Efficiency of Oxygen Evolution on Iridium Oxide Determined from the

Jul 21, 2017 - The oxygen evolution reaction (OER; 2H2O → O2 + 4H+ + 4e–) is being intensively studied to generate fossil fuel-independent energy ...
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Efficiency of Oxygen Evolution on Iridium Oxide Determined from the pH Dependence of Charge Accumulation Hideshi Ooka,†,‡ Akira Yamaguchi,§ Toshihiro Takashima,∥ Kazuhito Hashimoto,⊥ and Ryuhei Nakamura*,‡,# †

Department of Applied Chemistry, School of Engineering, The University of Tokyo, 7-3-1, Hongo, Bunkyo-ku, Tokyo 113-8656, Japan ‡ Biofunctional Catalyst Research Team, RIKEN Center for Sustainable Resource Science (CSRS), 2-1 Hirosawa, Wako, Saitama 351-0198, Japan § Department of Materials Science and Engineering, School of Materials and Chemical Technology, Tokyo Institute of Technology, 2-12-1, Ookayama, Meguro-ku, Tokyo 152-0033, Japan ∥ Clean Energy Research Center, University of Yamanashi, 4-3-11, Takeda, Kofu, Yamanashi 400-8511, Japan ⊥ National Institute for Materials Science (NIMS), 1-2-1, Sengen, Tsukuba, Ibaraki 305-0047, Japan # Earth-Life Science Institute, Tokyo Institute of Technology, 2-12-1, Ookayama, Meguro-ku, Tokyo 152-0033, Japan S Supporting Information *

ABSTRACT: The oxygen evolution reaction (OER; 2H2O → O2 + 4H+ + 4e−) is being intensively studied to generate fossil fuel-independent energy carriers. As 4d/5d rare metal catalysts, such as amorphous iridium oxide (IrOx), display higher activity than 3d metal catalysts, elucidating the critical mechanistic differences between these materials is important for the synthesis of cost-effective OER catalysts. Although most studies of OER catalysts have focused on O−O bond formation energetics, here, we examined the OER mechanism of IrOx based on charge accumulation, which was recently shown to determine the OER activity for Mn and Fe oxides. Kinetic analysis using Tafel and trumpet plots, along with the difference in the pH dependence between the OER onset potential and that of iridium valence change, showed that the valence change of iridium is more favorable than O−O bond formation. In situ evanescent wave spectroscopy revealed that an intermediate assignable to Ir5+ with oxygen ligands in opposite spin serves as the precursor of OER regardless of pH. As the generation of this species is not related to valence changes of iridium, these results confirm that charge accumulation is not rate-limiting for OER on IrOx, which is a key mechanistic difference between IrOx and less-efficient 3d metal electrocatalysts.

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INTRODUCTION The oxygen evolution reaction (OER, 2H2O → O2 + 4H+ + 4e−) has received considerable attention in recent years1−4 as a promising method for utilizing water as a sustainable electron source. By coupling the OER with a reduction half-reaction such as hydrogen evolution (2H+ + 2e− → H2)5,6 or CO2 reduction (CO2 + 8H+ + 8e− → CH4 + 2H2O),7−11 the intermittent energy from renewable sources, including solar energy or wind power, can be captured in the form of chemical fuels.12 Artificially driving the four-electron transfer process of water to dioxygen has proven difficult, however, due to the large overpotential and corrosive environment at the anode. Among artificial catalyst materials examined to date, RuO213−15 has the smallest overpotential (200 mV) in acidic conditions but also exhibits poor stability. One approach for increasing the stability of these materials is to construct mixed oxides using iridium oxide.16,17 Although amorphous iridium oxide (IrOx) has a slightly larger overpotential than RuO2,2,18,19 IrOx is more robust under anodic conditions and maintains activity across a wide pH range.2,20 However, the scarcity and high costs © 2017 American Chemical Society

associated with 4d/5d metals limit the large-scale applicability of these materials, and therefore it is highly desirable to substitute the rare metal catalysts with those derived from earth-abundant 3d metals. In contrast to 4d/5d metals, the majority of 3d metal oxides, such as MnO2,21,22 NiO2,4 Co3O4,23 and Fe2O3,24,25 operate most efficiently under alkaline conditions and with a few exceptions1,26−28 have markedly reduced catalytic properties at acidic and neutral pH. Not only will this impose limitations for the use of 3d metal oxides in acidic polymer electrolyte membrane (PEM) electrolyzers, which is considered to be one of the most efficient electrolysis systems,29−31 the inefficiency at lower pH is a fundamental restriction for anode materials because the pH in the vicinity of the anode becomes increasingly acidic after prolonged electrolysis,32,33 even in the presence of a buffered electrolyte. To facilitate the rational design of 3d metal catalysts for practical applications, it is Received: April 21, 2017 Revised: June 18, 2017 Published: July 21, 2017 17873

DOI: 10.1021/acs.jpcc.7b03749 J. Phys. Chem. C 2017, 121, 17873−17881

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iridium, these results indicate that charge accumulation proceeds efficiently on IrOx.

