Microkinetic assessment of electrocatalytic oxygen evolution reaction over iridium oxide in unbuffered conditions

Journal of Catalysis 391 (2020) 435–445

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Microkinetic assessment of electrocatalytic oxygen evolution reaction over iridium oxide in unbuffered conditions Takeshi Nishimoto, Tatsuya Shinagawa, Takahiro Naito, Kazuhiro Takanabe ⇑ Department of Chemical System Engineering, School of Engineering, The University of Tokyo, 7-3-1 Hongo, Bunkyo-ku, Tokyo, Japan

a r t i c l e

i n f o

Article history: Received 4 May 2020 Revised 25 August 2020 Accepted 8 September 2020 Available online 17 September 2020 Keywords: Electrocatalysis Oxygen evolution reaction Iridium oxide Near-neutral pH Reaction mechanism

a b s t r a c t Water electrolysis driven by electrical power generated from renewable energy sources will play a pivotal role in future sustainable societies, which requires adaptation of various reaction conditions as well as electrolyte identities. Regardless, the anodic half-reaction of the oxygen evolution reaction (OER) is considered a kinetic bottleneck. This study provides quantitative description of the OER kinetics based on rigorous microkinetic analyses including Tafel analysis, isotope effects and temperature dependence using an IrOx electrocatalyst in unbuffered solution at varying pH levels. The diffusional constraints of H+/OH
determine three distinctive kinetic regimes in the pH-potential-current relationships: below pH 5, between pH 5 and 10, and above pH 10 at appreciable current densities on the order of 1 mA cm
2. When shifting from alkaline to acidic solution, the complete consumption of local OH
near the electrode surface switches the OER proceeding as the oxidation of OH
to that of the water molecule at pH ~ 11 irrespective of the electrode identity. At pH 5–10, the diffusional constraints of H+ generated via oxidation reaction yield an environment with pH ~ 4 near the electrode surface even prior to the OER, resulting in a bulk pH-independent region for the OER performance. Under this unbuffered near-neutralpH condition, the isotope effect was diminished for the OER catalysis, which is consistent with the ratedetermining step (rds) being the sole electron-transfer step via the formation of O-O bonds, decoupled from proton transfer. This reaction mechanism is distinct from that under more acidic conditions (pH < 4), although the water molecule is the same reactant. Under acidic conditions, noticeable isotope effects were observable, which is consistent with the formation of O-O bonds being the rds on uncoordinated bare Ir sites as the most abundant surface species. This study provides a quantitative description of the reactant- and mechanistic-switching that points to concurrent optimization of both electrode materials and electrolyte for improved OER performance at near-neutral pH levels. Ó 2020 The Author(s). Published by Elsevier Inc. This is an open access article under the CC BY-NC-ND license (http://creativecommons.org/licenses/by-nc-nd/4.0/).

1. Introduction By utilizing the energy generated in renewable manners, our society can transition to a sustainable one in the future. In particular, the conversion of the thermodynamically stable chemicals, namely water and carbon dioxide, into energy carriers and valueadded products could play a pivotal role when driven by renewable electricity [1,2]. Accordingly, recent decades have witnessed significant research efforts dedicated to the development of devices to achieve such chemical reactions [3–5]. The oxygen evolution reaction (OER) is the anodic half-reaction commonly shared in most cases, which for instance is coupled with the hydrogen evolution reaction (HER) for water splitting and with the CO2 reduction reaction for CO2 valorization. Critically, the slow kinetics of the OER ⇑ Corresponding author. E-mail address: [email protected] (K. Takanabe).

causes significant loss in the energy efficiency in most systems [6–10], thus necessitating the development of an active electrocatalyst for this reaction. Past research efforts have elucidated the property-activity relationship of the OER. A variety of materials have been examined as OER catalysts, including noble metal oxides [11–13] as well as earth-abundant transition metal oxides and (oxy)hydroxides [8,14,15], which show a well-known ‘‘volcano” trend in the performance as a function of material properties such as the heat of oxide formation [16], the number of electrons in the outer orbital [17,18], the d-band center [19], and DG(*O)
DG(*OH) [7,20,21], among others. In particular, Rossmeisl and coworkers [7,22] used density functional theory (DFT) calculations to identify such a volcano trend in silico by quantifying the free energies of reaction intermediates, i.e., OH, O, and OOH species, which have linear relations with each other as the Koper’s group first reported [23]. The volcano-trend relationship is regarded as one example of the

https://doi.org/10.1016/j.jcat.2020.09.007 0021-9517/Ó 2020 The Author(s). Published by Elsevier Inc. This is an open access article under the CC BY-NC-ND license (http://creativecommons.org/licenses/by-nc-nd/4.0/).

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Journal of Catalysis 391 (2020) 435–445

the reactant induces a change in the reaction mechanism and thus the transition state, leading to alteration of the reaction performance. Nevertheless, there have been few articles comprehensively studying the OER in a wide pH range with a focus on the nearneutral pH medium [48–50]. Herein, we performed a rigorous microkinetic analysis of the OER using IrOx electrocatalyst as a model electrode in unbuffered conditions at varying pH levels. The unbuffered condition at near-neutral pH is employed in direct sea-water splitting, as well as photocatalysis, e.g., overall water splitting reaction, where a pure water is uniquely employed as the reaction medium [51]. In addition, the study in the absence of buffer ions forms a basis for understanding the buffering effects on OER kinetics, which may depend on the identity and activity of the buffer ions. We systematically assessed the impact of the electrolyte pH on the OER, which allowed us to divide the OER into three kinetic regimes: below pH 5, pH 5–10, and above pH 10. In each kinetic regime, we performed a detailed microkinetic analysis by digesting the potential-pH and potential-current relationships as well as the isotope effect. The investigation identifies not only the point of reactant switching from the hydroxide ion to the water molecule due to diffusional restriction of proton and hydroxide ion, but also the point of mechanistic switching between coupled and decoupled proton-electron transfers depending on the pH.

