rhodium, exposure of a reduced iridium surface to wet O2 at room temperature produced only weak SERS oxide features. Moreover, heating the surface in these two environments yielded largely similar behavior, weak bands centered at ca. 550 cm-1 appearing by 100 °C, and growing further by 200 °C (cf., ref 3). A pair of SER spectra obtained after heating a reduced iridium surface [spectrum (a)] to 200 °C in dry and wet O2 are shown as (b) and (c), respectively, in Figure 3B. Aside from the near-identical form of the spectra, a sharp band at 320 cm-1 is evident. This feature probably arises from adsorbed chloride present during the film electrodeposition.18 Unlike rhodium, the adsorbed chloride band is very difficult to remove, either by water rinsing or polarizing the potential to negative values in an electrochemical cell; this reflects the strength of the Ir-Cl bond.18 As mentioned below, it is not unlikely that the presence of such chloride contamination influences the conditions for forming and removing the oxide film in water-containing environments. Voltammetric Behavior of Surface Oxide. Given the marked differences in the thermal conditions required to form surface oxide in the dry O2 and aqueous electrochemical environments, it is of interest to compare the voltammetric properties of oxides formed via these two routes. This issue is now briefly considered for the palladium and rhodium surfaces of primary interest in the present work. Part A of Figure 4 shows cyclic voltammograms obtained (at 50 mV s-1) for palladium in deaerated 0.1 M HClO4. The electrode potential range, from -0.25 to 1.2 V vs SCE, spans the region between the onset of cathodic hydrogen evolution and the formation of 1-2 monolayers of anodic oxide. The solid trace was obtained for a freshly deposited palladium film onto a gold foil. The broad anodic current plateau above 0.6 V and the sharp ensuing cathodic peak at ca. 0.4 V signal the formation and subsequent removal of the palladium oxide, with the more symmetric cathodic/anodic peaks below 0 V corresponding to chemisorbed hydrogen formation and removal. Such voltammetric signatures are well-known.19 The charge density contained under the oxide reduction peak, 1500 µC cm-2 , corresponds to a ca. 3 ML (monolayer) film, assuming the presence of PdO. However, heating the surface to 350 °C in dry O2, accomplished by transferring the surface into the gas-phase chamber and subsequently back to the electrochemical cell, yielded markedly different initial voltammetric behavior. This is illustrated by the dashed trace in A, which refers to an initial positive-going potential sweep from 0 V following transfer back from the gas-phase cell. Strikingly, little fresh anodic oxide is formed during the positive-going sweep, and the usual oxide reduction peak at 0.4 V on the return scan is entirely absent, being replaced by a very large cathodic peak at ca. -0.2 V, close to the onset of H2 evolution. (Only the base of this sharp cathodic peak is shown in A, for clarity.) The cathodic charge contained under this peak, 3000 µC cm-2, is indicative of extensive (ca. 6 ML) PdO production, not unexpectedly given the high formation temperature. [Heating to lower temperatures yields smaller oxide reduction peaks (i.e., thinner oxide films) as exemplified by the dashed-dotted segment in A, referring

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Figure 4. Cyclic voltammograms (50 mV s-1) for palladium and rhodium films, showing effects of thermal oxidation in dry O2. (A) For palladium films (on gold foil) in deaerated 0.1 M HClO4. Solid trace: for freshly deposited film. Dashed trace: after heating to 350 °C in dry O2, for initial positive-going potential sweep from 0 V. (Dashed-dotted cathodic segment: after heating instead to 250 °C in dry O2.) Dotted trace: second anodic-cathodic cycle, obtained immediately after the dashed voltammogram. (B) Similarly to (A), but for rhodium films (on gold foil) in deaerated 0.1 M HClO4. Solid trace: for freshly deposited film. Dashed trace: after heating to 250 °C in dry O2. Dotted trace: subsequent voltammogram, obtained immediately after dashed trace.