necessary to understand the key properties of 4d/5d metal catalysts that allow these rare materials to operate over a wide pH range. Catalytic properties are often evaluated within the framework of the Sabatier principle, which states that an optimum catalyst must bind the substrate with intermediate strength.34 Such binding affinity would allow sufficient interaction between the substrate and catalyst but would not hinder the desorption process, which is necessary to prevent poisoning of the catalyst surface. In recent years, density functional theory calculations35−40 have allowed binding affinities to be calculated, thereby greatly enhancing the understanding of the factors that determine the overpotential of a given catalyst. Despite these advances, it remains unclear why IrOx has greater OER activity than MnO2 at neutral pH. For example, under the assumption that O−O bond formation proceeds through nucleophilic attack of water onto an oxo species, Man et al.36 predicted that rutile IrO2 and MnO2 have similar overpotentials. Similar activity between MnO2 and IrO2 was also predicted by Busch et al.,41 who calculated the OER overpotential assuming a direct coupling pathway for O−O bond formation. The fact that two independent studies36,41 predicted that these materials should have similar activities, despite the experimentally observed overpotential difference of approximately 500 mV being greater than typical DFT errors,40 indicates that key aspects of catalysis remain unresolved. The discrepancy between the theoretical and experimental overpotential values may be attributable to the emphasis of the Sabatier principle on the bond formation process. In the case of multielectron transfer catalysis, multiple charges must be stored prior to the reorganization of chemical bonds, such as O−O bond formation in the case of OER. Oxides and similar catalysts with localized electronic structures perform this charge accumulation process through a valence state change of the metal center.21,22,24,25 Such charge accumulation has been confirmed in photosystem II,42,43 a biological enzyme in which the stepwise oxidation of Mn3+ to Mn4+ allows accumulation of the four charges necessary for water oxidation. Notably, the charge accumulation process may even dictate the overall OER activity, such as in the case of MnO221,22 and Fe2O324,25 OER catalysts at neutral pH, where we have attributed the large OER overpotential to the generation of a metal center in the highspin d4 configuration (Mn3+ and Fe4+). An OER mechanism where valence change is rate-limiting is in stark contrast to that of IrOx in acidic conditions, where we proposed the ratelimiting step to be due to O−O bond formation based on in situ spectroscopy and DFT calculation.44 From this viewpoint, both the bond formation and charge accumulation processes should be evaluated to identify the limiting factors of OER catalysts. As IrOx and MnO2 were predicted to have similar energetics for chemical bond reorganization and interaction,36,41 we hypothesized that the difference in 3d and 5d metal catalysts may be due to their charge accumulation efficiency. Herein, we investigated the OER mechanism of IrOx to gain insight into the critical factors between 5d metal catalysts and the more abundant, but less active, 3d metal catalysts, particularly with respect to the efficiency of charge accumulation. Using in situ evanescent wave (EW) spectroscopy, an intermediate assignable to Ir5+ with oxygen ligands in opposite spin was identified as the precursor of OER across the entire pH range investigated (pH 2−12), suggesting that O−O bond formation proceeds through a binuclear mechanism. As the generation of Ir5+ was not related to the valence change of

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EXPERIMENTAL DETAILS Fabrication of IrOx Electrodes. IrOx electrodes were fabricated by anodic deposition of [Ir(OH)6]2− as previously reported.2,44 Briefly, 0.1 mmol of K2IrCl6 (48.3 mg) was added to 50 mL of NaOH solution (pH 12.1) to yield an opaque brown solution, which was then heated at a rate of 3 °C/min to 70 °C under constant stirring using a hot plate held at 80 °C. The solution was removed from the hot plate as necessary to maintain the rate of heating. The solution pH and temperature values during the heating process are listed in Table S1. Once reaching 70 °C, the solution was cooled immediately in an ice bath. During cooling to 16 °C, the pH of the solution increased from 8.81 to 9.5. The cooled iridium stock solution was stored at 4 °C in the dark. Prior to electrodeposition, 1 mL of the stock solution was diluted 10-fold in 3 M NaCl solution and was allowed to warm to room temperature. The electrodeposition of IrOx on tin-doped-In2O3 (ITO) and fluorinedoped tin oxide (FTO) electrodes was performed at a constant current density of 25 μA cm−2 for 60 s. In Situ EW Spectroscopy. In situ EW spectroscopy was conducted using a quartz optical waveguide covered with a thin surface layer of ITO. An IrOx film was deposited on the ITO layer as described above, and the waveguide was mounted into the sample chamber of a surface and interface spectrometer system (SIS5000, System Instruments). A custom-made, threeelectrode electrochemical cell was placed on top of the electrode. The optical waveguide located at the bottom of the electrochemical reactor acted as an internal multiple reflection element and generated an evanescent wave at the interface between the electrode and electrolyte solution. The incident light from a 150 W xenon lamp was directed at the reactor at an angle of 3° parallel to the surface of the ITO substrate. Electrochemical Analysis. The electrolyte was prepared by adding the components of Britton−Robinson buffer (0.4 M each of phosphate, borate, and acetate) to a 0.5 M Na2SO4 solution, and the pH was then adjusted to the desired value by addition of H2SO4 or NaOH. All glassware was sonicated in ultrapure water directly before performing electrochemical measurements. A platinum wire and an Ag/AgCl/KCl (saturated) electrode were used as the counter and reference electrodes, respectively, for all measurements. All electrochemical potentials were converted into the standard and reversible hydrogen scales (SHE and RHE, respectively) after conducting the measurements. An automatic polarization system (HZ-5000, Hokuto-Denko) was used as the potentiostat. Trumpet Plot Analysis. The rate constants of the electrochemical processes E1 and E2 were estimated from trumpet plots using a Microsoft Excel spreadsheet. The slopes va and vc at faster scan rates were evaluated using the data from 10 to 100 V/s for E1 and the data from 10 to 50 V/s for E2. The raw voltammograms are presented in Figures S1−S6.

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RESULTS AND DISCUSSION Redox Potentials and Kinetics of Charge Accumulation for IrOx. Cyclic voltammograms (CVs) of IrOx in a buffered 0.5 M Na2SO4 electrolyte were obtained at various pH values (Figure 1). Consistent with previous reports,2,19,44,45 the prepared electrodes showed typical redox behavior of hydrous 17874

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the observed pH dependence indicates that electrons and protons transfer with a noninteger ratio of approximately 1:1.5. As each redox wave corresponds to a single electron transfer, such a noninteger ratio could be due to the combined effects of multiple processes, such as the simultaneous contribution of redox changes from hydrous and anhydrous iridium oxide as previously suggested.46 In contrast, the onset potential of OER, which was defined as the potential at which the current density exceeded 10 μA cm−2, remained relatively constant between pH 2 and 12 (Figure 2, black squares), indicating that the ratelimiting step of OER on IrOx is a process involving an equal amount of protons and electrons. Although the Nernstian behavior (0 mV/pH with respect to the RHE scale) of the onset potential of OER was previously documented,2,45 the observed difference in pH dependence between redox transitions and the OER onset potential, along with the presence of distinct voltammetric peaks (E1 and E2) before the onset of OER, suggest that redox transitions of the iridium center are not rate-limiting for OER on IrOx. To confirm that the valence state change of iridium does not kinetically inhibit the OER, the standard rate constant k0 (rate constant at zero driving force) of E1, E2, and OER was evaluated using trumpet and Tafel plots. In the trumpet plots of E1 (Figure 3) and E2 (Figure 4), the differences in the peak potentials during the anodic and cathodic CV scans markedly increased at scan rates faster than 10 V/s at all examined pH values, indicating a transition from a fully reversible to irreversible redox conversion. This will be the basis for

Figure 1. Cyclic voltammograms of the IrOx film in a buffered 0.5 M Na2SO4 electrolyte. Each voltammogram was offset by 5 μA cm−2, and the reported potentials are referenced vs the standard hydrogen electrode (SHE) for clarity. Scan rate: 10 mV/s.