Sabatier principle [7,24], which predicts that the highest performance can be achieved by using a material with surface-oxygen interaction that is neither too strong nor too weak. Among others, the iridium oxide (IrOx) has been reported to exhibit the highest activity toward the OER in both acidic and alkaline pH conditions [13,25–27], which sits at the top of the experimentally obtained volcano trend for the OER according to Seitz et al. [28] or close to the top of the theoretical one in the case of IrO2 (1 1 0) surface [7,29], suggesting that IrOx has the optimal binding energy to the key reaction intermediate. Nonetheless, to reach a current density of 10 mA cm
2 corresponding to approximately 10% solar-tohydrogen efficiency in a combined solar water-splitting system [30–32], the IrOx requires a substantial overpotential of 270 mV in 1 M H2SO4 [8], which is considerably larger than that of the cathodic counterpart in water splitting, in which an overpotential of 40 mV is sufficient for hydrogen evolution reaction over a Pt electrode [33]. Recent studies focusing on IrOx have shed light on the reaction mechanism of the OER under such extreme pH conditions. Employing in-situ near ambient pressure X-ray photoelectron spectroscopy (NAP-XPS), Sanchez Casalongue et al. [34] experimentally observed a potential-dependent change in the valence state of Ir, which eventually becomes an oxidation state of V that drives O-O bond formation at the rate-determining step (rds). This observation is consistent with a report by the group of Mayrhofer, which revealed that the OER proceeds via IrIII-IrV transition by studying the current–potential relationship while concurrently measuring the dissolution of Ir using online electrochemical mass spectrometry (OLEMS) [12]. It is interesting to note that the OER activity over the IrOx has a tradeoff relationship with the stability; Danilovic et al. reported that an Ir site with a valence state of > +IV is active toward the OER, but it can form higher oxidation states during the catalytic cycle, resulting in the dissolution of Ir and a loss in overall performance [35]. This consideration was reinforced by the direct measurement of Ir dissolution reported by Kasian et al. [12], who indicated that the Ir dissolves mostly via IrV-IrVI transformation at highly anodic potentials [12]. To suppress this undesired dissolution of the Ir active phase, Seitz et al. developed an IrOx/SrIrOx catalyst via leaching of Sr from SrIrOx in situ, which required an overpotential of merely 270 to 290 mV to achieve a current density of 10 mA cm
2 during continuous testing for 30 h in 0.5 M H2SO4 [28]. Notably, the experimentally derived reaction mechanism agrees well with DFT calculations simulating IrO2 (1 1 0) surface, which predicted the surface *OOH species formation from an *O surface (i.e., O-O bond formation) to be the potential determining step for the single-site mechanism [22,29,36,37]. These experimental and theoretical studies were successful in elucidating the reaction mechanism in the common reaction conditions of acidic and alkaline pH levels. Near-neutral pH medium has recently emerged as a mild reaction condition compatible with renewable energy-driven devices [5,38,39], such as photovoltaic cells conjugated with electrolyzers [40–43], as well as photoelectrochemical [44,45] and photocatalytic water splitting [46,47]. In this context, studying the OER at various pH levels is essential for understanding and designing an active catalyst. Considering the two possible ways of describing the OER in an equation, namely, the oxidation of the water molecule (Equation (1)) and the oxidation of the hydroxide ion (Equation (2)):

6H2 O O2 þ 4H3 Oþ þ 4e


ð1Þ

4OH
O2 þ 2H2 O þ 4e


ð2Þ

2. Experimental All chemicals used in this study were purchased with the following purities from Sigma-Aldrich: HClO4 (assay 70%), NaClO4xH2O (assay 99.99%), NaOH (99.99%), Na3IrCl6xH2O, H2C2O4 (99%), Na2CO3H2O (99.5%), KCl (99.999%), D2O (99 atom% D), DClO4 (68 wt% in D2O, 99 atom% D), and NaOD (30 wt% in D2O, 99 atom% D). We fabricated IrOx electrodes by electrochemical deposition following a reported protocol [52] with a titanium disk electrode as the substrate. The deposition bath contained 4 mM of Na3IrCl6xH2O and 20 mM of H2C2O4, and the pH level was adjusted to 10 by adding Na2CO3H2O, which was kept at 308 K for 4 days prior to the deposition. On a separate note, the Ti disk electrode had a diameter of 3.0 mm (BAS, Inc.) and was first polished with 1 lm diamond and then with 0.05 lm alumina (both purchased from BAS, Inc.). The electrochemical deposition was conducted using a three-electrode configuration with Ir wire and Hg/Hg2Cl2 (saturated with KCl) as the counter and reference electrodes, respectively. IrOx was deposited onto the titanium disk electrode as a working electrode immersed in the prepared deposition bath by applying a constant current density of 140 lA cm
2 for 10 ks (see Figure S1a for the representative potential profile obtained during the deposition). Before using it for catalytic testing, the fabricated electrode was assessed by cyclic voltammetry (CV) in 0.1 mol kg
1 HClO4 to ensure the identical properties of the electrodes (Figure S1b). A variety of unbuffered aqueous solutions were employed as the electrolyte. In a typical procedure for the preparation of 500 g of 0.1 mol kg
1 aqueous NaClO4 solution at pH greater than 6, 6.12 g of NaClO4xH2O was firstly put in approximately 400 g of ultrapure water (18.2 MX cm) with vigorous stirring. At this stage, the pH value of the solution was adjusted to the target value by adding 0.1 mol kg
1 NaOH while concurrently measuring the pH using a pH meter (D-71 and 9625, HORIBA). After pouring the ultrapure water into the solution until the total weight reached 500 g, the pH level of the thus prepared solution was measured, which is referenced throughout the manuscript. For the preparation of the solution at pH below 7, the electrolyte was prepared in the same manner as the alkaline ones except that 0.1 mol kg
1

one already sees a significant impact of the electrolyte pH, which can trigger the switching of the reactant for the OER between the hydroxide ion and the water molecule. Importantly, the change in 436

