Oxygen-Enhanced Dissolution of Platinum in Acidic Electrochemical

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Oxygen-Enhanced Dissolution of Platinum in Acidic Electrochemical Environments Masashi Matsumoto,* Takashi Miyazaki, and Hideto Imai* Green Innovation Research Laboratories, NEC Corp., 34 Miyukigaoka, Tsukuba, 305-8501, Japan ABSTRACT: We investigated the influence of oxygen on the dissolution of platinum in acidic electrochemical environments, which is closely related to the degradation of the cathodes in operating polymer electrolyte fuel cells (PEFCs), by using carbon-supported platinum nanoparticles (Pt/C), bulk polycrystalline Pt electrodes, and Pt(111) single crystals. We found that the platinum dissolution under potential cycling was significantly enhanced by the presence of oxygen even in the lower potential regions (OH adsorption potentials under inert atmospheres): the amount of dissolved platinum increased by factors of 1.2, 18, and 19 on the Pt/C, polycrystalline Pt, and Pt(111) surfaces, respectively, when the potential of the electrodes was cycled under O2. The results of electrochemical and electrochemical scanning tunneling microscopy measurements revealed that a few monolayers of the platinum surface were oxidized, and the surface was roughened by the exposure to oxygen, even in the low potential regions, where no surface reconstruction occurs under N2. The electrochemical dissolution of such surface oxides most likely enhances the dissolution of the platinum under the potential cycling in an oxygen atmosphere.

’ INTRODUCTION Due to the recent commercialization of polymer electrolyte fuel cells (PEFCs), although they are still in limited areas and have limited usages, the improvement of their reliability and durability has been recognized as a major research interest.1
3 Among the performance-reduction factors of PEFCs in longterm operation, such as the dissolution or sintering of platinum electrocatalysts, carbon-support corrosion, and membrane degradations, the electrochemically active surface area (ECSA) loss that arises from effective platinum area loss is one major factor, accounts for at least one-third of the total loss.4,5 Even though platinum is the most stable materials, its dissolution to electrolytes is innegligible under severe PEFC cathode conditions, viz., under strongly acidic and at high potential environments. The stability of bulk platinum in aqueous solutions has been described on the basis of the chemical equilibrium. According to Pourbaix, platinum electrochemically dissolves above 0.85 V vs reversible hydrogen electrode (RHE) in aqueous systems,6 and the amount of dissolved platinum increases exponentially with respect to the applied potential, obeying Nernst equation.7 For PEFC-cathode stability analyses, however, such simple treatment with the Nernstian relationship for bulk Pt is insufficient, because the surface properties of catalyst nanoparticles are different from those of the bulk8,9 and the electrochemical environments of the electrodes—the potential and atmospheres—are in nonequilibrium states.10 Indeed, the concentration of dissolved platinum from the Pt nanoparticles (NPs) has been found to be greater than that of the bulk:11 One reason is that the dissolution equilibrium of Pt NPs shifts to lower potentials due to the Gibbs
Thomson effect, r 2011 American Chemical Society

which is caused by the surface energy increase as the particle size decreases.8,9 The very thin oxide layer formed on the surface of Pt NPs might no longer have a passivative nature. This also could accelerate the dissolution. In addition to such intrinsic surface modifications of NP surfaces, the results of recent research suggested the importance of the dynamical reconstruction of the NP’s surfaces accompanied with oxide formation and reduction during the potential-cycling12,13 in fuel cell operation, to understand the degradation phenomena.15
17 Another factor that could influence the Pt dissolution is the oxygen that is present as an oxidant in the PEFC cathodes, although the degradation analyses that take into consideration the influence of oxygen have been very limited to date.14 If the dynamical surface reconstructions associated with the oxidation
reduction cycles truly enhance the dissolution, the oxygen molecules that dissociate on platinum forming surface oxides possibly work as the promoters of the dissolution even at a lower potential where no surface oxidation occurs in an inert atmosphere.15
17 We, therefore, investigated the influence of oxygen on the platinum dissolution during potential cycling, specifically focusing on the lower potential regions where no surface reconstruction occurs in inert atmosphere (the upper potential limit was chosen at the OH adsorption region), to establish a more realistic picture that describes the platinum dissolution phenomena involving the nonequilibrium states relevant to PEFC cathodes. Received: March 1, 2011 Revised: April 27, 2011 Published: May 16, 2011 11163