IrOx films. Specifically, the redox transition from Ir3+ to Ir4+ (E1) was clearly detectable between pH 2 and 12, whereas that of Ir4+ to Ir5+ (E2) was more obvious at alkaline pH. In addition, E1 and E2 shifted with a pH dependence of approximately 30 mV/pH on the reversible hydrogen electrode (RHE) scale (Figure 2, red and blue squares), which is consistent with the report by Steegstra et al.45 According to the Nernst equation,

Figure 2. pH dependence of the onset of OER current (black), E1 (redox transition between Ir3+ and Ir4+, red), and E2 (redox transition between Ir4+ and Ir5+, blue). The onset of OER current was defined as the potential at which the OER current exceeded 10 μA cm−2, and the values for E1 and E2 were calculated from the average of anodic and cathodic sweeps during the CVs.

Figure 3. Peak potential of E1 with respect to the logarithm of the scan rate [V/s] during the CVs at pH 2 (A), 4 (B), 6 (C), 8 (D), 10 (E), and 12 (F). The raw cyclic voltammograms are presented in the Supporting Information (Figures S1−S6). Black: anodic peak; red: cathodic peak. 17875

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Table 1. Trumpet Plot of E1 at pH 2, 4, 6, 8, 10, and 12 pH

va[V/log(V/s)]a

vc[V/log(V/s)]b

αc

k0 [1/s]d

2 4 6 8 10 12

0.203 0.237 0.207 0.263 0.215 0.238

0.235 0.276 0.267 0.213 0.192 0.192

0.464 0.462 0.438 0.552 0.528 0.553

4.21 4.94 4.50 4.55 3.93 4.11

a

Slope of the anodic peak in the trumpet plot (Figure 3). bSlope of the cathodic peak in the trumpet plot (Figure 3). cElectron transfer coefficient. dStandard rate constant of the redox transition between Ir3+ and Ir4+.

Table 2. Trumpet Plot of E2 at pH 2, 4, 6, 8, 10, and 12 pH

va[V/log (V/s)]a

vc[V/log (V/s)]b

αc

k0 [1/s]d

2 4 6 8 10 12

0.124 0.140 0.198 0.157 0.122 0.357

0.153 0.077 0.111 0.158 0.114 0.109

0.447 0.653 0.639 0.498 0.517 0.766

2.65 1.92 2.76 3.05 2.27 3.24

a

Slope of the anodic peak in the trumpet plot (Figure 4). bSlope of the cathodic peak in the trumpet plot (Figure 4). cElectron transfer coefficient. dStandard rate constant of the redox transition between Ir4+ and Ir5+.

Figure 5 shows the Tafel plots used to estimate the rate constant of the OER catalyzed by IrOx. The lack of hysteresis in

Figure 4. Peak potential of E2 with respect to the logarithm of the scan rate [V/s] during the CVs at pH 2 (A), 4 (B), 6 (C), 8 (D), 10 (E), and 12 (F). The raw cyclic voltammograms are presented in the Supporting Information (Figures S1−S6). Black: anodic peak; red: cathodic peak.

calculating the standard rate constant, based on Laviron’s theory of nondiffusion systems.47 In the case of a diffusionless electrochemical system, the rate constant k0 can be estimated from the equation k0 =

αnFvc (1 − α)nFva = RT RT

(1)

where R is the gas constant, T represents the absolute temperature, α is the electron transfer coefficient, n represents the number of electrons transferred, and va and vc are the anodic and cathodic slopes of the trumpet plot, respectively. The accuracy of eq 1 increases when the scan rate is sufficiently fast relative to the electron transfer rate, and the overpotential is sufficiently large such that the current from the reverse reaction can be neglected. An experimental benchmark for the validity of eq 1 is for the peak separation in the CVs to be larger than 200 mV,47 which is satisfied at scan rates >10 V/s. Therefore, we can obtain α and k0 based on the slopes of the trumpet plots (va and vc) based on a rearranged form of eq 1: v α = a 1−α vc (2)

Figure 5. Tafel plots at pH 2, 4, 6, 8, 10, and 12 constructed based on the current−potential characteristics of IrOx obtained during forward CV scans.

the CVs (Figure 1) shows that the OER current is controlled not by mass transport but by electron transfer kinetics. The Tafel slope, exchange current density, iridium coverage, and standard rate constants determined from the Tafel plots are summarized in Table 3. The iridium coverage was estimated by the charge transferred under the redox wave E1, based on the assumption that all Ir sites which undergo the redox transition between Ir3+ and Ir4+ are OER-active. At all examined pH values, the standard rate constant for the OER was several orders of magnitude smaller than that of E1 and E2, indicating that the redox transition of iridium is kinetically faster than that of OER on IrOx. This conclusion holds true even if the number of active sites changes 2-fold depending on whether OER proceeds through a mononuclear or binuclear mechanism. Although the exchange current density estimated in this work was several orders of magnitude smaller than the value reported by Zhao et al. ((4−8) × 10−10 A cm−2),2 the iridium coverage

The results of the trumpet plot analysis are presented in Tables 1 and 2. The standard rate constants for both redox transitions (E1 and E2) were in the order of 1/s. The kinetics of the redox transitions in the context of OER catalysis can be evaluated by comparison with the standard rate constant of OER, which was determined from the Tafel plot. 17876

