Cathode Platinum Degradation in Membrane Electrode Assembly

Oct 23, 2008 - Minoru Umeda,*,† Takahiro Maruta, Mitsuhiro Inoue, and Akira Nakazawa. Department of Materials Science and Technology, Faculty of ...
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18098

J. Phys. Chem. C 2008, 112, 18098–18103

Cathode Platinum Degradation in Membrane Electrode Assembly Studied Using a Solid-State Electrochemical Cell Minoru Umeda,*,† Takahiro Maruta, Mitsuhiro Inoue, and Akira Nakazawa Department of Materials Science and Technology, Faculty of Engineering, Nagaoka UniVersity of Technology, Kamitomioka 1603-1, Nagaoka, Niigata 940-2188, Japan ReceiVed: August 12, 2008; ReVised Manuscript ReceiVed: September 26, 2008

The electrochemical dissolution of the Pt cathode catalyst, which determines the lifetime of polymer electrolyte fuel cells (PEFCs), has been studied. A newly developed solid-state electrochemical cell, which consisted of a Nafion membrane sandwiched by Pt/C-based catalyst layers, was employed for the electrochemical measurement. When humidified O2 and H2 gases were supplied to working and counter electrodes, respectively, the Pt electrochemical surface area (ESA) decreased during the potential cycle for 20 h. The highest rate of ESA decrease was recorded for an alternating potential cycle between (i) a positive-potential region of 0.2-0.65 V vs Ag/Ag2SO4 and (ii) a negative-potential region of -0.73 to -0.4 V vs Ag/Ag2SO4. From an SEM-EDS observation of the cross section of a membrane sample after the potential cycle, Pt particles are found in the Nafion membrane, indicating that the decrease in ESA caused by Pt dissolution brings about the Pt deposition in the membrane. Under the assumption that O2 reduction at the negative-potential region generates a large amount of H2O2, EQCM measurement was conducted in an acid solution, which revealed the fact that H2O2 accelerates the Pt dissolution. According to this demonstration, the Pt dissolution in the solid state was carried out by adding H2O2. As a result, the potential range for the Pt dissolution in the presence of H2O2 was determined to be an alternating potential cycle between (i) a positive-potential region of 0.4-0.65 V vs Ag/ Ag2SO4 and (ii) a negative-potential region of -0.2-0.15 V vs Ag/Ag2SO4. It is considered that the mechanism for the Pt degradation in the solid state is the generation of H2O2, which can reduce the Pt-oxide that protects Pt from dissolution. 1. Introduction A carbon-supported Pt electrocatalyst (Pt/C) is widely used at the anode and cathode of polymer electrolyte fuel cells (PEFCs). Pt is known to be a noncorrosive metal; however, it degrades after a long period of operation in PEFCs.1,2 With regard to the Pt degradation, the following phenomena have been reported: (i) a decrease in the Pt electrochemical surface area (ESA),3 (ii) Pt particle growth based on Ostwald ripening,4 and (iii) Pt deposition in the form of a line in a polymer electrolyte membrane, called the “Pt band”.5–7 Each phenomenon has been attributed to the cathode Pt degradation. Many studies on the Pt degradation in electrolytic solutions have been reported.8–17 It can be predicted from a Pourbaix′s potential-pH diagram that Pt dissolves and becomes Pt2+ in a strong acidic solution (pH e 0).8 However, in an acidic solution containing Cl-, Pt dissolves to form soluble complexes such as PtCl42- and PtCl62-.9-11 For 0.5-5 mol/dm3 H2SO4 solution, the Pt dissolution is enhanced by a potential cycle.13-15 In this case, Pt dissolves during the formation and reduction of Pt-oxide.13-16 Kodera et al. reported that Pt drastically dissolves by a potential cycle in concentrated H2SO4 solution of 16 mol/ dm3.17 In these reports, Pt degradation is studied in the liquid phase; however, a solid-state Pt-based electrocatalyst is used in PEFCs. Therefore, the solid-state Pt/C degradation in PEFCs must be investigated. With regard to electrochemical measurements in the solid state, Parthasarathy et al. evaluated the O2 reduction activity of * Corresponding author. E-mail: [email protected]. Tel.: +81-258-47-9323. Fax: +81-258-47-9300. † ACS active member.