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The Journal of Physical Chemistry C The platinum dissolution analyses in an aqueous H2SO4 solution in either a N2 or O2 atmosphere were conducted using carbonsupported platinum nanoparticles (Pt/C), polycrystalline Pt, and Pt(111) single crystals. Furthermore, atomic-level surface analyses were carried out by using Pt(111) single crystals via both electrochemical (by using surface-selective hydrogen adsorption behaviors) and EC-STM measurements. The results showed that the oxygen in the electrolytes enhances the platinum dissolution in the lower potential regions below the OH adsorption oxidation levels [approximately, 0.95 V vs RHE for Pt/C and bulk Pt, 1.25 V vs RHE for Pt(111) surface]. The oxygen molecules dissociate at the platinum surface and form subsurface oxides that correspond to the initial place-exchange oxidation level even in the low potential regions. These surface oxides, which are formed by the exposure to oxygen during the oxygenreduction reaction, are likely responsible for the oxygenenhanced platinum dissolution during potential cycling.

’ EXPERIMENTAL SECTION Carbon-Supported Platinum Nanoparticles. Highly dispersed carbon-supported platinum nanoparticles (Pt/C, 50 wt % metal-loading) were prepared using an oxide-colloidal method.18 The average particle size was approximately 2 nm. The catalysts were fixed on a carbon electrode (HOPG 10 mm in diameter) using Nafion for the electrochemical measurements. Polycrystalline Platinum Bulk Electrodes. Platinum polycrystalline electrodes with a diameter of 10 mm were obtained from Nilaco. Prior to the electrochemical measurements, the platinum polycrystalline electrodes were annealed in a hydrogen flame and quenched into ultrapure water saturated with hydrogen to obtain the clean surfaces. Pt Single Crystals. Pt(111) single crystal electrodes were made from a platinum wire (purity 99.99%) by using the Clavilier method.19,20 The single-crystal bead orientation was determined by using a laser-beam reflection method, and the (111) plane was exposed by using mechanical polishing with finer grade diamond pastes down to 0.25 μm at an accuracy of 0.2°. Then, atomically flat surfaces were created by using hydrogen flame annealing that is similar to that used for the polycrystalline electrodes. Electrochemical Measurements. The electrochemical measurements of Pt/C, polycrystalline Pt, and platinum (111) single crystal electrodes were carried out in a hanging meniscus configuration by using a standard three-electrode electrochemical cell with a gold-gauze counter electrode (to avoid platinum dissolution and redeposition onto the working electrodes) and a RHE reference electrode in a 0.5 M H2SO4 aqueous solution. The Pt/C and polycrystalline electrodes were electrochemically cleaned and activated by repeating CV cycles between 0.05 and 1.0 V at a scan rate of 50 mV/s. The electrochemically active surface areas (ECSA) were estimated via the hydrogen adsorption charge in the cyclic voltammograms (CVs). Hydrogen adsorption charge was estimated to be 210 μC/cm2 for Pt/C and polycrystalline Pt electrodes from hydrogen adsorption waves in the potential region from 0.05 to 0.4 V and 240 μC/ cm2 for Pt(111) single crystal electrode in the potential region from 0.05 to 0.5 V. Then oxidation charges were determined by integrating the corresponding oxidation currents in the potential regions above 0.75 V for Pt/C and Pt electrodes and above 0.88 for Pt(111) single crystal electrode. The degree of oxidation was evaluated by QO/QH assuming successive oxidation reactions