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approximately 450 nm (hereafter denoted as species A450) was observed at all examined pH values, suggesting that the pH-independent onset potential of OER on IrOx (constant at approximately 1.5 V vs RHE) is due to the existence of a single intermediate that initiates the OER in a pH-independent manner. The absorption maxima of 450 nm coincides with the absorption maxima of Ir5+ reported by Castillo-Blum et al. (447 nm),48 indicating that species A450 is an Ir5+ related species. This Ir5+ is in a different chemical state from the Ir5+ generated during the redox process E2, as can be seen from the difference in absorption maxima. In each panel of Figure 6, the absorption maxima at 1.2 V, corresponding to the Ir5+ generated at E2, is approximately 410 nm. However, scanning the potential anodically to 1.5 V resulted in a 40 nm red-shift. We have previously performed DFT calculations and have rationalized this red-shift based on a spin transition of an oxygen atom ligated to Ir5+. This transition shifts the spin of two adjacent oxygen ligands from symmetric to asymmetric spin, allowing O−O bond formation to occur spontaneously. Having a ratelimiting O−O bond formation process via the binuclear mechanism is also consistent with the Tafel slope of 30 mV/ dec (Table 3). Therefore, the spin transition corresponding to the red-shift of the EW spectra is expected to be the elementary step which dictates OER activity on IrOx. Further support for the difference in the chemical origin of the two spectral features was provided by the difference in the pH dependence of their formation (Figure 7). The pH dependence of species A450 formation (red circles) exhibited a close relationship with the onset of OER current (black squares) and remained constant at approximately 1.4 V vs RHE. However, the potential necessary for oxidizing Ir4+ to Ir5+ (absorption maximum 410 nm, green circles) shifted with a pH

Table 3. Exchange Current Density and Rate Constants of the OER Current at pH 2, 4, 6, 8, 10, and 12 for IrOx pH 2 4 6 8 10 12

Tafel slope [mV/dec] 29 28 28 23 24 23

j0 [μA/cm2]a 1.6 1.2 5.0 4.7 2.6 3.4

× × × × × ×

−7

10 10−7 10−8 10−10 10−10 10−11

Γ [μC/cm2]b 15.6 16.4 17.6 9.39 7.78 10.6

k0 [1/s]c 2.6 1.9 7.1 1.3 8.6 8.1

× × × × × ×

10−9 10−9 10−10 10−11 10−12 10−13

a

Exchange current density obtained by extrapolating the Tafel plot shown in Figure 5 to 1.23 V vs RHE. bCoverage of iridium calculated from the charge of E1. All electrochemically active iridium sites were assumed to be participate during OER. cRate constant at equilibrium potential obtained using the equation j0 = nk0Γ.

on the electrode was also several orders of magnitude smaller than their reported values (4 mC cm−2). The numbers reported by Zhao et al.2 yield standard rate constants in the order of 10−8 [1/s], which is also orders of magnitude smaller than those of E1 and E2. Therefore, by combining the kinetic analysis with the thermodynamic information on E1 and E2 (Figure 2), we confirmed that the valence change of IrOx is more favorable than the rate-limiting step of OER. Analysis of the Rate-Limiting Step of OER Using in Situ EW Spectroscopy. To gain direct insight into the factors influencing the rate-limiting step of OER on IrOx, in situ EW spectroscopy was performed for IrOx between pH 2 and 12. Figure 6 shows the potential dependence of the EW subtraction spectra when the potential was scanned in the positive direction from 1.2 to 1.5 V vs RHE under acidic, neutral, and alkaline pH conditions. A species with an absorption maximum at

Figure 6. Potential dependence of the EW spectrum for IrOx between 1.2 and 1.5 V at pH 2 (A), 4 (B), 6 (C), 8 (D), 10 (E), and 12 (F). The spectrum at 1.2 V was used as the reference spectrum in all cases. Spectra are shown in 20 mV intervals, and the spectra at 1.5 and 1.4 V are colored red and blue, respectively. 17877

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photoelectrochemical systems.52,53 This property of MnO2 and Fe2O3 is due to the instability of the high-spin d4 configuration to charge disproportionation (2d4 → d3 + d5), which arises from the instability of having only one electron in the degenerate eg orbital. This instability provides the driving force for electron exchange between two nearby metal ions and would lead to dissipation of the accumulated oxidization energy, resulting in an inefficient charge accumulation process. In contrast, IrOx exhibits markedly different properties because of its low-spin electron configuration. A catalytic cycle consisting of Ir3+, Ir4+, and Ir5+ valence states corresponds to an OER mechanism with low-spin d6, d5, and d4 electronic states.54,55 The absence of eg electrons from the catalytic cycle preemptively suppresses charge disproportionation, allowing for a smooth charge accumulation process (Figure 8). While we Figure 7. pH dependence of the onset of OER current (black), absorption at 450 nm (red), 410 nm, (green), and 580 nm (blue). The absorption at 410 and 580 nm correspond to E2 and E1, respectively (Figure S7). The onset of OER current was defined as the potential at which the OER current exceeded 10 μA cm−2 and is the same as in Figure 2. The onset potential for the absorption at 450 nm was defined as the potential where the absorption at 450 nm in Figure 6 exceeded 0.02. The onset potential for the absorptions at 410 and 580 nm were defined as the potential where the absorption at the specified wavelength increased by more than 0.005 compared to the spectrum taken at a potential 100 mV more negative. The difference in definition of the onset potential with the absorption at 450 nm is because the pH dependence of E1 and E2 did not allow a single potential to serve as a reference spectrum for all pH.

Figure 8. Comparison of MnO2 (A) and IrOx (B) catalysts based on their electron configuration and susceptibility to charge disproportionation.

dependence of 30 mV/pH on the RHE scale, in accordance with the pH dependence of E2 (Figure 2). The pH dependence of redox potentials reflects the extent of coupling between proton transfer and electron transfer. Therefore, the formation of species A450 is due to a different process than the valence change of Ir4+ to Ir5+ (redox process E2). These results indicate that the rate-determining step of the OER on IrOx is not the valence change of iridium, but the activation of the oxygen ligand. This may proceed by a spin transition mechanism, which we have proposed previously based on DFT calculations.44 The present assignment of the spin state transition is also supported by a recent DFT study by Ping et al.,49 who reported that a spin transition occurs within the adsorbed oxygen atom upon O−O bond formation. There is also experimental evidence based on in situ X-ray absorption spectroscopy (XAS)50 where the oxidation state of oxygen was found to change from O2− to O− at OER potentials. Although the authors have interpreted this change in the context of the mononuclear mechanism for OER based on the electrophilicity of the O− site, their results nonetheless indicate the active participation of the ligated oxygen atom in the rate-limiting step of OER on IrOx. The detection of a pH-independent OER precursor that is not generated through valence change indicates that charge accumulation on IrOx is more efficient than O−O bond formation. This behavior is in stark contrast to the OER mechanism of Mn21,22 and Fe oxides.24,25 Implications of an Efficient Charge Accumulation Process. The efficiency of the charge accumulation process may be influenced by a lack of unstable valence states within the catalytic cycle of IrOx. For MnO2 and Fe2O3 catalysts, the formation of the catalytic intermediates in the high-spin d4 electronic state (Mn3+ and Fe4+, respectively) was found to dictate OER activity in both electrochemical21,22,24,25,51 and