Pt using a solid-state electrochemical cell (SSC) that was connected to a Pt disk electrode and a polymer electrolyte membrane.18 Basura et al. also used a similar SSC system to evaluate the permeability of O2 gas from its solubility and diffusion coefficient in a polymer electrolyte.19 Recently, Jiang and Kucernak utilized an SSC composed of nanostructured Pt deposited on an Au electrode and a Nafion 117 membrane to evaluate the oxidation of methanol using the Pt catalyst.20 However, the degradation of a Pt/C catalyst has not yet been investigated using an SSC. Thus, we developed a new SSC system to evaluate the degradation of Pt/C. The SSC consists of a Pt/C-based anode and cathode, a reference electrode, and a Nafion membrane as a polymer electrolyte. In the SSC, H2 and O2 gases can be supplied to the anode and cathode, respectively, through gas diffusion electrodes. The cathode Pt/C was degraded by a continuous potential cycle under a controlled potential. The extent of the Pt/C degradation was evaluated by monitoring the ESA obtained under a deaerated condition. The cross-sectional observation of the electrode and membrane was conducted by scanning electron microscopy (SEM) and energy-dispersive X-ray spectrometry (EDS). On the basis of the obtained results, the Pt/C degradation mechanism will be discussed. 2. Experimental Section Figure 1 shows a schematic illustration of the SSC used in this study. A membrane-electrode assembly (MEA) was installed in the SSC, which was prepared as follows. Nafion 112 (Dupont) was used as a polymer electrolyte membrane. The membrane was boiled in 0.5 mol/dm3 H2SO4 and washed twice by boiling

10.1021/jp8071958 CCC: $40.75  2008 American Chemical Society Published on Web 10/23/2008

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ESA ) Q ⁄ 210

(1)

where Q is the coulomb charge associated with H adsorption/ desorption (in µC) and 210 is the coulomb charge per unit surface area of Pt (in µC/cm2). After the potential cycle, the cross section of the catalystcoated membrane was observed by SEM (JSM-6060A, JEOL). After the carbon papers were peeled from the MEA, the center of the sample was cut using a microtome (cut 4060, MEIWAFOSIS). The cross section of the sample was coated with carbon using an ion coater (CADE-E, MEIWAFOSIS) to suppress the generation of charge. An elemental analysis was then conducted by EDS (JED-2300, JEOL). Moreover, measurement using an electrochemical quartz crystal microbalance (EQCM; QCM922, Seiko EG&G) was performed in 0.5 mol/dm3 H2SO4 as a supporting electrolyte. The frequency of the QCM was 9 MHz. Pt dissolution was evaluated by the weight change of a Pt disk electrode (geometric surface area, 0.196 cm2). The measurement was conducted at a sweep rate of 50 mV/s. The Pt electrode potential was controlled by a potentiostat (model 283, PAR) with respect to the Ag/ Ag2SO4 reference electrode.

Figure 1. Schematic illustration of the solid-state cell used in this study.

in pure water each for 1 h. A commercially available Pt/C electrocatalyst (amount of deposited Pt is 45.9 wt %, Tanaka Kikinzoku) was used for a working electrode (WE) and a counter electrode (CE). The catalyst was ground in a mortar with 5 wt % Nafion solution and a mixed solvent of methanol, 2-propanol, and Millipore water (1:1:1) was added. The obtained slurry was spread over a circular area with 8 mm diameter on both sides of the Nafion 112 membrane. The amount of Pt used was adjusted to 0.2 mg/cm2. Then, the catalyst-coated membrane was sandwiched between carbon papers of 8 mm diameter then hot-pressed at 4.5 kN and 140 °C for 15 min. In addition, a Au mesh in combination with a Au wire was connected to the WE and CE as a current collector, as shown in Figure 1. A Ag/ Ag2SO4 reference electrode (RE) was set on the Nafion membrane. The RE was prepared in accordance with a procedure reported in our previous work.21 Briefly, a Pt wire of 100 µm diameter and 3 mm length was inserted inside the tip of a glass tube, followed by heat-sealing, cutting, and polishing to a mirror finish. Finally, a silver wire deposited with Ag2SO4 was introduced into the glass tube, which was filled with saturated K2SO4 aqueous solution. The prepared RE has been previously demonstrated to perform effectively in the SSC.21 It should be noted that humidified N2 gas was always supplied to the SSC. The cathode Pt/C degradation was conducted by a potential cycle using a potentiostat (HA301, Hokuto Denko). During the operation, humidified O2 and H2 gases were supplied to the WE and CE at a rate of 20 mL/min, respectively. The potential scan speed was 10 mV/s. Before and after the potential cycle, background cyclic voltammograms for the evaluation of the ESA were obtained at a scan rate of 10 mV/s under humidified N2 gas flowing to the WE and CE at a rate of 200 mL/min. The ESA (in cm2) was estimated using the following equation22