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(see eqs 1
3). The electrochemical dissolution measurements were carried out in both the potential cycling [from 0.6 to 0.95 V for Pt/C and polycrystalline Pt electrodes and from 0.6 to 1.25 V for Pt(111) single crystal] and constant potential holding modes in either N2-purged or O2-saturated 0.5 M H2SO4 solutions. The upper limit of the potential cycling was chosen from below the OH adsorption potential regions [0.95 V for Pt/C and polycrystalline Pt and 1.25 V for Pt(111) single crystal: the OH coverage was approximately 0.4
0.45]. The amounts of dissolved platinum were measured by using a Agilent 4500 (Agilent Technologies) for the inductively coupled plasma mass spectrometry (ICP-MS) Electrochemical Scanning Tunneling Microscopy (ECSTM) Measurements. Morphological changes caused by electrochemical platinum dissolution were monitored at the atomiclevel by using EC-STM for the Pt(111) single crystals. The ECSTM observations were performed by using a Nanoscope E (Digital Instruments) and homemade Teflon electrochemical cell with a RHE reference electrode and a gold wire counter electrode in an airtight chamber. The STM tips were prepared by electrochemically etching a Pt
Ir wire (0.25 mm diameter) in a 1 M CaCl2 and 0.25 M HCl aqueous solution. The sidewalls of the tips were coated with an electrically insulating polymer to reduce the residual background currents. The EC-STM observations were carried out at 0.6 V in a N2 atmosphere without exposure to air after potential cycling using a homemade airtight chamber.21

’ RESULTS AND DISCUSSION Electrochemical Analyses on Pt Dissolution for Pt/C. Now, let us look at the influence of oxygen on the platinum dissolution on the carbon-supported platinum catalysts (Pt/C). Figure 1 shows the amount of dissolved platinum during the potential cycling carried out between 0.6 and 0.95 V vs RHE in 0.5 M H2SO4 solution in either a N2 or an O2 atmosphere. The upper limit of the potential cycling was set at 0.95 V, assuming PEFC cathodes. The degree of oxidation for Pt/C at 0.95 V remains in the OH adsorption potential regions with a ca. 0.45 OH coverage in a N2 atmosphere, assuming successive oxidation reactions, such as

Pt þ H2 O f Pt
OHad þ Hþ þ e


ðOH adsorptionÞ ð1Þ

Pt
OH f Pt
Oad þ Hþ þ e
ðatomic O adsorptionÞ ð2Þ PtO þ H2 O f PtO2 þ 2Hþ þ 2e


ð3Þ

for platinum oxidation (see the inset in Figure 1). No surface reconstruction due to place-exchange oxidation occurs at this oxidation level. In these potential regions, the amounts of dissolved platinum (normalized by the initial ECSA) linearly increased as the number of potential cycles increased (up to 8000 cycles) for both N2 and O2. The amount of dissolved Pt was greater in the O2 than that in N2 throughout the potential cycles, indicating that oxygen truly enhanced the Pt dissolution. As shown in Figure 1b, the ECSAs were also linearly decreased as the number of potential cycles increased, and the ECSA loss was larger in O2. Electrochemical Analyses on Pt Dissolution on Polycrystalline Pt Electrodes. Next, we examined the dissolution of polycrystalline Pt bulk electrodes eliminating the contribution 11164

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Figure 1. Platinum dissolution experiments on carbon-supported platinum nanoparticles. (a) Amount of dissolved platinum during potential cycling from 0.6 to 0.95 V vs RHE in 0.5 M H2SO4 under N2 (blue) and O2 (red). The scan rate was 100 mV/s. The left inset shows a cyclic voltammogram before potential cycling. The right inset shows the potential variation of oxidation charge relative to the hydrogen adsorption charge. The OH coverage was 0.45 at 0.95 V (the upper limit of potential cycles). (b) ECSA losses during the potential cycles. The inset exhibits CVs before cycling (blue) and after 4000 cycles (red).