do not expect bulk properties to directly influence surface phenomena such as catalysis, the preferential localization of electrons of MnO2 and Fe2O3 (charge disproportionation) seems to be consistent with their lower electronic conductivity.56 In addition to OER, 4d and 5d metal catalysts are more active than their 3d metal counterparts for other multielectron transfer reactions in general. This property of 4d and 5d metal catalysts may be because their low-spin electronic configuration promotes charge accumulation, which is necessary prior to bond-formation and -breaking processes. It should also be noted here that efficient charge accumulation also implies favorable reaction kinetics. The Tafel slope of IrOx was found to be approximately 30 mV/dec in the present study. Although this value is lower than the 40 mV/dec reported by Zhao et al.,2 it is nonetheless markedly smaller than the value of 120 mV/dec reported for MnO2.28,57 The Tafel slope is inversely proportional to the effective transfer coefficient (αeff) of the total reaction: Tafel slope =

∂E RT ln 10 = ∂ log |j| αeff F

(3)

where R, T, and F correspond to the gas constant, temperature, and Faraday constant, respectively. A higher effective transfer coefficient will allow for the generation of a larger current at a given overpotential, and therefore a smaller Tafel slope is a desirable property of an active catalyst material. This property may be achieved when the concentration of the precursor for the rate-limiting step is potential dependent, as occurs for hydrogen evolution on platinum in acidic electrolytes.58 The implications of a favorable charge accumulation process on the 17878

DOI: 10.1021/acs.jpcc.7b03749 J. Phys. Chem. C 2017, 121, 17873−17881

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The Journal of Physical Chemistry C effective transfer coefficient αeff and Tafel slopes can be illustrated mathematically. When [Crds] denotes the coverage of the intermediate responsible for the rate-limiting step, the current (j) can be expressed as

0.5 and r = 1, again giving a smaller Tafel slope (40 mV/dec) due to the potential dependence of [Crds]. Therefore, a catalyst with facile charge accumulation steps preceding a rate-limiting step is expected to show enhanced reaction kinetics through an increase in the effective transfer coefficient αeff. This speculation is supported by the present experimental observations for IrOx, as a mechanism with the rate-limiting step (species A450 formation) being preceded by a process (Ir5+ generation, k ∼ 1/s) with a rate constant of approximately 10 orders of magnitude larger indeed yielded a small Tafel slope. A reaction mechanism with the rate-limiting step at later steps of the catalytic cycle is in stark contrast to the mechanism of MnO2, where the rate-limiting step (Mn2+ →Mn3+ + e−) is the first electron transfer event, as evidenced by the Tafel slope of 120 mV/dec. Thus, the difference in the Tafel slope and onset potential between IrOx and MnO2 further highlights the importance of charge accumulation for efficient OER, not only from the viewpoint of lowering the overpotential requirement but also with respect to enhancing the electrode kinetics.

⎛ ⎛ α F (E − E ) ⎞ OER j = nFk[Crds]r ⎜exp⎜ rds ⎟ ⎠ RT ⎝ ⎝ ⎛ (1 − αrds)F(E − EOER ) ⎞⎞ − exp⎜ − ⎟⎟ ⎝ ⎠⎠ RT

(4)

or when the reverse reaction is negligible ⎛ α F(E − EOER ) ⎞ j = nFk[Crds]r exp⎜ rds ⎟ ⎠ ⎝ RT

(5)

In eqs 4 and 5, αrds is the transfer coefficient of the rate-limiting elementary reaction and r is the reaction order with respect to [Crds]. If [Crds] is not a function of the electrochemical potential, then αeff = αrds. Assuming αrds ≈ 0.5, the Tafel slope would be 120 mV/dec, indicating that the current increases 10fold for every 120 mV of overpotential. However, it is wellknown that the Tafel slope (larger αeff) is smaller when the rate-limiting intermediate is involved in an electrochemical preequilibrium step.58−60 For example, if a one-electron equilibrium exists between the rate-limiting intermediate Crds and its precursor Cpre (Cpre → Crds + e−, equilibrium potential Eeq), [Crds] can be expressed using the Nernst equation: [C ] RT ln rds F [Cpre]

(6)

⎛ F(E − Eeq ) ⎞ ⎟⎟ ∴ [Crds] = [Cpre] exp⎜⎜ RT ⎝ ⎠

(7)

E = Eeq −

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CONCLUSIONS We investigated the OER mechanism of IrOx by assessing charge accumulation to elucidate the critical mechanistic differences between IrOx and less-efficient 3d metal OER catalysts, such as MnO2 and Fe2O3. The charge accumulation process, which occurs through valence state changes of the metal center, was found to proceed efficiently on IrOx, as the redox peaks of Ir3+ to Ir4+ (E1) and Ir4+ to Ir5+ (E2) were much more negative than the onset potential of the OER. In addition to favorable thermodynamics, the kinetics of charge accumulation were also more efficient compared to that of O−O bond formation based on the rate constants estimated from the trumpet and Tafel plots of IrOx. To obtain direct evidence for the rate-limiting factor of OER on IrO x, in situ EW spectroscopy was used to observe the electronic state change of IrOx. This analysis led to the identification of a common intermediate species, designated species A450, which appeared concurrently with the onset of OER current regardless of electrolyte pH. This species was previously assigned as two oxygen atoms with opposite spins that are ligated to two adjacent Ir5+ ions. As the pH dependence of species A450 formation clearly differs from that of Ir5+ generation (E2), the valence change of iridium is not rate-limiting for the OER. Taken together, these results show that in contrast to 3d metal catalysts, such as MnO2 or Fe2O3, charge accumulation on IrOx is a highly efficient process. The mechanistic differences between 3d and 5d metal catalysts revealed by studying the pH dependence of charge accumulation IrOx is expected to aid in the design of active catalysts for electrolyzers and artificial photosynthetic systems that operate under different pH environments.