3. Results and Discussion 3.1. Decrease in ESA for the Cathode Pt by a Potential Cycle in the Solid-State Cell. In Figure 2, the upper diagram shows voltammograms obtained in the SSC. In this figure, curve (a) represents the background cyclic voltammogram obtained by supplying humidified N2 gas to the WE and CE at a scan rate of 10 mV/s. A typical profile of the Pt electrode similar to that obtained in an acidic solution23 is observed, implying that the prepared SSC can be used in the subsequent experiments. Next, an O2 reduction voltammogram, shown by curve (b), is obtained by feeding humidified O2 and H2 gases into the WE

Figure 2. Voltammograms taken in the solid-state cell: (upper figure) (a) is a background cyclic voltammogram under N2 atmosphere, (b) is an O2 reduction voltammogram, and (c) is a background cyclic voltammogram after a potential cycle between -0.73 and 0.65 V vs Ag/ Ag2SO4 for 20 h under the conditions of curve (b). Scan rate: 10 mV/s. The lower figure shows the rate of ESA decrease versus the potential cycle range. The rate of ESA decrease shown by line C is obtained from the difference between those given by curves (a) and (c). Regions 1 and 2 are described in the text.

18100 J. Phys. Chem. C, Vol. 112, No. 46, 2008 and CE, respectively. The onset potential of O2 reduction is 0.2 V vs Ag/Ag2SO4, which is almost the same as that for the onset of the Pt-oxide reduction.24 Then, a successive potential sweep was conducted between a potential range of -0.73 and 0.65 V vs Ag/Ag2SO4 at a scan rate of 10 mV/s, during which humidified O2 and H2 gases were supplied. After the 20 h potential cycle, the gases were switched to N2 to obtain the background cyclic voltammogram again. The obtained voltammogram, curve (c), is almost the same as that of curve (a). However, currents originating from the H adsorption/desorption are smaller than those in curve (a). According to a calculation based on eq 1, the ESA is estimated to be 108.0 cm2 for curve (a) and 48.5 cm2 for curve (c). For comparison, another experiment was conducted by supplying N2 to the WE, resulting in an ESA of 88.1 cm2. It is clear that the O2 gas affects the decrease in ESA. Subsequently, the relationship between the rate of ESA decrease and the potential cycle range was investigated. In Figure 2, the lower diagram shows the rate of ESA decrease versus the potential cycle range. As for line A, the potential cycle was conducted between -0.73 and 0.2 V vs Ag/Ag2SO4, and the rate of ESA decrease was 0 cm2/h. In the case of line B, the ESA decrease only occurs upon changing the anode reverse potential from 0.2 to 0.4 V. Upon further changing the anode reverse potential to 0.65 V, line C, which corresponds to the result in the upper part of Figure 2, a much larger rate of decrease is observed. When we compare lines A, B, and C, a factor causing the decrease in ESA exists in the potential region between 0.2 and 0.65 V vs Ag/Ag2SO4, which is expressed as region 1 in Figure 2. It should be noted that region 1 is much more anodic than the onset potential of O2 reduction. However, the rate of ESA decrease shown by line E, where the potential cycle is conducted in region 1, is low. This implies that region 1 is merely one factor causing the decrease in ESA. Next, another set of similar experiments was performed for the fixed anode reverse potential of 0.65 V vs Ag/Ag2SO4. The obtained results are shown as lines C, D, and E. It is evident that the ESA decreases in the order of line C > line D > line E. In the same way, the potential region of -0.73 to -0.4 V vs Ag/Ag2SO4, expressed as region 2, contributes to the decrease in ESA. However, the result shown by line A indicates that the ESA does not decrease at all; that is, the decrease in ESA is not only induced during the potential cycle in region 2. According to these results, regions 1 and 2 are found to synergistically induce the decrease in ESA. In other words, the ESA is markedly reduced by an alternating potential cycle between the two regions. 3.2. Cross-Sectional SEM Observation of the CatalystCoated Membrane. After the decrease in ESA, the MEA was removed from the SSC, and the current collector was peeled off. The resulting carbon-coated membrane was cut at the center of the sample to observe the cross section by SEM. Figure 3A shows an image of the whole sample whose rate of ESA decrease is expressed by line C in the lower part of Figure 2. The areas represented by dotted squares are shown as magnified images in Figures 3B-D (B, WE side; C, center; D, CE side). In these images, bright particles of ca. 150-300 nm size can be seen at the center of the Nafion 112 membrane. In addition, according to EDS elemental analysis, Pt is detected around the particles, indicating that the bright particles are made of Pt. This implies that the decrease in ESA is due to the Pt dissolution, a part of which is then deposited as large particles in the Nafion membrane.