to oxide formed at 250 °C in dry O2.] However, subsequent voltammetric cycles following reduction of the “thermal oxide” yielded current-potential profiles reverting almost entirely to the original morphology, as illustrated by the dotted trace in A, obtained in the second anodic-cathodic cycle, after the dashed voltammogram. This indicates that the no irreversible changes in the palladium surface morphology are incurred as a consequence of the “oxidative heating” step. Note, however, that a small cathodic peak at 0.9 V appears in the subsequent voltammograms (dashed, dotted traces). This feature, arising from gold oxide reduction,9 may partly reflect film “pinholes” caused by the heating step.20 Clearly, then, the oxide produced by elevated-temperature exposure to oxygen undergoes reduction at markedly, ca. 0.5 V, higher cathodic overpotentials than the oxide formed by ambient-temperature water electrooxidation. This difference may well reflect dissimilarities in the palladium oxide structure as well as composition. Indeed, the marked differences seen in the SER spectra for the electrochemical and thermally produced oxidesscompare spectra (c) and (d) in Figure 2Asindicate a dissimilarity in their microscopic structure. It is worth mention-

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ing that a form of palladium oxide, observed to form upon exposure to solution-phase NO, and characterized by a sharp 530 cm-1 SERS band, was also found to be surprisingly resistant to electroreduction.4 Further interpretation of these results, however, is probably unwarranted at present. Corresponding voltammetric data obtained on a rhodium film electrode are contained in part B of Figure 4. Again, the solid voltammetric trace was obtained (at 50 mV s-1) in 0.1 M HClO4 for freshly deposited rhodium on a gold foil. The broad oxide formation wave commencing at ca. 0.6 V, and the main oxide reduction peak located at ca. 0 V on the return potential sweep are characteristic of polycrystalline rhodium electrodes.19 The corresponding dashed trace is an anodic-cathodic cyclic voltammogram obtained after heating the surface to 250 °C in dry oxygen. Again, the anodic current is markedly depressed, indicating the initial presence of thermal oxide. The return sweep features a broad cathodic peak, arising therefore chiefly from electroreduction of the thermal oxide, which is again shifted to lower potentials compared with the corresponding trace for the electrochemically produced oxide. The subsequent cyclic voltammogram, shown as the dotted trace in B, has a form closer to the original (solid) trace, although several cycles are required for the current-potential fingerprint to revert entirely to the latter. Evidently, then, the thermally produced rhodium oxide undergoes electroreduction more sluggishly than the electrochemical oxide, although the differences are rather smaller than in the palladium case. Metal Surface Oxidation Pathways in Wet OxygenContaining Environments. Two features of the results described above are particularly noteworthy. First, the ability of water vapor to trigger surface oxide production on rhodium and ruthenium in gaseous oxygen is restricted to ambient or mildly elevated temperatures, say e50 °C. Second, the catalytic potency of gaseous water/oxygen mixtures, or, similarly, water rinsing in laboratory air, is apparently dependent on the thermodynamic oxide stability, oxide formation under these conditions being extensive only on rhodium and ruthenium, and minor on palladium. We now consider the likely oxide formation mechanisms responsible for these observations, along with additional pertinent electrochemical experiments. Two distinct pathways can be envisaged for the formation of surface oxide in the presence of both oxygen and water:3

O2 f 2Oad

-(1)

or,

H2O f Oad + 2H+ + 2e-

-(2)

Oad + M f MO

-(3)

followed by

where Oad denotes adsorbed oxygen atoms, and MO refers to surface metal oxide. Forming the latter necessarily entails penetration of at least the top layer of metal atoms by oxygen, most likely involving some sort of metal-oxygen “place exchange.” The notion of interfacial M/O place exchange is a familiar one in discussions of electrochemical oxide formation.21 The production of metal oxide from gaseous oxygen, on the other hand, is commonly envisaged to occur via the formation of “subsurface” oxygen from chemisorbed oxygen atoms, followed eventually by diffusion to yield a crystalline surface oxide phase.3