Combining eqs 5 and 7 gives the following current expression using [Cpre]: ⎛ αrdsF(E − EOER ) + rF(E − Eeq ) ⎞ ⎟⎟ j = nFk[Cpre]r exp⎜⎜ RT ⎝ ⎠ (8)

⎛ (r + αrds)FE − αrdsFEOER − rFEeq ) ⎞ ⎟⎟ = nFk[Cpre]r exp⎜⎜ RT ⎝ ⎠ (9)

Based on this equation, αeff would be increased to (r + αrds) and the Tafel slope would become smaller. For example, if the coverage of adsorbed oxygen (Oads) was in equilibrium with its precursor, and the rate-limiting step of OER was the rearrangement of chemical bonds with two adjacent Oads (binuclear or direct coupling mechanism), then αeff would become 2 based on αrds = 0 and r = 2. This would yield a Tafel slope of 30 mV/dec, such as that observed in this study. In contrast to the case without a pre-equilibrium step, only 30 mV is necessary to generate a 10-fold current enhancement. The enhanced sensitivity of the reaction rate with respect to the electrode potential upon introducing a potential dependence in [Crds] holds true qualitatively, regardless of the reaction mechanism. For example, different O−O bond formation mechanisms, such as that through nucleophilic attack of water to adsorbed oxygen, would yield an αeff of 1.5, based on αrds =

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.7b03749. Time dependence of the solution pH during alkaline hydrolysis of K2IrCl6, CVs used for making trumpet plots, and the pH dependence of absorbance derived from Ir4+ (580 nm) and Ir5+ (410 nm) (PDF) 17879

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The Journal of Physical Chemistry C

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Study of Nanoparticles and Bulk Materials. ACS Catal. 2012, 2, 1765− 1772. (15) Lee, Y.; Suntivich, J.; May, K. J.; Perry, E. E.; Shao-Horn, Y. Synthesis and Activities of Rutile IrO2 and RuO2 Nanoparticles for Oxygen Evolution in Acid and Alkaline Solutions. J. Phys. Chem. Lett. 2012, 3, 399−404. (16) Yeo, R. S.; Orehotsky, J.; Visscher, W.; Srinivasan, S. Ruthenium-Based Mixed Oxides as Electrocatalysts for Oxygen Evolution in Acid Electrolytes. J. Electrochem. Soc. 1981, 128, 1900− 1904. (17) Sardar, K.; Petrucco, E.; Hiley, C. I.; Sharman, J. D. B.; Wells, P. P.; Russell, A. E.; Kashtiban, R. J.; Sloan, J.; Walton, R. I. WaterSplitting Electrocatalysis in Acid Conditions Using Ruthenate-Iridate Pyrochlores. Angew. Chem., Int. Ed. 2014, 53, 10960−10964. (18) Seitz, L. C.; Dickens, C. F.; Nishio, K.; Hikita, Y.; Montoya, J.; Doyle, A.; Kirk, C.; Vojvodic, A.; Hwang, H. Y.; Norskov, J. K.; et al. A Highly Active and Stable IrOx/SrIrO3 Catalyst for the Oxygen Evolution Reaction. Science 2016, 353, 1011−1014. (19) Geiger, S.; Kasian, O.; Shrestha, B. R.; Mingers, A. M.; Mayrhofer, K. J. J.; Cherevko, S. Activity and Stability of Electrochemically and Thermally Treated Iridium for the Oxygen Evolution Reaction. J. Electrochem. Soc. 2016, 163, F3132−F3138. (20) Yagi, M.; Tomita, E.; Kuwabara, T. J. Remarkably High Activity of Electrodeposited IrO2 Film for Electrocatalytic Water Oxidation. Electroanal. Chem. 2005, 579, 83−88. (21) Takashima, T.; Hashimoto, K.; Nakamura, R. Mechanisms of pH-Dependent Activity for Water Oxidation to Molecular Oxygen by MnO2 Electrocatalysts. J. Am. Chem. Soc. 2012, 134, 1519−1527. (22) Yamaguchi, A.; Inuzuka, R.; Takashima, T.; Hayashi, T.; Hashimoto, K.; Nakamura, R. Regulating Proton-Coupled Electron Transfer for Efficient Water Splitting by Manganese Oxides at Neutral pH. Nat. Commun. 2014, 5, 4256. (23) Xu, Q.-Z.; Su, Y.-Z.; Wu, H.; Cheng, H.; Guo, Y.-P.; Li, N.; Liu, Z.-Q. Effect of Morphology of Co3O4 for Oxygen Evolution Reaction in Alkaline Water Electrolysis. Curr. Nanosci. 2014, 11, 107−112. (24) Takashima, T.; Ishikawa, K.; Irie, H. Efficient Oxygen Evolution on Hematite at Neutral pH Enabled by Proton-Coupled Electron Transfer. Chem. Commun. 2016, 52, 14015−14018. (25) Takashima, T.; Ishikawa, K.; Irie, H. Detection of Intermediate Species in Oxygen Evolution on Hematite Electrodes Using Spectroelectrochemical Measurements. J. Phys. Chem. C 2016, 120, 24827−24834. (26) Esswein, A. J.; Surendranath, Y.; Reece, S. Y.; Nocera, D. G. Highly Active Cobalt Phosphate and Borate Based Oxygen Evolving Catalysts Operating in Neutral and Natural Waters. Energy Environ. Sci. 2011, 4, 499−504. (27) Huynh, M.; Bediako, D. K.; Liu, Y.; Nocera, D. G. Nucleation and Growth Mechanisms of an Electrodeposited Manganese Oxide Oxygen Evolution Catalyst. J. Phys. Chem. C 2014, 118, 17142−17152. (28) Huynh, M.; Bediako, D. K.; Nocera, D. G. A Functionally Stable Manganese Oxide Oxygen Evolution Catalyst in Acid. J. Am. Chem. Soc. 2014, 136, 6002−6010. (29) Carmo, M.; Fritz, D. L.; Mergel, J.; Stolten, D. A Comprehensive Review on PEM Water Electrolysis. Int. J. Hydrogen Energy 2013, 38, 4901−4934. (30) Puthiyapura, V. K.; Pasupathi, S.; Su, H.; Liu, X.; Pollet, B.; Scott, K. Investigation of Supported IrO2 as Electrocatalyst for the Oxygen Evolution Reaction in Proton Exchange Membrane Water Electrolyser. Int. J. Hydrogen Energy 2014, 39, 1905−1913. (31) Nakamura, A.; Ota, Y.; Koike, K.; Hidaka, Y.; Nishioka, K.; Sugiyama, M.; Fujii, K. A 24.4% Solar to Hydrogen Energy Conversion Efficiency by Combining Concentrated Photovoltaic Modules and Electrochemical Cells. Appl. Phys. Express 2015, 8, 107101. (32) Auinger, M.; Katsounaros, I.; Meier, J. C.; Klemm, S. O.; Biedermann, P. U.; Topalov, A. A.; Rohwerder, M.; Mayrhofer, K. J. J. Near-Surface Ion Distribution and Buffer Effects During Electrochemical Reactions. Phys. Chem. Chem. Phys. 2011, 13, 16384−16394. (33) Katsounaros, I.; Meier, J. C.; Klemm, S. O.; Topalov, A. A.; Biedermann, P. U.; Auinger, M.; Mayrhofer, K. J. J. The Effective