from the carbon support degradation and/or the NP migration that could occur on carbon-supported Pt catalysts. Figure 2 shows the Pt degradation behaviors for the Pt polycrystalline electrodes. The polycrystalline Pt electrodes dissolved almost linearly with the increase in the number of potential cycles, which is similar to the NPs in both O2 and N2, but the enhancement of the platinum dissolution under O2 was clearer. The oxygen-enhanced dissolution behaviors were affected by the lower-limit potentials in both the potential-cycling and constant-potential polarization when oxygen was present in the electrolytes. The graph in Figure 3 plots the amount of dissolved platinum after 4000 potential cycles with an upper potential of 0.95 and a lower potential of 0.2, 0.4, 0.6, and 0.8 V in either N2 or O2. The amount of dissolution remains unchanged irrespective of the lower potentials: this is consistent with the previous report with N2.13 Since the surface is not reconstructed at 0.95 V (OH adsorption region), the OH species were reversibly formed and reduced, regardless of the lower-limit potentials. On the other hand, the amounts of Pt dissolution in O2 clearly increased and depended on the potentials of the lower limits: the amount of dissolved Pt was 4 times larger at a lower potential of 0.8 V and more than 15 times at a lower potential below 0.6 V. These results suggest that the dissolution mechanism of Pt in O2 differs from that in N2. A similar trend can be seen in the constant-potential polarization experiments. Figure 4 exhibits the potential dependence of the amount of dissolved platinum under a static potential polarization (the holding time was 50 h). While the amount of

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Figure 2. Platinum dissolution behaviors for polycrystalline platinum electrodes. (a) Amount of dissolved platinum during potential cycling from 0.6 to 0.95 V vs RHE in 0.5 M H2SO4 under N2 (blue) or O2 (red). The scan rate was 100 mV/s. The OH coverage was 0.45 at 0.95 V (the upper limit of potential cycles), which is similar to that for Pt/C. (b) ECSA losses during the potential cycles. The inset exhibits CVs before cycling (blue) and after 4000 cycles (red). The measurable ECSAs look unchanged for both N2 and O2, since a new surface was created on the polycrystalline surfaces when Pt dissolved during potential cycling.

Figure 3. Amount of dissolved platinum in potential cycling experiments (after 4000 potential cycles) under N2 (blue) and O2 (red). The upper-limit potential was fixed at 0.95 V vs RHE and the lower-limit potential was changed to 0.2, 0.4, 0.6, and 0.8 V. Overall polarization times were 3.3, 7.8, 12.2, and 16.7 h.

dissolved platinum in N2 was very small and does not depend on the polarization potentials, the amount of dissolved Pt in O2 was larger than that in N2, especially at lower potentials (except for 0.8 V). As we discuss below, the Pt oxides are stable in O2, even at lower potentials, and the oxidized Pt surfaces, which presumably have a passivative nature, were not fully reduced to a bare Pt 11165

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Figure 4. Amount of dissolved platinum in constant potential polarization experiments under N2 (blue) and O2 (red). Overall polarization time was 50 h.

Figure 6. Cyclic voltammograms of Pt(111) single crystals in 0.5 M H2SO4 taken after potential cycling between 0.6 and 1.25 V under N2 (a) and O2 (b).

Pt Dissolution on Pt(111) Single Crystals: Electrochemical and EC-STM Analyses. We then investigated the dissolution

Figure 5. Results from platinum dissolution tests on Pt(111) single crystal electrodes: (a) Amount of dissolved platinum during potential cycling from 0.6 to 1.25 V vs RHE in 0.5 M H2SO4 under N2 (blue) and O2 (red). The scan rate was 100 mV/s. The OH coverage was 0.41 at 1.25 V (the upper limit of potential cycles), which is similar to that in Pt/C. (b) ECSA losses during the potential cycles.

surface above 0.8 V. This could be one of the possible causes for the potential dependence of the Pt dissolution in O2, and this also suggests that the surface structural changes during the potential cycles across the Pt oxidation and reduction boundaries might be crucial for the oxygen-enhanced dissolution phenomena. We note that despite the fact that the overall duration of dissolution tests was much longer for the constant-potential experiments (50 h), the amount of dissolved Pt is much larger for the potential-cycling tests (16.7 h at most). This, again, indicates that the surface reduction process (potential cycling) is important for the enhanced dissolution.