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Figure 3. Cross-sectional SEM images of the sample whose rate of ESA decrease is represented by line C in the lower part of Figure 2 (A, whole; B, WE side; C, center; D, CE side). The arrows indicate deposited Pt.

3.3. Weight Change of Pt Electrode in H2O2-Containing H2SO4 Solution. From the lower part of Figure 2, it is considered that O2 reduction in potential region 2 induces the Pt dissolution at the layered structure of the Pt/C-catalyst/Nafion 112 membrane. Region 2 corresponds to the H adsorption/ desorption potential region. Inaba et al. reported that a large amount of H2O2 is formed by O2 reduction at the H adsorption/ desorption region in an acidic solution.25 From these results, the Pt dissolution in region 2 is expected to be related to the generation of H2O2. To confirm this, the weight change of the Pt electrode was measured using the EQCM in 0.5 mol/dm3 H2SO4 containing 10 mmol/dm3 H2O2. Figure 4A exhibits the results obtained by simultaneous measurements of the cyclic voltammetry and the EQCM using a Pt disk electrode (φ 5 mm, geometric surface area is 0.196 cm2) in 0.5 mol/dm3 H2SO4 without H2O2. The shape of the voltammogram does not change upon a successive potential sweep. In the same way, the profile of the EQCM data is independent of the cycle number. In the potential range of 0-0.7 V vs Ag/ Ag2SO4, a weight gain of the Pt electrode is observed during the anodic sweep, and its loss is seen during the cathodic sweep. These phenomena have been explained as being due to the formation of Pt-oxide and its reduction, respectively.12,26 In the case of H2O2 addition, the results are shown in Figure 4B. The shape of the cyclic voltammogram is the same as that measured at the Pt electrode in an electrolytic solution containing H2O2.27 Also, a successive potential sweep does not change the voltammogram shape. However, as for the EQCM data, the weight of the Pt electrode decreases with the number of potential cycles in the measured potential region. This strongly supports the above expectation that H2O2 induces the Pt dissolution. 3.4. Pt Dissolution at the Solid-State Cell in the Presence of H2O2. According to the above demonstration that H2O2 induces the Pt dissolution in an acidic solution, we will now discuss the Pt dissolution in the SSC by adding H2O2. The results are shown in the upper part of Figure 5. Curve (a) in the figure shows a background cyclic voltammogram obtained under humidified N2 gas. The profile is almost the same as that shown

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Figure 4. Cyclic voltammogram and EQCM data obtained at the Pt electrode (A, in 0.5 mol/dm3 H2SO4; B, in 0.5 mol/dm3 H2SO4 + 10 mmol/dm3 H2O2). The geometric surface area of the Pt working electrode is 0.196 cm2. The scan rate is 50 mV/s.

TABLE 1: Rate of the ESA Decrease for Pt/C Obtained by the Potential Cycle with and without H2O2 in the SSCa potential cycle range/V vs Ag/Ag2SO4 0.15-0.65 -0.4-0.65

decreasing rate of the ESAb/cm2/h with

without

1.01c 2.02c

0.34 0.43

a The potential cycle was conducted at 10 mV/s for 5 h while supplying humidified N2 and H2 gases to the WE and CE, respectively. b The background cyclic voltammogram for the ESA was measured at 10 mV/s under flowing the humidified N2 gas to the WE and CE. c The same results are expressed as lines C (0.15-0.65 V) and E (-0.4-0.65 V) in the lower part of Figure 5.