AUTHOR INFORMATION

Corresponding Author

*E-mail [email protected]; Tel +81-(0)48-467-9539; Fax +81-(0)48-462-4639 (R.N.). ORCID

Ryuhei Nakamura: 0000-0003-0743-8534 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors thank Dr. Hiromi Takahashi of System Instruments for technical support with the in situ EW spectroscopy measurements. This work was supported by JSPS Grant-in-Aid for Scientific Research no. 26288092. H.O. acknowledges Grant-in-Aid for JSPS Research Fellows no. 15J10450.

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REFERENCES

(1) Kanan, M. W.; Nocera, D. G. In Situ Formation of an OxygenEvolving catalyst in Neutral Water Containing Phosphate and Co2+. Science 2008, 321, 1072−1075. (2) Zhao, Y.; Vargas-Barbosa, N. M.; Hernandez-Pagan, E. A.; Mallouk, T. E. Anodic Deposition of Colloidal Iridium Oxide Thin Films from Hexahydroxyiridate (IV) Solutions. Small 2011, 7, 2087− 2093. (3) Morris, N. D.; Suzuki, M.; Mallouk, T. E. Kinetics of Electron Transfer and Oxygen Evolution in the Reaction of [Ru(bpy)3]3+ with Colloidal Iridium Oxide. J. Phys. Chem. A 2004, 108, 9115−9119. (4) Trzesniewski, B. J.; Diaz-Morales, O.; Vermaas, D. A.; Longo, A.; Bras, W.; Koper, M. T. M.; Smith, W. A. In Situ Observation of Active Oxygen Species in Fe-Containing Ni-Based Oxygen Evolution Catalysts: The Effect of pH on Electrochemical Activity. J. Am. Chem. Soc. 2015, 137, 15112−15121. (5) Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Hydrogen Oxidation and Evolution Reaction Kinetics on Platinum: Acid vs Alkaline Electrolytes. J. Electrochem. Soc. 2010, 157, B1529−B1536. (6) Greeley, J.; Jaramillo, T. F.; Bonde, J.; Chorkendorff, I. B.; Norskov, J. K. Computational High-Throuput Screening of Electrocatalytic Materials for Hydrogen Evolution. Nat. Mater. 2006, 5, 909− 913. (7) Hori, Y.; Kikuchi, K.; Suzuki, S. Production of CO and CH4 in Electrochemical Reduction of CO2 at Metal Electrodes in Aqueous Hydrogen Carbonate Solution. Chem. Lett. 1985, 14, 1695−1698. (8) Hori, Y.; Kikuchi, K.; Murata, A.; Suzuki, S. Production of Methane and Ethylene in Electrochemical Reduction of Carbon Dioxide at Copper Electrode in Aqueous Hydrogen Carbonate Solution. Chem. Lett. 1986, 15, 897−898. (9) Schouten, K. J. P.; Gallent, E. P.; Koper, M. T. M. The Influence of pH on the Reduction of CO and CO2 to Hydrocarbons on Copper Electrodes. J. Electroanal. Chem. 2014, 716, 53−57. (10) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Norskov, J. K. How Copper Catalyzes the Electroreduction of Carbon Dioxide into Hydrocarbon Fuels. Energy Environ. Sci. 2010, 3, 1311− 1315. (11) Tang, W.; Peterson, A. A.; Varela, A. S.; Jovanov, Z. P.; Bech, L.; Durand, W. J.; Dahl, S.; Norskov, J. K.; Chorkendorff, I. The Importance of Surface Morphology in Controlling the Selectivity of Polycrystalline Copper for CO2 Electroreduction. Phys. Chem. Chem. Phys. 2012, 14, 76−81. (12) Tributsch, H. Energy Materials: What Strategies did Nature Pursue? Phys. Status Solidi A 2014, 211, 2052−2062. (13) Burke, L. D.; Murphy, O. J.; O’Neill, J. F.; Venkatesan, S. J. The Oxygen Electrode. Part 8.-Oxygen Evolution at Ruthenium Dioxide Anodes. J. Chem. Soc., Faraday Trans. 1 1977, 73, 1659−1671. (14) Reier, T.; Oezaslan, M.; Strasser, P. Electrocatalytic Oxygen Evolution Reaction (OER) on Ru, Ir, and Pt Catalysts: A Comparative 17880