phenomena on Pt(111) single crystals by conducting electrochemical and EC-STM measurements to obtain the atomic level information on the oxygen-enhanced dissolution behaviors. Figure 5 shows the results from dissolution analyses on the potential cycling in either N2 or O2. The upper limit of the applied potential was set at 1.25 V, so the OH coverage was equal to 0.41 (see the right inset in Figure 5a). Similar to polycrystalline Pt, we can see that the amounts of Pt linearly increased, and the amount dissolved in O2 was more than 10 times larger than that in N2. Since single-crystal surfaces exhibit surface-specific hydrogen adsorption waves in electrochemical responses,22 we are able to monitor atomic-level morphological changes in the dissolution processes using the CV measurements. Figure 6 shows the CV curves taken in a H2SO4 solution after the potential cycling between 0.05 and 1.2 V. The shape of the CV (hydrogen adsorption and desorption waves, 0.05 and 0.6 V, respectively) cycled in N2 did not change for up to 8000 times, indicating that atomically flat surfaces were maintained at this oxidation level (at most 0.41 OH coverage) and no surface roughening occurred. In O2, however, the hydrogen adsorption and desorption peaks that are specific to (110) and (100) surfaces (see Figure 6 for the details) gradually grew with the potential cycling, indicating that the (111) surface was roughened by the increasing CV cycles. Since the appearance of clear (110) and (100) waves requires severe surface reconstruction, the OH or O species could possibly enter into the subsurface areas when the surface was exposed to oxygen during the potential cycles, viz., during the oxygen-reduction reaction. We confirmed the surface roughening behaviors caused by the O2 by using EC-STM. Figure 7a,d displays the STM images of the Pt(111) single-crystal surfaces observed after the potential cycling between 0.6 and 1.2 V in either N2 or O2. We can see that an atomically flat surface and a step edge (with 0.23 nm height corresponding Pt-monolayer) were maintained even after 200 cycles when the cycling was done in a N2 atmosphere. Note that 11166

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Figure 7. STM images of Pt(111) single crystal surfaces after potential cycling (200 cycles) under N2 (a) and after the potential cycling under O2 (d) in 0.5 M H2SO4. (b and e) The cross section height along the red the arrows in parts a and d, respectively. (c and f) Schematics of the surface reconstruction models for Pt(111) surfaces.

Figure 8. STM images of Pt(111) single crystals at 0.6 V vs RHE under N2 after being maintained at (a) 0.6, (b) 1.2, (c) 1.45, and (d) 1.5 V (2 h) under N2 (upper images in parts a
d) and O2 (lower images in parts a
d) in 0.5 M H2SO4 (area of STM images: 35 nm  17.5 nm). In an O2 atmosphere, many humps and pits were observed on the terraces of the Pt(111) surface. The heights and widths of the bumps are ca. 0.2
0.3 and 3
5 nm, indicating that monatomic islands formed on the terraces. The formation of these Pt islands is also seen in the reduction reaction of the place-exchange reconstructed surfaces formed at higher potentials in inert atmospheres.

since Pt could dissolve in these potential regions (Figure 5), the very small amount of Pt gradually dissolved from the step edge sites without surface reconstruction in the terrace.23,24 However, the Pt(111) surface was severely roughened after the potential cycles in O2: the formation of many pits with 0.25 nm depths or humps that were 0.2
0.5 nm tall were clearly observed (Figure 7b). These results indicate that roughening with a more-than-three-monolayer Pt height that accompanied the place-exchange oxidation occurred in a potential cycling with O2, being consistent with the appearance of clear (110) and (100) features in the electrochemical measurements (Figure 6). Such irreversible surface reconstruction would originate in formation of higher-order oxides that requires large platinum atom displacements due to atomic oxygen diffusions.15
17 We further investigated the morphology of the Pt(111) surfaces held under static potentials in N2 or O2 up to the higher potential regions to clarify the influence of oxygen on the surface morphological changes more clearly. The holding potentials were set at 0.6, 1.2, 1.45, and 1.5 V, which corresponds to the QO/QH values (under N2) of 0 (double-layer region), 0.3 (OH adsorption region), 1.2 (O adsorption region), and 2.0 (PtO region), respectively. The STM images taken in N2 at 0.6 V after constant potential polarization are given in Figure 8. The flat surface was maintained to at least below the OH adsorption levels (1.25 V) in the N2 (upper images of Figure 8a, b),25 but the surface was roughened due to the partial oxidation with water above the PtO oxidation levels (QO/QH = 1.2): Pt islands were observed at 1.45 V, indicating that the adsorbed O could partly enter the subsurface regions (corresponding R-PtO2 formation by place-exchange16) (upper image of Figure 8c). Above 1.5 V (QO/QH = 2.0), all the surface regions were severely roughened, suggesting that a place-exchange process occurred on all of them (upper image of Figure 8d). On the other hand, in an O2 atmosphere, the surface was severely roughened, even at as low as 0.6 V (the double layer potential region under N2 shown in 11167