Figure 5. Voltammograms taken in the solid-state cell: (upper figure) (a) is a background cyclic voltammogram under N2 atmosphere, (b) is a voltammogram after adding H2O2 to the WE, and (c) is a background cyclic voltammogram after a potential cycle between 0.15 and 0.65 V vs Ag/Ag2SO4 for 5 h under the conditions of curve (b). Scan rate: 10 mV/s. The lower figure shows the rate of ESA decrease versus the potential cycle range. The rate of ESA decrease shown by line C is obtained from the difference between those given by curves (a) and (c). Regions 1′ and 3 are described in the text.

in the upper part of Figure 2. Then, 100 µL of 5 wt % H2O2 was added to the WE under an N2 gas supply, and the gas supplied to the CE was changed to humidified H2. After the alternation, the electrode potential was swept from the rest potential of 0.16 V vs Ag/Ag2SO4 to the anodic direction and then to the cathodic direction. The obtained voltammogram is shown as curve (b). The profile of the curve compared with that in Figure 4B suggests that the electrode reaction of H2O2 takes place. Subsequently, a potential cycle was conducted

between 0.15 and 0.65 V vs Ag/Ag2SO4 for 5 h. After the potential cycle, the background cyclic voltammogram was obtained again by supplying humidified N2 gas to the WE and CE. The obtained voltammogram is expressed as curve (c), which is the same as curve (a). It is evident that the H2O2 added to the WE is consumed during the potential cycle of 5 h. When the ESAs obtained from curves (a) and (c) are compared, it is apparent that the ESA decreases during the potential cycle in the presence of H2O2. To clarify the effect of H2O2, the decrease in ESA during the potential cycle was compared for the cases with and without H2O2. Table 1 shows a summary of the rate of ESA decrease. It is clear from the table that the rates of ESA decrease without H2O2 are smaller than those in the presence of H2O2. Thus, the ESA decrease is shown to accelerate by the added H2O2. The lower part of Figure 5 shows the dependence of the rate of ESA decrease on the potential cycle range. First, we will compare the results shown by lines A-E, in which all the anode reverse potential is fixed at 0.65 V vs Ag/Ag2SO4. The rate of ESA decrease is in the order of line A < line B < line C < line D e line E. From this comparison, one factor causing the decrease in ESA is presumed to be the existence of a potential region between -0.2 and 0.4 V vs Ag/Ag2SO4. Next, lines E and F, which both have a cathode reverse potential of -0.4 V vs Ag/Ag2SO4, are compared. On the basis of the result that the rate of ESA decrease shown by line E is

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Figure 6. Cross-sectional SEM images of the sample whose rate of ESA decrease is represented by line C in the lower part of Figure 5 (A, whole; B, WE side; C, center; D, CE side). The arrows indicate deposited Pt.

larger than that for line F, another potential region causing a decrease in ESA exists around 0.15-0.65 V, which is almost the same as region 1 of Figure 2. According to the result shown by line G, where the potential cycle is conducted at 0.15-0.4 V vs Ag/Ag2SO4, the rate of ESA decrease is low. This indicates that the decrease in ESA is not induced in this potential region. From these results, the factors causing the decrease in ESA are found to exist at 0.4-0.65 V vs Ag/Ag2SO4 (region 1′) and -0.2-0.15 V Ag/ Ag2SO4 (region 3). According to lines A, B, and F, the rates of ESA decrease are small. Consequently, an alternating potential cycle between regions 1′ and 3 is needed to produce the H2O2accelerated decrease in ESA. After the decrease in ESA in the presence of H2O2, the cross section of a sample was observed by SEM. Figure 6 shows SEM images of the sample represented by line C in Figure 5. The images show that Pt particles of ca. 20-150 nm are deposited in the Nafion membrane on the WE side. This reveals that the decrease in ESA in the presence of H2O2 is also due to the Pt dissolution. 3.5. Prediction of the Cathode Pt Dissolution Mechanism. To consider the cathode Pt dissolution mechanism, the following experiment was conducted. An O2 reduction voltammogram was obtained from the SSC using another membrane-electrode system shown in the inset of Figure 7. The system is constructed as follows. A Pt/C electrode for O2 reduction (denoted as “generator”, φ 5 mm) is set on one side of a Nafion 112 membrane. On the opposite side, a disk electrode (denoted as “collector”, φ 2 mm) and a ring electrode (outer diameter, 7 mm; inner diameter, 5 mm) are set. The Pt amount of the Pt/C was 0.2 mg/cm2. The ring electrode is used as a CE. This system was installed in the SSC shown in Figure 1, and the electrochemical measurement was performed by supplying humidified O2 gas to the generator side and humidified N2 gas to the disk and ring sides. The generator and collector potentials were controlled by a dual potentiostat (700B, ALS). Figure 7 shows the obtained results. The lower voltammogram was obtained at the generator for O2 reduction at a potential