Consequently, then, the microscopic nature of reaction 3 may well be distinctly different in gaseous oxygen and aqueous electrochemical environments (also see below). Indeed, the likely occurrence of reaction 3 via a direct M/O place-exchange step in the latter case has been surmised to be facilitated by the interfacial water solvent, acting to stabilize the incipient metal cations, Mδ+, formed in this process.3 This notion can account for the facile production of metal oxide in the electrochemical case even at room temperature, contrasting the markedly elevated temperatures required in water-free gaseous oxygen.3 Another possible factor may be a differing availability of chemisorbed oxygen (Oad) in the two environments, given its likely role as the precursor species required for metal oxide formation [eq 3]. However, while this species is apparently not readily detectable by SERS,3 there is strong evidence for the formation of at least moderate coverages of Oad upon exposing Pt-group surfaces to dioxygen even at ambient temperatures.22 The lack of facile oxide production in dry O2 cannot therefore be accounted for from a nonavailability of Oad. On the basis, then, the observed occurrence of surface oxide formation on rhodium and ruthenium in water-containing gaseous oxygen at near-ambient temperatures can be ascribed to the facilitation of the M/O place-exchange step by the presence of interfacial water. This notion is also consistent with the observation that the presence of water vapor in the O2 stream triggers oxide formation on rhodium and ruthenium only at temperatures of ca. 50 °C or below. Although quantitative assessments are hampered by a lack of information concerning water adsorption thermodynamics at near-ambient pressures, the presence of a condensed film of water on Pt-group metals in the presence of water vapor is expected to desorb at temperatures, Tdes, only marginally higher than that, Tsat , corresponding to equilibrium with the liquid, i.e., saturated H2O vapor. Since for the water vapor pressure used here (ca. 12 Torr), Tsat ∼ 10 °C, and given that a water monolayer is known to desorb from Pt-group metals in UHV at temperatures only ca. 40-60 K higher than for multilayer water films,23 one anticipates that roughly Tdes ∼ 50-70 °C. We therefore deduce that the presence of at least a condensed monolayer of water, rather than merely gaseous water molecules, is necessary to trigger surface oxide formation of Pt-group metals in gaseous oxygen. This deduction is consistent with the aforementioned notion that the “catalytic” role of water vapor in triggering oxide film formation on rhodium and ruthenium in gaseous O2 at nearambient temperatures arises by lowering the energy barrier to M/O place exchange. However, this idea by itself fails to account for the observed ineffectiVeness of water vapor in triggering oxide film formation on palladium and platinum under these conditions. A more complete picture that accounts satisfyingly for these experimental observations involves considering also the electrochemical half-reactions that can occur at the metal surfaces in contact with gaseous water-oxygen mixtures. In deaerated aqueous solutions, metal oxide formation can only occur via water oxidative deprotonation coupled with oxygen adsorption and lattice incorporation, reactions 2 and 3. In conventional electrochemical environments, the removal of electrons released by reaction 2 is obviously provided by external current flow. In the gaseous H2O/O2 system, as for aqueous O2 solutions, the electrons can be consumed instead by a coupled O2 electroreduction to form water: 1

/2O2 + 2e- + 2H+ f H2O

(4)

The latter half-reaction, well-known in aqueous corrosion

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phenomena, can be combined with the oxidative half reaction 2 plus reaction 3, i.e.,

H2O + M f MO + 2H+ + 2e-

-(5)

to yield an “electrochemical catalytic cycle” featuring the net consumption of oxygen to yield metal oxide, i.e.,

/2O2 + M f MO

1

-(6)