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the lower image of Figure 8a). The oxidation level was almost identical to the surface held at 1.5 V in N2, and the degree of oxidation was almost identical, regardless of the potentials (lower images of Figures 8a,b). Namely, the oxygen molecules in the electrolyte react with the Pt surfaces and formed a few monolayers of surface oxides, while an extensive oxygen-reduction reaction should simultaneously occur. Thermodynamically, PtO2 is the most stable phase in an oxygen atmosphere, even at room temperature.26 Indeed, the formation of subsurface oxides including the place-exchange process has been observed on a Pt single crystal in the oxygen gas phase.27,28 In addition, the density functional theory (DFT) calculations predicted that subsurface oxide states are favorable29 and that the OH adsorbate decreases the energy barrier for the O migration from the Pt surface (hcp site) to the subsurface and stabilizes the subsurface oxygen compared to the same on-surface oxygen coverage.30 Molecular dynamics simulation also showed that surface oxides could be formed by dosing oxygen on Pt surface with water molecules.31 Oxygen dissolved in the electrolytes, therefore, could react with the platinum surfaces and form the platinum oxides on Pt surfaces. Our results, consequently, suggest that the oxygen-enhanced dissolution originates in the instability of the surface oxides that are formed by the chemical reaction between the oxygen in the electrolytes, viz., in the PEFC cathode conditions during the oxygen-reduction reaction. Possible Dissolution Mechanism of Platinum Catalysts in the Presence of Oxygen. The dissolution of Pt catalysts has been generally explained by a direct electrochemical (anodic) dissolution path, Pt f Pt2þ þ 2e


E ¼ 1:188 V þ 0:0295 log½Pt2þ  vs SHE ð4Þ

or the indirect chemical dissolution path via the electrochemical oxidation of platinum, viz., Pt þ H2 O f PtO þ 2Hþ þ 2e
E0 ¼ 0:980 V vs SHE PtO þ 2Hþ f Pt2þ þ H2 O

ð5Þ ð6Þ

or, if higher order oxides are formed, via PtO þ H2 O f PtO2 þ 2Hþ þ 2e
E0 ¼ 1:045 V vs SHE ð7Þ

PtO2 þ 4Hþ f Pt4þ þ 2H2 O

ð8Þ 2,3

without taking the influence of oxygen into account. Our present results, however, clearly indicated that oxygen dissolved in the electrolytes significantly affects the platinum dissolution phenomena, simultaneously exhibiting characteristic behaviors; platinum surfaces were oxidized at least in the subsurface oxidation level in an oxygen atmosphere, even at lower potentials where no oxidation occurred in N2. Furthermore, the extent of the dissolution was more significant at the lower potentials, indicative of cathodic dissolution mechanism. Thus, anodic direct dissolution mechanisms (eqs 4
8) are unlikely for the present case. Instead, the cathodic dissolution of such platinum oxides most likely occurs at low potential regions. Thermodynamically, higher-order oxides could dissolve cathodically below 0.837 V, via32,33 PtO2 þ 4Hþ þ 2e
f Pt2þ þ 2H2 O E ¼ 0:837 V þ 0:0295 log½Pt2þ  vs SHE

ð9Þ

Furthermore, Gu and Balbuena examined the dissolution pathways by using the DFT calculations,34 and concluded that the chemical dissolution mechanism, such as 2PtOH þ 2Hþ f Pt2þ þ Pt þ 2H2 O

ð10Þ

PtOOH þ PtOH þ 4Hþ f 2Pt2þ þ 3H2 O

ð11Þ

and

is thermodynamically unfavorable, but the electrochemical dissolution of higher order oxides (PtOOH (Pt3þ)) PtOOH þ PtO þ 5Hþ þ e
f 2Pt2þ þ Pt þ 3H2 O