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around 3 mA cm
2 at approximately 0.9 V vs. SHE, which started increasing further at a potential of ca. 1.2 V vs. SHE. In alkaline solutions of pH 12 and 13, only monotonically increasing anodic current densities were again observed with the onset potential shifted by
59 mV pH
1, similar to the acidic cases. The onset potential (defined as the potential reaching + 1 mA cm
2) is plotted as a function of pH in Fig. 1b (see Figure S3 for potentialpH diagram at + 10 mA cm
2 and Figure S4 for the variation of the Tafel slope as a function of pH). This figure clearly shows that the onset potential for the OER follows the Nernstian expectation of
59 mV pH
1 at pH levels below 5 and above 10, while it remains constant at pH values in between. Thus, three pH regimes were categorized for the OER: (1) below pH 5, (2) pH levels between 5 and 10, and (3) alkaline pH values of pH 11 and 13. We address the origin of the transition between alkaline and near-neutral conditions by assessing the current–potential relationship at the threshold pH of 11. Figure S5 shows dependence of the current density on the disk-rotation speed in the RDE configuration over the IrOx electrode at pH 11. The plateau of the current densities increased at faster disk-rotation speeds independently of the applied potential, which is consistent with the diffusion limitation of OH
reactant. Applying further overpotentials (Fig. 1a and S5) resulted in a further increase in the current densities exceeding the OH
-diffusion-limited current densities, which indicates the reactant switching to the water molecule in Equation (1). Since it is the property of not the electrode but the electrolyte that induces this reactant-switching between the hydroxide ion and the water molecule, the transition pH of 11 is anticipated to be universal irrespectively of the identity of electrode. In fact, the same pH value of 11 was observed as the transition pH of the j-E relationship over nickel- and cobalt-based electrodes [48], as well as the gold RDE (see Figure S6a for LSVs and Figure S6b for the potential-pH diagram). Taken together, it is concluded that the reactant for the OER switches between the water molecule and hydroxide ion at approximately pH 11 because of diffusional constraint of OH
that cannot catch up with its consumption at appreciable current densities, which causes an apparent transition when compiling the OER as a function of pH. Fig. 1c provides a close look at the redox peaks of IrOx in CVs curves at various pH levels, in which multiple oxidation and reduction peaks appeared in the investigated pH range. The redox peaks prior to the OER catalysis suggested the formation of Ir(V) irrespective of pH, which were previously ascribed by Ooka et al. to the transition of the iridium oxidation state from III to IV and from IV to V at lower and higher potentials, respectively, by in-situ evanescent wave (EW) spectroscopy [27]. These iridium redox processes involve proton-coupled processes; therefore, the apparent positions of the redox peaks remained identical in the nearneutral pH range of 4–10, due to the shift in the local pH to ~ 4 and ~ 11 during the oxidation and reduction events, respectively. Consistently, a drastic change of the local pH in the solutions of bulk near-neutral pH (pH 4–10) was reported by Katsounaros et al., who discussed that a reaction rate of merely 25 mA cm
2 was sufficient to trigger a significant shift of the local pH in the absence of buffering species [53]. Therefore, it is concluded that these small currents resulting from redox events were sufficient to shift the local pH near the electrode surfaces, even prior to the OER catalysis. At higher current densities, the proton generated during the OER contributes to the electrocatalytic performance though its own diffusion, i.e., the local pH gradient, under unbuffered nearneutral pH conditions. Fig. 2 presents the LSV profiles at pH 1 and 7 over the IrOx RDE at varying disk-rotation speeds. While the j-E relationship obtained at pH 1 was insensitive to the diskrotation speed, the anodic current density at pH 7 increased with the disk-rotation speed, suggesting that there is the

HClO4 was used to adjust the pH values while maintaining the total
1 concentration of ClO
. For studies with a 4 anion at 0.1 mol kg deuterium system, D2O, DClO4, and NaOD were used instead of H2O, HClO4, and NaOH, respectively. The pD of the solution was adjusted to the target value by measuring the pH value based on the relationship that pD is equal to the pH value + 0.4. Electrocatalytic measurements were conducted using the IrOx/ titanium electrodes (denoted as IrOx hereafter for simplicity) and gold disk electrodes in a rotating-disk electrode (RDE) configuration and in the three-electrode system with Ir wire and Hg/Hg2Cl2 (saturated with KCl) or Hg/HgO (1.0 M NaOH) electrodes as the counter and reference electrodes, respectively. Notably, Hg/Hg2Cl2 (saturated with KCl in D2O) was used as a reference electrode for the study using D2O. Before and during all measurements, O2 gas (99.99995%) was continuously supplied to the cell. CV, linear sweep voltammetry (LSV), chronopotentiometry (CP), and potentiostatic electrochemical impedance spectroscopy (PEIS) were performed using a 16-channel research-grade potentiostat system (VMP3, BioLogic Science Instruments). All LSV profiles reported in this study were obtained by the cathodic scan at a negative scan rate to isolate the contribution from oxidation of iridium species to the overall electric current. These measurements were conducted at various temperatures of 298, 313, 328 and 343 K, and electrochemical cells employed at latter three temperatures were equipped with a water jacket (Water-Jacket glass cell, BAS Inc.) connected to an external temperature-control system (NCB1210A, EYELA). For measurements at varying temperatures, the reference electrode was connected to the electrochemical cell through an extension jacket to maintain the temperature of the reference electrode at 298 K, and the potential in the present study is denoted with respect to the standard hydrogen electrode (SHE) scale at 298 K. Overpotential for the OER was calculated based on the equilibrium potential for the OER at each temperature, e.g. 1.23 V vs. SHE at 298 K and 1.19 V vs. SHE at 343 K at pH 0. The error bars reported are the standard deviation and the difference of the data values obtained from three and two independently performed experiments, respectively. All current densities are expressed in terms of the geometric electrode surface area unless otherwise noted, and all current–potential relationships were iRcorrected using impedance values measured at a frequency  100 kHz with an amplitude of 10 mV.