Figure 7. O2 reduction voltammograms obtained in the integrated MEA shown in the inset. Lower: O2 reduction at the generator at a scan rate of 1 mV/s. Upper: collector current response measured at 0.6 V vs Ag/Ag2SO4. The geometric surface areas of the generator and collector are 0.196 and 0.031 cm2, respectively.

sweep rate of 1 mV/s, while the collector electrode potential was held at 0.6 V vs Ag/Ag2SO4, where H2O2 can be detected.25 For the generator, O2 reduction begins at 0.15 V vs Ag/Ag2SO4, and a limited current is observed at -0.4 V vs Ag/Ag2SO4. In addition, although the collector is placed at the opposite side to the generator, as shown in the inset, H2O2 can be detected at the collector. This implies that the H2O2 generated at the generator reaches the collector through the Nafion membrane. Thus, H2O2 moves in the Nafion membrane. On the basis of the above results, the Pt dissolution mechanism at the cathode of the MEA is now discussed. It is wellknown in the O2 reduction reaction that Pt-oxide is immediately formed on the surface of the cathode Pt, according to the following reaction.28

Pt + H2O f PtO + 2H++2e-

(E0 ) 0.980 V vs RHE)

(2) It has been reported that even though O2 is reduced on a clean Pt surface the Pt-oxide layer is generated under a small O2 reduction current.28 The existence of the Pt-oxide layer can prevent the Pt from dissolution.15 According to Figure 7, O2 reduction causes the generation of H2O2. In a polymer electrolyte fuel cell, the O2 supplied to

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the cathode makes a cross-leak through the polymer electrolyte membrane, which generates H2O2 at the anode during the oxidation of hydrogen.29 Because H2O2 passes through the Nafion membrane, the H2O2 generated at the anode can attack the cathode in the fuel cell. At the cathode, the Pt-based electrocatalyst is surrounded by H2O2. In the presence of H2O2, the following reactions are considered to occur thermodynamically.

PtO + 2H2O2+2H++2e- f Pt + 3H2O + O2 (E0 ) 2.074 V vs RHE) (3) PtO + 2H2O2+5H++5e- f Pt + 4H2O + OH (E0 ) 2.492 V vs RHE) (4) PtO + H2O2+4H++4e- f Pt + 3H2O (E0 ) 2.756 V vs RHE) (5) Once reaction 2 occurs, the Pt-oxide is reduced by reactions 3, 4, and 5. According to this consideration, it is expected that the resulting amount of Pt-oxide is decreased by the H2O2-based reactions. As a result of the diminished amount of Pt-oxide, the prevention of Pt dissolution is ineffective. The Pt dissolution may take place anodically.2,15 In the future, we plan to investigate the detailed dissolution mechanism. 4. Conclusions In the present study, we have investigated the Pt degradation in the layered structure of a Pt/C catalyst/Nafion membrane system focusing on the cathode electrocatalyst of PEFCs. The obtained results are summarized as follows. (1) Regarding O2 supplied to the WE in the solid-state cell, the ESA of the Pt/C markedly decreases during an alternating potential cycle between (i) a positive-potential region of 0.20.65 V vs Ag/Ag2SO4 and (ii) a negative-potential region of -0.73 to -0.4 V vs Ag/Ag2SO4. (2) The decrease in ESA leads to the deposition of large Pt particles in the Nafion membrane, which causes the Pt dissolution. (3) To clarify the role of the negative-potential region (-0.73 to -0.4 V vs Ag/Ag2SO4), EQCM measurement was conducted in an acidic solution, leading to the result on the basis of the assumption that H2O2 generated during O2 reduction induces the Pt dissolution. (4) When H2O2 is supplied to the solid-state cell, Pt dissolution occurs by an alternating potential cycle between 0.40.65 V vs Ag/Ag2SO4 and -0.2-0.15 V vs Ag/Ag2SO4. (5) It is considered that the mechanism for the Pt degradation in the solid state is the generation of H2O2, which can reduce the Pt-oxide that protects the Pt from dissolution.