Given that the standard free energy of formation, ∆Gf°, of at least the bulk-phase oxides of all Pt-group metals are negative3 or, equivalently, the standard reversible potentials of reaction 5 under these conditions are lower than for reaction 4, reaction 6 is thermodynamically favorable for each of the systems considered here. However, the catalytic formation of metal oxide by this electrochemical half-reaction mechanism is limited not only by the kinetics of the anodic reaction 5, but also by the notably sluggish rates of the cathodic reduction of oxygen to water [reaction] or peroxide. Even though the Pt-group metals are notable electrocatalysts for O2 electroreduction,24 the onset of measurable cathodic currents for this process commonly requires overpotentials of several tenths of a volt.24 For Pt-group surfaces in contact with oxygen and water solvent, then, the ability of the cathodic reaction 4 to trigger the coupled anodic formation of surface oxide via reaction 5 will depend critically on the ability of the former process to shift the surface potential to sufficiently high values (via interfacial charging from electron consumption) so that the latter step becomes kinetically as well as thermodynamically favorable. These possibilities were examined by measuring currentpotential characteristics for O2 electroreduction at the rhodium and palladium film surfaces employed for the present SERS studies. Measurements were undertaken for O2-saturated solutions at pH 1 (0.1 M HClO4) and pH 7 (phosphate buffer), the latter so to mimic the condition anticipated for solvent films formed at the metal-gaseous water/oxygen interface. On rhodium, the onset of measurable O2 electroreduction, at ca. 0.6 and 0.3 V vs SCE at pH 1 and 7, respectively, occurs at potentials sufficiently above the reversible values for bulk-phase Rh2O3 formation (0.43 and 0.07 vs SCE, respectively) to trigger the anodic production of monolayer oxide via reaction 5. [Although electrochemical polarization to higher potentials, so to form more extensive oxide, yielded more sluggish O2 electroreduction kinetics (and hence smaller cathodic currents) close to the measurable onset of this process, the ensuing opencircuit potentials were not affected significantly by such alterations in the electrode potential history.] The presence of monolayer levels (1-2 ML) of oxide was confirmed by recording cathodic voltammograms commencing at the zerocurrent (i.e., “open-circuit”) potential, close to the onset of cathodic O2 electroreduction, following N2 purging to remove oxygen, and evaluating the charge under the oxide reduction wave. (As might be anticipated, the anodic-cathodic voltammetric features for oxide formation-removal at pH 1 and 7 were similar, the latter features being shifted by ca. 0.35 V to lower potentials, in rough accordance with thermodynamic expectations.) These findings are nicely consistent with the extensive formation of oxide on rhodium both in gaseous water-oxygen mixtures (Figure 1B) or upon rinsing with either neutral or acidified aerated water at near-ambient temperatures, as deduced above from SERS. On palladium, the onset potentials for O2 electroreduction, ca. 0.65 and 0.4 V in the pH 1 and 7 electrolytes, respectively,

are comparable to (if slightly higher than) on rhodium, reflecting the especially facile oxygen electrode kinetics on this metal. However, these potentials are only slightly more positive than the reversible values for bulk PdO formation at pH 1 and 7, 0.50 and 0.14 V, respectively.3,25 Again, the extent of oxide formation under these open-circuit conditions was deduced by cathodic voltammetry after removing the oxygen by N2 sparging. Significantly, only small amounts of PdO, ca. 0.15-0.2 ML, were formed. This finding is consistent with the observation of only weak SERS signals for PdO obtained upon exposure to gaseous water-oxygen mixtures (Figure 2A,B). Consequently, then, it appears likely that in the presence of water-oxygen mixtures under conditions where a condensed water film is present on the metal, surface oxidation occurs via an inherently electrochemical mechanism, involving water electrooxidation (reaction 5) being triggered (in the absence of external current flow) by O2 electroreduction (reaction 4). Note that the net reaction involves the consumption of O2 to form oxide, but with water acting as the Oad source. Since the water is regenerated by O2 reduction, it acts as a catalyst in the overall reaction mechanism. On this basis, then, water plays a role in triggering surface oxide formation in H2O/O2 mixtures beyond facilitating the M/O place exchange process by solvating incipient Mδ+ cations. Definite proof that oxygen lattice incorporation arises via O-transfer from interfacial water molecules, rather than via Oad formed by O2 dissociation, is lacking. However, given that Oad (along with adsorbed OH, OHad) can apparently be produced reversibly on Pt-group electrode-solution interfaces from water via potential-dependent equilibria,21 any Oad and OHad formed by O2 dissociative adsorption in the aqueous environment will be interconverted rapidly with interfacial water, probably in a preequilibrium step. Therefore, the “chemical oxidative” pathway, involving O2 dissociative chemisorption followed by metal oxide formation (reactions 1 and 3), will not be distinguishable from the alternative “coupled electrochemical redox” route (reactions 4 and 5) if, as is likely, steps (1) and/or (2) are rapid and reversible. It is worth mentioning in this context that the varying ability of oxygen-containing acidic aqueous solutions to trigger surface oxide formation on different Pt-group surfaces on open circuit was noted long ago on the basis of electrochemical meaurements,26 the propensity order being Ru ∼ Rh > Ir > Pd ∼ Pt, in essential harmony with the present findings. However, these early workers mistook oxygen chemisorption for oxide formation, and only considered the occurrence of reaction 1. In distinguishing conceptually between the alternative “chemical oxidative” and “coupled electrochemical redox” routes, we emphasize again that the latter (but not the former) is inherently dependent on the surface potential, arising from the exquisite sensitivity of half-reaction thermodynamics and kinetics to this parameter. In the present case, the critical controlling factor is the O2 electroreduction kinetics. The electrochemically irreversible (or “high overpotential”) nature of this process limits the surface potentials that will be attained, which in turn determines the extent of oxide formation as controlled by the (more mildly irreversible) water electrooxidation step. Conversely, the occurrence of such electrochemical half-reactions at metal-gas interfaces, altering markedly the surface potential via interfacial electron transfer and hence surface charging, can also yield substantial changes in chemisorbate structure and bonding. This point has been discussed recently by Weaver et al. in connection with the utilization of chemisorbate vibrational frequencies as a means of monitoring the shifts in surface potential of metal-