3. Results and discussion 3.1. Identification of kinetic regimes for the OER at different pH levels. The IrOx electrode was selected for investigation as it was tolerant to various pH levels ranging from extremely acidic to alkaline conditions. Fig. 1a shows the CV profiles on the SHE scale in the RDE configuration at a fixed rotation speed of 4900 rpm and pH levels from 1 to 13 (see Figure S2 for raw voltammograms). While anodic current densities ascribable to the OER were observed with increasing overpotentials at all pH levels of 1–13, the j-E relationship clearly differed according to the bulk pH values of the electrolyte. At acidic pH levels below 5, the anodic current density monotonically increased with the applied potential, which accompanied a shift of the onset potential with the pH value by approximately
59 mV per pH in accordance with the Nernst equation (see Supplementary Information for thermodynamic considerations). At pH levels of 5–10, while the current densities increased with the potential similarly to the acidic ones, the onset potential remained unchanged, and in turn, the j-E relationship apparently overlapped. Notably, under weakly alkaline conditions of pH 11, a two-step anodic event was observed; with increasing overpotential, the current density first increased until reaching a plateau 437

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Fig. 1. Oxygen evolution reaction (OER) and redox properties of IrOx at various pH levels. Cyclic voltammograms (CVs) highlighting (a) OER and (c) redox properties over the IrOx disk electrode in a rotating disk electrode (RDE) configuration are shown, which were recorded in 0.1 mol kg
1 Na-perchlorate electrolyte solutions at pH levels from 1 to 13, a disk-rotation speed of 4900 rpm, a scan rate of 1 mV s
1, and reaction temperature of 298 K. The numbers in the figure indicate the electrolyte pH in the bulk. (b) Onset potential defined as potential reaching + 1 mA cm
2 in the panel (a) was plotted as a function of electrolyte pH in the bulk. The solid line labeled as Eeq indicates the equilibrium potential for the OER. In panels (b) and (c), dashed lines are provided as a visual aid.

for the OER over CoOx in near-neutral pH conditions to highlight the effects of the buffer [54]. The results indicate the necessity of isolating the kinetic current from diffusion artifacts. Thus far, we have discussed how the pH levels of the electrolyte bulk impacts the CVs and OER, which revealed that (1) the OER at the electrolyte bulk pH above 10 proceeds via the oxidation of the hydroxide ion, (2) the OER below pH 5 involves the oxidation of the water molecule, and (3) the OER in the range of pH 5–10 is also dictated by the oxidation of the water molecule, where the generated proton induces the shift in the local pH value to approximately 4. In the next sections, we present the microkinetic analysis on the OER for each case to elucidate the complicated four-electron reaction, with focus placed on mass-transport-free current density region from 1 to 20 mA cm
2. 3.2. Alkaline pH: Hydroxide ion oxidation. In this section, a microkinetic analysis is discussed for the OER on the IrOx electrode at alkaline pH levels. With OH
as a sole reactant, the apparent rate expression is given by an equation below [55,56]:

Fig. 2. Current-potential relationship for water oxidation over IrOx. LSVs over the IrOx disk electrode in the RDE configuration are shown, which was recorded in 0.1 mol kg
1 Na-perchlorate electrolyte solutions at pH 1 and 7, with disk-rotation speeds from 1600 to 4900 rpm, a scan rate of
1 mV s
1, and reaction temperature of 298 K.

j ¼ A0 cm OH
expffEg

ð3Þ

´ is the pre-exponential factor, c
where A OH is the concentration of the hydroxide ion, m is the reaction order with the respect to the hydroxide ion, f denotes F/RT (F is the Faraday constant, R is the gas constant, and T is absolute temperature), and E is the electrode

mass-transport contribution of the proton from the electrode surface to the bulk electrolyte. Similar considerations were reported 438

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from 0.23 to 0.30 V at acidic pH, from 0.24 to 0.32 V at alkaline pH, and from 0.43 to 0.54 V at near-neutral pH). Using the reaction order and the Tafel slope, we below elaborate the reaction mechanism of the oxidation of the hydroxide ion over the IrOx electrode. Using a commonly reported framework of the reaction network [22,37,62–66], we consider the single active site mechanism, which is composed of the following elementary steps:

potential. In addition, our theoretical considerations included the following [57]:




@E @pH




 j



 @log j @E ¼
@pH E @log j pH

@log j @log cHþ

 E



 @log j @log j ¼
¼
@log cOH
E @pH E

ð4Þ

ð5Þ

Notably, these equations contain the terms of (oE/opH)j and (oE/olog j)pH, which can be elucidated by the E-pH diagram and the E-log(j) plot, respectively. Regarding the former, we evaluated the slope of the E-pH diagram in Fig. 1b and obtained a slope value of
64 mV pH
1 in the alkaline regime of pH 11–13, which is consistent with the literature using IrOx and a-Fe2O3 electrodes [27,50]. As for the E-log(j) relationship, the following equation has been widely confirmed experimentally and is known as the Tafel relation [55,58]:

g ¼ blogjjj þ a

 þ OH
OH þ e


ð7Þ

OH þ OH
O þ H2 O þ e


ð8Þ

O þ OH
OOH þ e


ð9Þ

OOH þ OH
 þ O2 þ H2 O þ e


ð10Þ

where * denotes a site on the surface. The spectroscopic observation of *OOH adsorbate over an Au electrode experimentally corroborated this single active site mechanism for the OER [67]. Because the formation and consumption rate of each *OH, *O, and *OOH would differ depending on the electrode potential, one may anticipate the change in the rds with increasing overpotential [68–70]. However, we observed constant Tafel slope values at each pH level in the range of current density from 1 to 20 mA cm
2 shown in Fig. 3. We therefore consider single rds to account for the OER kinetics in this current regime, and in the following address three possible rds, namely the formation of *O in Equation (8), the formation of *OOH in Equation (9), and the decomposition of *OOH in Equation (10) (see [58] and Supplementary Information for the derivation of rate expression). When the formation of *O is the rds, there are only two types of the surface adsorbate expected prior to the rds; * (empty site) or *OH. Based on microkinetic analysis, we obtain rate expression for the oxidation of hydroxide ion, which is limited by the formation of *O on the surface predominantly comprising * and *OH, respectively:

ð6Þ

where g is the overpotential, which is the difference between the electrode and equilibrium potentials, and a and b are constants. The slope b is called the Tafel slope and indicates the sensitivity of the electric current response to the applied potential, which provides insights into the reaction mechanism and the rds as discussed below [58–60]. In Fig. 3c, we present the Tafel plot of the OER in this pH regime (see Figures S2 and S7 for LSVs and Koutecky´-Levich (KL) analysis), revealing that the slope value remained almost identical irrespectively of the bulk pH values and was 40 ± 8 mV dec
1 at pH 13, in reasonable agreement with the reported value of 47.7 mV dec
1 in 1 M NaOH solution, contrasting with that of 23 mV dec
1 obtained in buffered solution at pH 12 [27,61]. The values of (oE/o pH)j =
64 mV pH
1 and (oE/olog j)pH = 40 mV dec
1 resulted in (olog j/opH)E = 1.6 dec pH
1 according to Equations (4), which corresponds to a reaction order with respect to the hydroxide ion of 1.6 in the rate expression for the OER based on Equation (5), i.e., m = 1.6 in Equation (3). Notably, because the Tafel slope in Fig. 3 and the pH dependence in Fig. 1 were obtained in the consistent range of current density from 1 to 20 mA cm
2, our discussion on the kinetics is valid in this current range (the overpotential range

n o 0 j ¼ nFkð8Þ K 0ð7Þ c2OH
exp f gð7Þ þ f ð1
að8Þ Þgð8Þ

ð11Þ

n o 0 j ¼ nFkð8Þ cOH
exp f ð1
að8Þ Þgð8Þ

ð12Þ

Fig. 3. Tafel plots of the OER over IrOx at various pH levels. Electrode potential is compiled as a function of the kinetic current densities over the IrOx disk electrode recorded in 0.1 mol kg
1 Na-perchlorate electrolyte solutions at (a) pH levels from 1 to 3, (b) pH 7, and (c) pH levels from 11 to 13 at a scan rate of
1 mV s
1 and reaction temperature of 298 K. 439

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n o 0 j ¼ nFkð10Þ cOH
exp f ð1
að10Þ Þgð10Þ

where k0 is the standard rate constant, a is the electron transfer coefficient, ci is the concentration of species i, and K0 is the ratio of the standard rate constant for the forward and backward reactions (see our previous microkinetic study for the derivation of the equations [58]). The subscripts indicate the corresponding equation number. These equations predict the Tafel slope values of 30 and 120 mV dec
1 as well as the reaction orders of 2 and 1 toward the hydroxide ion at the full-coverage of * and *OH, respectively, at an electron transfer coefficient of 0.5 [58]. When comparing these theoretical values with the experimental ones obtained herein, we found that the experimental Tafel slope of 40 ± 8 mV dec
1 and reaction order of 1.6 fall between two cases, indicating that the formation of *O on the surface composed of both * and *OH agrees with the experimental observation. This scenario, however, could be discarded based on existing spectroscopic study (vide infra). We then consider the formation of *OOH species being the rds, which can proceed on the surface comprising *O, *OH, and/or *OOH. According to a previous study on Tafel analysis [58], a smaller Tafel slope value is obtainable when the rds appears at a later stage of the reaction sequence while the most abundant surface species forms at an earlier stage of the reaction. More precisely, when the rds is the formation of *OOH (Equation (9)), the theoretical Tafel slope values of 21, 40, and 120 mV dec
1 are expected if the most abundant surface species are *, *OH, and *O, respectively, at an electron transfer coefficient of 0.5. Accordingly, we can formulate the rate expression for each case [58], and the equations below indicate the rate of the hydroxide ion oxidation, which is limited by the formation of *OOH when the most abundant surface species are *, *OH, and *O, respectively:

j¼

j¼

n o 0 nFkð9Þ K 0ð7Þ K 0ð8Þ c3OH
exp f ðgð7Þ þ gð8Þ Þ þ f ð1
að9Þ Þgð9Þ cH2 O n o 0 nFkð9Þ K 0ð8Þ c2OH
exp f gð8Þ þ f ð1
að9Þ Þgð9Þ c H2 O