Acknowledgment. The present work was supported by the New Energy and Industrial Technology Department Organization (NEDO). References and Notes (1) Yu, X.; Ye, S. J. Power Sources 2007, 172, 145. (2) Kim, L.; Chung, C. G.; Sung, Y. W.; Chung, J. S. J. Power Sources 2008, 183, 524. (3) Xie, J.; Wood, D. L., III.; More, K. L.; Atanassov, P.; Borup, R. L. J. Electrochem. Soc. 2005, 152, A1011. (4) Kinoshita, K.; Lundquist, J. T.; Stonehart, P. J. Electroanal. Chem. 1973, 48, 157. (5) Ferreira, P. J.; la O’, G. J.; Shao-Horn, Y.; Morgan, D.; Makharia, R.; Kocha, S.; Gasteiger, H. A. J. Electrochem. Soc. 2005, 152, A2256. (6) Akita, T.; Taniguchi, A.; Maekawa, J.; Shiroma, Z.; Tanaka, K.; Kohyama, M.; Yasuda, K. J. Power Sources 2006, 159, 461. (7) Bi, W.; Gray, G. E.; Fuller, T. F. Electrochem. Solid-State Lett. 2007, 10, B101. (8) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions, 2nd ed.; National Association of Corrosion Engineers: Houston, TX, 1974; p 378. (9) Bard, A. J. Encyclopedia of Electrochemistry of the Elements; Marcel Dekker: New York, 1976; Vol. 6. (10) Kanzaki, Y.; Takahashi, M. J. Electroanal. Chem. 1978, 90, 305. (11) Benke, G.; Gnot, W. Hydrometallurgy 2002, 64, 205. (12) Sugawara, Y.; Yadav, A. P.; Nishikata, A.; Tsuru, T. Electrochemistry 2007, 75, 359. (13) Ota, K.; Nishigori, S.; Kamiya, N. J. Electroanal. Chem. 1988, 257, 205. (14) Rand, D. A. J.; Woods, R. J. Electroanal. Chem. 1972, 35, 209. (15) Mitsushima, S.; Kawahara, S.; Ota, K.; Kamiya, N. J. Electrochem. Soc. 2007, 154, B153. (16) Birss, V. I.; Chang, M.; Segal, J. J. Electroanal. Chem. 1993, 355, 181. (17) Kodera, F.; Kuwahara, Y.; Nakazawa, A.; Umeda, M. J. Power Sources 2007, 172, 698. (18) Parthasarathy, A.; Martin, C. R. J. Electrochem. Soc. 1991, 138, 916. (19) Basura, V.; Beattie, P. D.; Holdcroft, S. J. Electroanal. Chem. 1998, 458, 1. (20) Jiang, J.; Kucernak, A. J. Electroanal. Chem. 2005, 576, 223. (21) Umeda, M.; Mohamedi, M.; Uchida, I. Langmuir 2001, 17, 7970. (22) Bard, A. J.; Faulker, L. R. Electrochemical Methods, 2nd ed.; Wiley: New York, 2001; p 166. (23) Mayrhofer, K. J. J.; Strmcnik, D.; Blizanac, B. B.; Stamenkovic, V.; Arenz, M.; Markovic, N. M. Electrochim. Acta 2008, 53, 3181. (24) Kuzume, A.; Herrero, E.; Feliu, J. M. J. Electroanal. Chem. 2007, 599, 333. (25) Inaba, M.; Yamada, H.; Tokunaga, J.; Tasaka, A. Electrochem. Solid-State Lett. 2004, 7, A474. (26) Shimazu, K.; Kita, H. J. Electroanal. Chem. 1992, 341, 361. (27) Rodrı´guez, J. L.; Pastor, E.; Zinola, C. F.; Schmidt, V. M. J. Appl. Electrochem. 2006, 36, 1271. (28) Appleby, A. J.; Foulkes, F. R. Fuel Cell Handbook; Van Norstrand Reinhold: New York, 1989; Chapter 12. (29) Aoki, M.; Uchida, H.; Watanabe, M. Electrochem. Commun. 2006, 8, 1509.

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