Surface Oxidation of Platinum-Group Metals ambient gas interfaces caused by electrochemical-like reactions.27 The notion that surface oxide formation in gaseous oxygen may commonly employ such a “coupled electrochemical redox” pathway is conspicuous by its absence in the literature. However, we have recently invoked such a mechanism to account for the pH-dependent occurrence of oxide films on copper surfaces exposed to aerated water or damp air.28 In that case, the anodic formation of surface Cu2O at higher pH values (>6) is triggered by O2 electroreduction to water and peroxide, although the oxide is spontaneously removed by rinsing with acidic solution, where Cu2O electroreduction is coupled with copper anodic dissolution as Cu2+.28 In contrast, at most Pt-group metal-solution interfaces, anodic oxide formation usually competes effectively with oxidative metal dissolution, even in acidic media. Thus oxide formation on rhodium can be triggered (and maintained) by rinsing with aerated water having acidic as well as neutral pH (vide supra), since the O2-induced anodic formation of oxide is favored on kinetic as well as thermodynamic grounds over metal cation dissolution. A different situation, however, apparently applies to iridium given that, as mentioned above, rinsing with acidic water attenuates the oxide SERS feature. This apparent similarity with the behavior of copper surfaces suggests that the electroreductive removal of oxide may be driven by the anodic dissolution of iridium. Although the electrochemical surface oxidation of iridium is more complex than for the other Pt-group metals, the voltammetric behavior of iridium does not support this possibility.29 A significant complication is the likely presence of adsorbed chloride remaining from the iridium film electrodeposition (vide supra). It is therefore feasible that the formation of Ir-Cl surface complexes may drive the reductive removal of oxide. However, further speculation is unwarranted in the absence of more detailed experimental information. Concluding Remarks: Mechanistic Differences in Wet versus Dry Gaseous Environments. Although partly qualitative in nature, the present findings are believed to provide clear-cut evidence demonstrating the important role played by gaseous water in facilitating the O2-induced surface oxidation of Ptgroup metals. Two aspects of these findings are noteworthy. First, the “anti-Arrhenius” nature of the temperature-dependent catalysis implicates the need for a condensed water film, or at least a water monolayer. Second, the close connection established between the electrode potential-dependent formation of surface oxide in O2-containing aqueous solution and the occurrence of the same process in gaseous O2-water mixtures strongly support the occurrence of electrochemical half-reaction steps in the latter as well as former environments. The second finding implies that the role of interfacial water in the gaseous O2-induced metal oxidation goes beyond facilitating the incorporation of O atoms into the metal lattice by surface solvation. As already mentioned, we proposed in ref 3 that the interfacial solvent acted to facilitate the metal-oxygen placeexchange step by solvating the incipient Mδ+ cation created upon the formation of monolayer surface oxide. This notion is certainly consistent with the findings noted here on rhodium and ruthenium, where the presence of water vapor triggers O2induced oxidation at near-ambient temperatures, which otherwise requires heating to 200 °C or above. However, how does one account for the lack of effectiveness of water vapor in catalyzing O2-induced oxidation on palladium on this basis? If the interfacial water is merely facilitating M-O place exchange, enabling this step to occur at small driving forces at ambient temperatures instead of the markedly (200