n o 0 j ¼ nFkð9Þ cOH
exp f ð1
að9Þ Þgð9Þ

According to these equations, we expect the theoretical Tafel slope values of 22, 30, 60, and 120 mV dec
1 with corresponding reaction orders of 4, 3, 2, and 1 toward the hydroxide ion on the surface comprising *, *OH, *O, and *OOH, respectively, at an electron transfer coefficient of 0.5 [58]. Our experimentally obtained value of 40 ± 8 mV dec
1 is consistent with the surface comprising both *OH and *O, which however contradicts with that the reaction order of 1.6 is evidenced by the surface covered by both *O and *OOH. This inconsistency reasonably allows us to exclude this possibility in the present condition. The above discussion suggests the rds of the oxidation of hydroxide ion over IrOx to be (1) the formation of *O on the surface composed of both * and *OH, or (2) the formation of *OOH on the surface predominantly covered by *OH. We remark here that a previous spectroscopic study identified IrV surface species using an IrOx catalyst at pH 12 during the OER [27], and in the framework of single-site mechanism, the elementary step occurring on IrV corresponds to the formation of *OOH in Equation (9). In addition, previous DFT calculations suggested this step to be thermodynamic potential determining step [22] that likely limits the overall reaction rate [71]. The consistent observation herein therefore allow us to pin-down the rds to be the formation of *OOH at the IrV site, demonstrating that the rigorous microkinetic analysis is a powerful tool for understanding the electrocatalytic reaction, which we rely on to elucidate the water oxidation reaction kinetics at acidic and near-neutral pH levels in subsequent sections. 3.3. Acidic pH: water oxidation

ð13Þ

We now turn to the oxidation of the water molecules as the OER occurring at acidic pH level over the IrOx surface. Similar to the previous case on the oxidation of the hydroxide ion at alkaline pH levels, we analyzed the slope in the potential-pH diagram and the Tafel plot. The Tafel plot in Fig. 3a reveals a Tafel slope value of 29 ± 1 mV dec
1 at pH 1 and a corresponding ln(j0) value being
14.2 ± 0.5, which remained almost unchanged with varying pH levels below 3 (see Figures S2 and S8 for LSVs and K-L analysis), in agreement with a value of 29 mV dec
1 reported by Ooka et al. [27] but slightly smaller than the value reported by Abbott et al. (38–45 mV dec
1) [72]. Together with the (oE/opH)j slope of
50 mV pH
1 at j = 1 mA cm
2 in the acidic pH region (pH 1– 3) of Fig. 1b, the reaction order (olog j/ologc+H)E was revealed to be
1.7 according to Equations (4) and (5). At a higher reaction rate of 10 mA cm
2, the (oE/opH)j slope slightly increased to
42 mV pH
1 as shown in Figure S3 that corresponds to the reaction order of
1.4, demonstrating the insensitivity of the reaction order to the reaction rate of our interest, from 1 to 20 mA cm
2. In the commonly adopted reaction network employing a single active site [22,37,62–66], the oxidation of the water molecules is considered to proceed via the following elementary steps:

ð14Þ ð15Þ

Equations (13), (14), and (15) predict the reaction order of the hydroxide ion as 3, 2, and 1, respectively. Comparing the experimental and theoretical values, we found that the experimentally observed Tafel slope of 40 ± 8 mV dec
1 agrees well with the theoretical one of 40 mV dec
1, and the experimentally observed reaction order of 1.6 is also close to the corresponding reaction order of 2. These quantitative agreements suggest that the OER via the oxidation of hydroxide ion on the IrOx surface can likely be limited by the formation of *OOH species (Equation (9)) with the surface site predominantly comprising *OH. Lastly, the case of the *OOH decomposition being the rds will be discussed, which has recently been claimed theoretically be Exner and Over [66]. The equations below dictate the rate of the oxidation of hydroxide ion, which is limited by the decomposition of *OOH on surfaces covered by *, *OH, *O, and *OOH, respectively:

j¼

n o 0 nFkð10Þ K 0ð7Þ K 0ð8Þ K 0ð9Þ c4OH
exp f ðgð7Þ þ gð8Þ þ gð9Þ Þ þ f ð1
að10Þ Þgð10Þ c H2 O ð16Þ

j¼

n o 0 nFkð10Þ K 0ð8Þ K 0ð9Þ c3OH
exp f ðgð8Þ þ gð9Þ Þ þ f ð1
að10Þ Þgð10Þ c H2 O

n o 0 j ¼ nFkð10Þ K 0ð9Þ c2OH
exp f gð9Þ þ f ð1
að10Þ Þgð10Þ

ð19Þ

ð17Þ

 þ H2 O OH þ Hþ þ e


ð20Þ

OH O þ Hþ þ e


ð21Þ

O þ H2 O OOH þ Hþ þ e


ð22Þ

OOH  þ O2 þ Hþ þ e


ð23Þ

In the same manner as the oxidation of the hydroxide ion in the previous section, we below discuss three possible rds of the formation of *O in Equation (21), the formation of *OOH in Equation (22), and the decomposition of *OOH in Equation (23).

ð18Þ 440

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T. Nishimoto et al.

In the case where the formation of *O limits the rate, we can formulate rate equations below for the water molecule oxidation occurring on the surface composed of * and *OH, respectively:

j¼

n o 0 nFkð21Þ K 0ð20Þ cH2 O exp f gð20Þ þ f ð1
að21Þ Þgð21Þ c Hþ

n o 0 j ¼ nFkð21Þ exp f ð1
að21Þ Þgð21Þ

j¼

n o 0 nFkð22Þ K 0ð20Þ K 0ð21Þ c2H2 O exp f ðgð20Þ þ gð21Þ Þ þ f ð1
að22Þ Þgð22Þ c2Hþ ð26Þ