0:9225 V vs SHE

ð12Þ

and PtOOH þ PtOH þ 4Hþ þ 2e
f Pt2þ þ Pt þ 3H2 O 0:6672 V vs SHE

ð13Þ

34

is favorable. These mechanisms are consistent with our present results on the oxygen-enhanced dissolution phenomena: although our present work could not directly detect the formation of higher-order oxides at lower potentials, the surface roughening of more than three Pt layers indirectly suggested the formation of higher-order oxides. To date, such oxygen-enhanced Pt dissolution has not been taken into careful consideration in the durability investigations of PEFC cathodes, although it seems to provide innegligible contributions on the actual degradation phenomena of the fuel cell cathode particularly in the lower potential regions if oxygen is present on the surface of the platinum catalysts. These results also implicate that an oxygen-reduction reaction could not occur on a bare platinum surface but could occur on an oxidized surface, where the dissociated oxygen atoms, which are the intermediates of the ORR, are buried in the subsurface regions.

’ CONCLUSION In conclusion, we investigated the influence of oxygen on the dissolution of platinum by utilizing carbon-supported platinum nanoparticles (Pt/C), bulk polycrystalline electrodes, and Pt(111) single crystals in order to establish a more realistic picture describing the platinum dissolution phenomena involving the nonequilibrium states relevant to PEFC cathodes. We found that the platinum dissolution was significantly enhanced by the presence of oxygen, even in the lower potential regions (OH adsorption potentials under inert atmospheres). The atomic-level surface analyses we conducted utilizing Pt(111) single crystals via both electrochemical and EC-STM measurements showed that several platinum surface monolayers were oxidized and the surfaces were roughened by being exposed to oxygen during potential cycling and static potential polarization, even in low potential regions (OH adsorption potentials under inert atmospheres), indicating that the oxygen in an electrolyte reacts with the Pt surfaces accompanied with a place-exchange surface reconstruction even at lower potentials. The surface oxides formed in an O2 atmosphere most likely cathodically dissolve during the oxygen reduction reactions. These mechanisms reasonably explain the observed oxygenenhanced dissolution behaviors, and we need to take account of such an oxygen influence for dissolution analyses. 11168

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’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected] (M.M.), [email protected]. nec.com (H.I.).

’ ACKNOWLEDGMENT This work was partially performed under the “Non-precious metal oxide-based cathode for PEFC Project” supported by New Energy and Industrial Technology Development Organization (NEDO).

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(26) Getman, R. B.; Xu, Y.; Schneider, W. F. J. Phys. Chem. C 2008, 112, 9559–9572. (27) Steininger, H; Lehwald, S.; Ibach, H. Surf. Sci. 1982, 123, 1–17. (28) Winkler, A; Guo, X.; Siddiqui, H. R.; Hagans, P. L.; Yates, J. T., Jr. Surf. Sci. 1988, 201, 419–443. (29) Gu, Z; Balbuena, P. B J. Phys. Chem. C 2007, 111, 9877–9883. (30) Gu, Z; Balbuena, P. B J. Phys. Chem. C 2007, 111, 17388–17396. (31) Callejas-Tovar, R.; Liao, W.; Martinez, J. M.; Hoz, D. L.; Balbuena, P. B. J. Chem. Phys. C. 2011, 115, 4104–4113. (32) Bard, A. J.; Persons, R.; Jordan, J. In Standard Potentials in Aqueous Solution; Macel Dekker: New York, 1985; 353 pp. (33) Kawahara, S.; Mitsushima, S.; Ota, K.; Kamiya, N ECS Trans. 2006, 3, 619. (34) Gu, Z; Balbuena, P. B J. Phys. Chem. A 2006, 110, 9783–9787.