DOI: 10.1021/acs.jpcc.7b03749 J. Phys. Chem. C 2017, 121, 17873−17881

Article

The Journal of Physical Chemistry C Surface pH During Reactions at the Solid-Liquid Interface. Electrochem. Commun. 2011, 13, 634−637. (34) Sabatier, P. Hydrogenations et Deshydrogenations par Catalyse. Ber. Dtsch. Chem. Ges. 1911, 44, 1984−2001. (35) Norskov, J. K.; Studt, F.; Abild-Pedersen, F.; Bligaard, T. Fundamental Concepts in Heterogeneous Catalysis; John-Wiley and Sons, Inc.: Hoboken, NJ, 2014. (36) Man, I. C.; Su, H. Y.; Calle-Vallejo, F.; Hansen, H. A.; Martinez, J. I.; Inoglu, N. G.; Kitchin, J.; Jaramillo, T. F.; Norskov, J. K.; Rossmeisl, J. Universality in Oxygen Evolution Electrocatalysis on Oxide Surfaces. ChemCatChem 2011, 3, 1159−1165. (37) Bajdich, M.; Garcia-Mota, M.; Vojvodic, A.; Norskov, J. K.; Bell, A. T. Theoretical Investigation of the Activity of Cobalt Oxides for the Electrochemical Oxidation of Water. J. Am. Chem. Soc. 2013, 135, 13521−13530. (38) Calle-Vallejo, F.; Diaz-Morales, O. A.; Kolb, M. J.; Koper, M. T. M. Why is Bulk Thermochemistry a Good Descriptor ofr the Electrocatalytic Activity of Transition Metal Oxides? ACS Catal. 2015, 5, 869−873. (39) Li, H.; Li, Y.; Koper, M. T. M.; Calle-Vallejo, F. Bond-Making and Breaking Between Carbon, Nitrogen, and Oxygen in Electrocatalysis. J. Am. Chem. Soc. 2014, 136, 15694−15701. (40) Xu, Z.; Rossmeisl, J.; Kitchin, J. R. A Linear Response DFT + U Study of Trends in the Oxygen Evolution Activity of Transition Metal Rutile Dioxides. J. Phys. Chem. C 2015, 119, 4827−4833. (41) Busch, M.; Ahlberg, E.; Panas, I. Water Oxidation on MnOx and IrOx: Why Similar Performance? J. Phys. Chem. C 2013, 117, 288−292. (42) Yano, J.; Yachandra, V. Mn4Ca Cluster in Photosynthesis: Where and How Water is Oxidized to Dioxygen. Chem. Rev. 2014, 114, 4175−4205. (43) Glockner, C.; Kern, J.; Broser, M.; Zouni, A.; Yachandra, V.; Yano, J. Structural Changes of the Oxygen-Evolving Complex in Photosystem II During the Catalytic Cycle. J. Biol. Chem. 2013, 288, 22607−22620. (44) Ooka, H.; Wang, Y.; Yamaguchi, A.; Hatakeyama, M.; Nakamura, S.; Hashimoto, K.; Nakamura, R. Legitimate Intermediates of Oxygen Evolution on Iridium Oxide Revealed by In Situ Electrochemical Evanescent Wave Spectroscopy. Phys. Chem. Chem. Phys. 2016, 18, 15199−15204. (45) Steegstra, P.; Busch, M.; Panas, I.; Ahlberg, E. Revisiting the Redox Properties of Hydrous Iridium Oxide Films in the Context of Oxygen Evolution. J. Phys. Chem. C 2013, 117, 20975−20981. (46) Steegstra, P.; Ahlberg, E. Influence of Oxidation State on the pH Dependence of Hydrous Iridium Oxide Films. Electrochim. Acta 2012, 76, 26−33. (47) Laviron, E. General Expression of the Linear Potential Sweep Voltammogram in the Case of Diffusionless Electrochemical Systems. J. Electroanal. Chem. Interfacial Electrochem. 1979, 101, 19−28. (48) Castillo-Blum, S. E.; Richens, D. T.; Sykes, A. G. Oxidation of Hexaaquairidium (III) and Related Studies: Preparation and Properties of Iridium (III), Iridium (IV), and Iridium (V) Dimers as Aqua Ions. Inorg. Chem. 1989, 28, 954−960. (49) Ping, Y.; Nielsen, R. J.; Goddard, W. A. The Reaction Mechanism with Free Energy Barriers at Constant Potentials for the Oxygen Evolution Reaction at the IrO2 (110) Surface. J. Am. Chem. Soc. 2017, 139, 149−155. (50) Pfeifer, V.; Jones, T. E.; Velasco-Velez, J. J.; Arrigo, R.; Piccinin, S.; Havecker, M.; Knop-Gericke, A.; Schlogl, R. In Situ Observation of Reactive Oxygen Species Forming on Oxygen-Evolving Iridium Surfaces. Chem. Sci. 2017, 8, 2143−3149. (51) Zandi, O.; Hamann, T. W. Determination of Photoelectrochemical Water Oxidation Intermediates on haematite Electrode Surfaces Using Operando Infrared Spectroscopy. Nat. Chem. 2016, 8, 778. (52) Le Formal, F.; Pastor, E.; Tilley, S. D.; Mesa, C. A.; Pendlebury, S. R.; Grätzel, M.; Durrant, J. R. Rate Law Analysis of Water Oxidation on a Hematite Surface. J. Am. Chem. Soc. 2015, 137, 6629−6637. (53) Klahr, B.; Hamann, T. Water Oxidation on Hematite Photoelectrodes: Insight into the Nature of Surface States Through

In-Situ Spectroelectrochemistry. J. Phys. Chem. C 2014, 118, 10393− 10399. (54) Komer, W. D.; Machin, D. J. Ternary and Quaternary Oxides of Ruthenium and Iridium. J. Less-Common Met. 1978, 61, 91−105. (55) Kuriyama, H.; Matsuno, J.; Niitaka, S.; Uchida, M.; Hashizume, D.; Nakao, A.; Sugimoto, K.; Ohsumi, H.; Takata, M.; Takagi, H. Epitaxially Stabilized Iridium Spinel Oxide Without Cations in the Tetrahedral Site. Appl. Phys. Lett. 2010, 96, 182103. (56) Attfield, J. P.; Bell, A. M. T.; Rodriguez-Martinez, L. M.; Greneche, J. M.; Cernik, R. J.; Clarke, J. F.; Perkins, D. A. Electrostatically driven charge-ordering in Fe2OBO3. Nature 1998, 396, 655−658. (57) Hayashi, T.; Yamaguchi, A.; Hashimoto, K.; Nakamura, R. Stability of Organic Compounds on the Oxygen-Evolving Center of Photosystem II and Manganese Oxide Water Oxidation Catalysts. Chem. Commun. 2016, 52, 13760−13763. (58) Kucernak, A. R.; Zalitis, C. General Models for the Electrochemical Hydrogen Oxidation and Hydrogen Evolution Reactions: Theoretical Derivation and Experimental Results Under Near Mass-Transport Free Conditions. J. Phys. Chem. C 2016, 120, 10721−10745. (59) Conway, B. E.; Tilak, B. V. Interfacial Processes Involving Electrocatalytic Evolution and Oxidation of H2 and the Role of Chemisorbed H. Electrochim. Acta 2002, 47, 3571−3594. (60) Koper, M. T. M. Analysis of Electrocatalytic Reaction Schemes: Distinction Between Rate-Determining and Potential-Determining Steps. J. Solid State Electrochem. 2013, 17, 339−344.

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