j¼

c Hþ

n o 0 j ¼ nFkð22Þ cH2 O exp f ð1
að22Þ Þgð22Þ

ð27Þ ð28Þ

which show reaction orders of
2,
1, and 0 with respect to the proton, respectively. The derivation of the equations is available in a previous microkinetic study [58]. By comparing the theoretical values with the experimentally obtained ones, we found that the Tafel slope of 29 ± 1 mV dec
1 lies between 21 and 40, and the reaction order of
1.7 falls between
2 and
1 predicted by the orders above. Generally, when transitioning from a surface state (e.g., *) to the other (e.g., *OH), the Tafel slope and reaction order also vary from one value to the other. When two adsorbate coexists, the values of the Tafel slope and the reaction orders become between these two states, as has been numerically demonstrated in our previous study [58]. With this discussion in mind, this quantitative agreement suggests that the oxidation of the water molecule at acidic pH levels is limited by the formation of *OOH species with the most abundant surface species being * and *OH if the single-site mechanism is adopted. When the decomposition of *OOH is the rds, the equations below dictate the rate for the oxidation of water molecule, which is limited by the decomposition of *OOH on the surface covered by *, *OH, *O, and *OOH, respectively (Equation (29)):

j¼

n o 0 nFkð23Þ K 0ð20Þ K 0ð21Þ K 0ð22Þ c2H2 O exp f ðgð20Þ þ gð21Þ þ gð22Þ Þ þ f ð1
að23Þ Þgð23Þ c3Hþ ð29Þ

j¼

ð31Þ

ð32Þ

These equations dictate the theoretical Tafel slope values of 22, 30, 60, and 120 mV dec
1 when the most abundant surface species are *, *OH, *O, and *OOH, respectively, at an electron transfer coefficient of 0.5 [58], and the corresponding reaction orders to the proton are
3,
2,
1, and 0. Comparison of these quantities with the experimentally obtained Tafel slope of 29 ± 1 mV dec
1 and reaction order of
1.7 suggests that, when the decomposition of *OOH is the rds, the surface is predominantly composed of *OH. Our rigorous microkinetic analysis thus far successfully pinned down as the rds of the water molecule oxidation either (1) the formation *OOH on the surface composed of both * and *OH, or (2) the decomposition of *OOH on the surface composed of *OH. While the recent computational study argued the latter to likely be the rds [66], experimental studies conclusively demonstrated the formation of O-O bond is the rds on the IrV species generated at the top surface layer of the IrOx catalyst under OER conditions at acidic pH [27,34,73], which was supported by a DFT study [36]. With these existing studies, we argue that the former is likely to be the case; that is, the water oxidation over IrOx at acidic pH levels is limited by the formation of *OOH on the surface comprising both * and *OH. We further examined the water oxidation reaction in acidic medium by studying the isotope effect. In general, the elementary step for electrocatalysis falls into one of the following categories:, namely the ET, PT, or CPET reactions [57,74]; and in the latter two cases, the reaction rate is sensitive to the isotope present in the reaction media. Fig. 4a shows the kinetic current densities as a function of the electrode potential in both H2O and D 2O systems at pH 1 and pD 1 (see Figures S2a, S7a, and S8 for LSV and K-L plots). In both electrolytes, the current densities increased with the applied anodic potential, although the current density in the H2 O was larger than that in the D 2O in the potential window investigated. In addition, the extent of the increase in the current density with the potential was greater in H2O than D 2O. We note in passing that the OER mechanism is likely identical in both systems because the Tafel slope of 35 mV dec
1 obtained with D2 O was similar to the slope of 29 ± 1 mV dec
1 observed with H2O (see Figure S9a for the Tafel plot). As a result, the ratio of the current density that corresponds to the apparent ‘‘kinetic isotope effect” (KIE) was larger than unity in the potential range investigated over the IrO x electrocatalyst. The apparent KIE increased with increasing overpotential, as shown in Fig. 4a, which was likely due to the difference in the electron transfer coefficient and the overpotential for preequilibrium reaction according to Equation (26). The apparent KIE greater than unity indicates that the rds involves the transfer of proton, i.e., either PT or CPET. In this context, we argue that the possibility of the OER occurring via the dimerization [60] can be discarded. More precisely, although the dimerization of the surface metal-oxo species can agree with the experimentally obtained Tafel slope of 30 mV dec
1 [60], it involves the coupling of *O species without the dissociation of O-H bonds and therefore is expected to be insensitive to the isotope. Taken together, it is concluded that the rate of water oxidation over the IrOx surface was likely limited by the formation of *OOH from the surface composed of * (empty site) and *OH, involving the direct dissociation of O-H bond.

ð25Þ

n o 0 nFkð22Þ K 0ð21Þ cH2 O exp f gð21Þ þ f ð1
að22Þ Þgð22Þ

cHþ

n o 0 j ¼ nFkð23Þ exp f ð1
að23Þ Þgð23Þ

ð24Þ

With these expressions, we expect the theoretical Tafel slope values of 30 and 120 mV dec
1 if the most abundant surface species are * and *OH, respectively, at an electron transfer coefficient of 0.5 [58], whose corresponding reaction orders is
1 and 0 with respect to the proton. While the experimentally observed value of 29 ± 1 mV dec
1 was consistent with the value of 30 mV dec
1, the kinetic order of
1.7 apparently disagreed with those theoretically anticipated values. This observation clearly excludes the possibility that the formation of *O is the rds. When rds is the formation of *OOH (Equation (22)), one can expect theoretical Tafel slope values of 21, 40, and 120 mV dec
1 at an electron transfer coefficient of 0.5, if the most abundant surface species are *, *OH, and *O, respectively [58]. The rate expressions corresponding to each case are:

j¼

n o 0 nFkð23Þ K 0ð22Þ cH2 O exp f gð22Þ þ f ð1
að23Þ Þgð23Þ

n o 0 nFkð23Þ K 0ð21Þ K 0ð22Þ cH2 O exp f ðgð21Þ þ gð22Þ Þ þ f ð1
að23Þ Þgð23Þ c2Hþ ð30Þ 441