J. Phys. Chem. B, Vol. 104, No. 34, 2000 8257 °C) higher values required in dry O2 , then one might expect a similar facility on palladium as for rhodium and ruthenium. We suggest that the inability of interfacial water alone to trigger extensive oxidation on palladium reflects a further fundamental difference in the microscopic nature of monolayer oxide formation occurring via “electrochemical” (EC) and “thermal chemical” (TC) routes. In the former (EC), the direct production of metal oxide can occur by M-O place exchange, driven by the electrochemical thermodynamics and kinetics of water electrooxidation. In the latter (TC), oxide formation is envisaged to require initial production of subsurface oxygen (O ss), entering the metal lattice via vacancies/defects, etc., rather than concerted M-O place exchange, and subsequently yielding oxide crystallites only upon Oss diffusion. Unlike EC, this TC process may not be facilitated by interfacial water. A practical consequence of these notions is that the role of interfacial water in triggering the initial stages of metal oxidation is inexorably intertwined with the occurrence of water electrooxidation, whether driven by external applied electrode potentials or brought about by coupled O2 electroreduction. On this basis, the presence of interfacial water will only engender oxide formation in either gaseous or solution-phase oxygen if the O2 electroreduction is sufficiently facile to shift the potential up to values where water electrooxidation to oxide can proceed. Evidently, these coupled requirements are met for rhodium and ruthenium, but not entirely for palladium, accounting for the experimental findings. In closing, it is also worth mentioning that such electrochemical-type oxidation mechanisms may well also apply to less noble metals, since the onset of electrooxidation requires less positive potentials to be accessed. Such processes may well occur not only in damp air, but possibly at elevated temperature and pressures (e.g., in steam-air mixtures), even though the present mechanism requires the presence of at least a condensed monolayer of water. Acknowledgment. This work is supported by the National Science Foundation (Analytical and Surface Chemistry Program). References and Notes (1) For example: Masel, R. I. Principles of Adsorption and Reaction on Solid Surfaces; Wiley: New York, 1996; Chapter 16. (2) For example: (a) Parsons, R.; VanderNoot, D. J. Electroanal. Chem. 1988, 257, 9. (b) Beden, B.; Leger, J. M.; Lamy, C. In Modern Aspects of Electrochemistry; Bockris, J. O’M., White, R. E., Eds.; Plenum: New York, 1992; Vol. 22, p 97. (3) Chan, H. Y. H.; Zou, S.; Weaver, M. J. J. Phys. Chem. B 1999, 103, 11141. (4) Zou, S.; Chan, H. Y. H.; Williams, C. T.; Weaver, M. J. Langmuir 2000, 16, 754. (5) (a) Leung, L.-W. H.; Weaver, M. J. J. Am. Chem. Soc. 1987, 109, 5113. (b) Leung, L.-W. H.; Weaver, M. J. Langmuir 1988, 4, 1076. (6) Zhang, Y.; Gao, X.; Weaver, M. J. J. Phys. Chem. 1993, 97, 8656. (7) For example: (a) Wilke, T.; Gao, X.; Takoudis, C. G.; Weaver, M. J. Langmuir 1991, 7, 714. (b) Tolia, A. A.; Smiley, R. J.; Delgass, W. N.; Takoudis, C. G.; Weaver, M. J. J. Catal. 1994, 150, 58. (c) Williams, C. T.; Tolia, A. A.; Chan, H. Y. H.; Takoudis, C. G.; Weaver, M. J. J. Catal. 1996, 163, 63. (d) Williams, C. T.; Chen, E. K-Y.; Takoudis, C. G.; Weaver, M. J. J. Phys. Chem. B. 1998, 102, 4785. (8) For recent overviews, see: (a) Weaver, M. J. Top. Catal. 1999, 8, 65. (b) Weaver, M. J.; Zou, S.; Chan, H. Y. H. Anal. Chem. 2000, 72, 38A. (9) Zou, S.; Weaver, M. J. Anal. Chem. 1998, 70, 2387. (10) Bockris, J. O’M.; Khan, S. U. M. Surface Electrochemistry; Plenum Press: New York, 1993; Chapter 8. (11) Wilke, T.; Gao, X.; Takoudis, C. G.; Weaver, M. J. J. Catal. 1991, 130, 62. (12) Gao, X.; Zhang, Y.; Weaver, M. J. Langmuir 1992, 8, 668. (13) Gao, P.; Gosztola, D.; Leung, L-W. H; Weaver, M. J. J. Electroanal. Chem. 1986, 209, 377. (14) Zou, S.; Gomez, R.; Weaver, M. J. Langmuir 1997, 13, 6713.