’ REFERENCES (1) Borup, B.; Meyers, J. P.; Pivovar, B.; Kim, Y. S.; Mukundan, R.; Garland, N.; Myers, D.; Wilson, M.; Garzon, F.; Wood, D.; Zeleney, P.; More, K.; Stroh, K.; Zawodzinski, T.; Boncella, J.; McGrath, J. E.; Inaba, M.; Miyatake, K.; Hori, M.; Ota, K.; Ogumi, Z.; Miyata, S.; Nishikata, A.; Siroma, Z.; Uchimoto, Y.; Yasuda, K.; Kimijima, K.; Iwashita, N. Chem. Rev. 2007, 107, 3904–3951. (2) Zhang, S.; Yuan, X-. Z.; Hin, J. N. C.; Wang, H.; Friedrich, K. A.; Schulze, M. J. Power Sources 2009, 194, 588–600. (3) Kim, L.; Chung, C. G.; Sung, Y. W.; Chung, J. S. J. Power Sources 2008, 183, 524–532. (4) Shao-Horn, Y; Sheng, W. C.; Ferreira, P. J.; Holby, E. F.; Morgan, D. Top. Catal. 2007, 46, 285–305. (5) Rinaldo, S.; Stumper, J.; Eikerling, M. J. Phys. Chem. C 2010, 114, 5773–5785. (6) Pourbaix, M. In Atlas of Electrochemical Equilibrium in Aqueous Solutions; Pergamon Press: Oxford, 1966. (7) Bindra, P.; Clouser, S. J.; Yeager, E. J. Electrochem. Soc. 1979, 126, 1631–1632. (8) Tang, L.; Li, X.; Cammarata, R. C.; Frisen, C.; Sieradski, K. J. Am. Chem. Soc. 2010, 132, 11722–11726. (9) Tang, L.; Han, B.; Persson, K.; Friesen, C.; He, T.; Sieradski, K.; Ceder, G. J. Am. Chem. Soc. 2009, 132, 596–600. (10) Mitsushima, S.; Kuwahara, S.; Ota, K.; Kamiya, N. J. Electrochem. Soc. 2007, 154, B153–B158. (11) Wang, X. P.; Kumar, R.; Myers, J. P. Electrochem. Solid-State Lett. 2006, 9, A225–A227. (12) Kinoshita, K.; Lundquis, J. T.; Stonehart, P. J. Electroanal. Chem. Interfacial Electrochem. 1973, 48, 157–166. (13) Yasuda, K.; Taniguchi, A; Akita, T.; Ioroi, T.; Shiroma, Z. Phys. Chem. Chem. Phys. 2006, 8, 746–752. (14) Ota, K.; Koizumi, Y.; Mitsushima, S.; Kamiya, N. ECS Trans. 2006, 3, 619–624. (15) Wakisaka, M.; Ashizawa, S.; Uchida, H.; Watanabe, M. Phys. Chem. Chem. Phys. 2010, 12, 4184–4190. (16) Imai, H.; Izumi, K.; Matsumoto, M.; Kubo, Y.; Kato, K.; Imai, Y. J. Am. Chem. Soc. 2009, 131, 6293–6300. (17) Imai, H.; Matsumoto, M.; Miyazaki, T.; Kato, K.; Tanida, H.; Uruga, T. Chem. Commun. 2011, 47, 3538–3540. (18) Watanabe, M.; Uchida, M.; Motoo, S. J. Electroanal. Chem. 1987, 229, 395–406. (19) Clavilier, J.; Faure, R.; Durand, R.; Guinet, G. J. Electroanal. Chem. 1979, 107, 205–209. (20) Itaya, K. Prog. Surf. Sci. 1998, 58, 121–247. (21) Matsumoto., M.; Manako., T.; Imai, H. J. Electrochem. Soc. 2009, 156, B1208–B1211. (22) Clavilier, J.; Achi, K. E.; Rodes Chem. Phys. 1990, 141, 1–8. (23) Komanicky, V.; Chang, K. C.; Menzel, A.; Markovic, N. M.; You, H.; Wang, X.; Myers, D. J. Electrochem. Soc. 2006, 153, B446. (24) Jinnouchi, R.; Toyoda, E.; Hatanaka, T.; Morimoto, Y. J. Phys. Chem. C 2010, 114, 17557–17568. (25) Motoo, S.; Furuya, N. J. Electroanal. Chem. 1984, 172, 339. 11169

dx.doi.org/10.1021/jp201959h |J. Phys. Chem. C 2011, 115, 11163–11169