8258 J. Phys. Chem. B, Vol. 104, No. 34, 2000 (15) Chang, H. Y. H.; Takoudis, C. G.; Weaver, M. J. J. Catal. 1997, 172, 336. (16) Zou, S.; Weaver, M. J.; Li, X. Q.; Ren, B.; Tian, Z. Q. J. Phys. Chem. B 1999, 103, 4218. (17) A band at 290 cm-1, evident in our earlier studies,6,7b is usually weak or absent upon judicious water rinsing as in (a). While originally attributed to an oxide,6,7b we now believe this spectral feature arises from chloride contamination introducing during the metal electrodeposition.3 A similar SER feature is indeed observed for chloride adsorbed on rhodium.18 (18) Mrozek, M. F.; Weaver, M. J. J. Am. Chem. Soc. 2000, 122, 155. (19) For example, see: Woods, R. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1976; Vol. 9, p 1. (20) It is also likely, however, that this peak arises, at least partly, from the exposure of a portion of the gold foil to the electrolyte solution (upon remounting the surface into the electrode holder) that was not coated with palladium in the original electrodeposition step, due to a slight change in the position of the O-ring seal. (21) For example, see: Conway, B. E. Prog. Surf. Sci. 1995, 49, 331.

Luo et al. (22) For example, see: Nolan, P. D.; Wheeler, M. C.; Davis, J. E.; Mullins, C. B. Acc. Chem. Res. 1998, 31, 798. (23) Thiel, P. A.; Madey, T. E. Surf. Sci. Rep. 1987, 7, 211. (24) For example see: Appleby, A. J. In Modern Aspects of Electrochemistry; Conway, B. E., Bockris, J. O’M., Eds.; Plenum: New York, 1974; Vol. 9, Chapter 5. (25) Note that the reversible potential for anodic oxide formation, reaction 5, will necessarily shift to lower values (versus a fixed reference electrode, such as SCE) by 60 mV per increasing pH unit. (26) Rao, M. L. B.; Damjanovic, A.; Bockris, J. O’M. J. Phys. Chem. 1963, 67, 2508. (27) (a) Weaver, M. J.; Williams, C. T.; Zou, S.; Chan, H. Y. H.; Takoudis, C. G. Catal. Lett. 1998, 52, 181. (b) See also: Weaver, M. J. J. Mass Spectrom. 1999, 182/183, 403. (28) Chan, H. Y. H.; Takoudis, C. G.; Weaver, M. J. Electrochem. Solid State Lett. 1999, 2, 189. (29) (a) Mozota, J.; Conway, B. E. Electrochim. Acta 1983, 28, 1; (b) Mozota, J.; Conway, B. E. J. Electrochem. Soc. 1981, 128, 2142.