T. Nishimoto et al.

Journal of Catalysis 391 (2020) 435–445

Fig. 4. Kinetic current densities and the ratio of the kinetic current densities as a function of the electrode potential over the IrOx at pH (or pD) 1 and 7. Kinetic current densities obtained from K-L plots over the IrOx disk electrode and the ratio of the kinetic current densities in H2O and D2O are shown, which was recorded in 0.1 mol kg
1 Naperchlorate electrolyte solutions of (a) pH and pD 1 and (b) pH and pD 7 at a scan rate of
1 mV s
1 and reaction temperature of 298 K.

current–potential relations, the kinetic current densities were obtained as a function of the applied potential in both H2O and D2O systems at pH and pD 7, as shown in Fig. 4b (see Figures S2g and S7g for the LSVs and K-L plots in H2O and Figure S12 for D2O). In both electrolytes, the anodic current densities increased with the applied potential, and the j-E relationships were nearly identical. The Tafel slope value at pD 7 was 97 mV dec
1, which is similar to the value of 100 ± 8 mV dec
1 obtained at pH 7 (see Figure S9b for the Tafel plot), suggesting a common reaction mechanism and rds in both electrolytes. Correspondingly, the ratio of the current densities, or the apparent KIE value, was unity in the investigated potential window. Similarly at pH and pD 4, in which the local pH alteration was minimal, the ratio of the current densities was unity (see Figure S13 for the LSVs, K-L plots, Tafel plots, and ratio of current densities). These finding clearly indicates that the rds during the water oxidation in near-neutral pH medium does not involve the transfer of hydrogen atom, i.e., ET. Based on these findings, we explore the reaction mechanism of the water oxidation in the unbuffered near-neutral medium. Fig. 5 illustrates the possible reaction network occurring during the oxidation of the water molecule. At acidic pH levels, the rds was the formation of O-O bond via CPET step (route 3 in Fig. 5) influenced by two preequilibrium steps (the route 1 and 2 in Fig. 5), when the single-site mechanism was adopted. Previous studies reported on the O-O bond formation occurring at Ir(V) [27,34,73], which is consistent with the oxidation peaks prior to the OER as observed in the CVs shown in Fig. 1c. Very similar CV behavior in the near-neutral pH regime to the ones at acidic pH suggests that the formation of Ir (V) prevails prior to OER also at the near-neutral pH levels, while the absence of the apparent KIE at near-neutral pH indicated the rds to be ET. To gain further insight into the rds during the water oxidation, the sensitivity of the reaction rate to the reaction temperature was examined. The obtained Arrhenius’ plot is presented in Fig. 6, which reveals that the apparent activation energies at pH 7 and 1 are 36 and 6 kJ mol
1, respectively (see Figures S14, S15 and S16 for LSVs, K-L plots and Tafel plots). Of note, the figure shows the larger j0 values at pH 7 than pH 1, which does not immediately indicate the higher catalytic activity at the neutral pH; in fact, because of the larger Tafel slope value at pH 7 (Fig. 3), we observed a superior OER performance at pH 1 (Figure S3). We remark here that the apparent activation energies for the OER were found to be small with respect to the values reported for other reaction such

3.4. Unbuffered near-neutral pH: water oxidation with induced local pH shift To investigate the OER at near-neutral pH levels in unbuffered medium, we first isolated the contribution to the overall reaction rate from the mass-transport of the generated proton by using the K-L equation [55]:

1 1 1 þ ¼ j jkin jlim

ð33Þ

where j, jkin, and jlim are the overall, kinetic, and diffusion-limited current densities, respectively. A Tafel plot for the jkin is shown in Fig. 3b at a pH value of 7 (see Figures S2g and S7g for LSVs and K-L analysis). The figure reveals a Tafel slope value of approximately 100 mV dec
1, irrespective of bulk pH values in the near-neutral pH medium from pH 4 to 10 (see Figure S4 and S10 for Tafel slope values and Tafel plots), and corresponding ln (j0) value of
8.6 ± 0.6. The observed Tafel slope was larger than that at acidic pH levels in spite of the common reactant of water molecule, indicating that the reaction rate was less influenced by preequilibrium than those observed in acidic conditions. This observation is similar to a previous study by Surendranath et al. [56], who reported the absence of contribution from preequilibrium to the overall reaction rate over a cobalt-phosphate catalyst at pH 8 in unbuffered condition due to incapability of the base present to sustain the preequilibrium before the rds [56]. The OER was insensitive to pH at these pH levels, since the OER under such conditions experiences an acid shift of the local pH value to ~ 4 regardless of the bulk electrolyte pH in the range of 4–10, as discussed in the first section. Our subsequent study using the isotope in this condition disclosed a distinct kinetics from the extreme pH conditions. Figure S11 compares the LSVs for the OER in H2O and D2O at pH and pD 7, respectively. In both electrolytes, the anodic current density increased with the applied potential. The current density in H2O was found to be larger than that in the D2O system across the potential window shown in the figure, and the ratio of the current density in H2O to that in D2O was found to be constant (1.3). Notably, the diffusion coefficients of H+ and D+ are 9.311  10
9 m2 s
1 and 6.655  10
9 m2 s
1, respectively [75], whose ratio is coincidingly 1.4. This quantitative agreement strongly supports our previous argument that the mass-transport of proton to the bulk of the electrolyte is slow relative to the appreciable current density. After isolating the diffusion contribution to the 442

Journal of Catalysis 391 (2020) 435–445

T. Nishimoto et al.

Fig. 5. Illustration of the reaction network for water oxidation at near-neutral pH levels. CPET and ET on route 3 are the possible rds at acidic and near-neutral pH levels, respectively. The structure containing an iridium site coordinated with five or six oxygen atoms is adopted in this study following previous reports [36,73,76].

The CPET at the acidic pH has smaller apparent activity than ET at the near-neutral pH levels. In other words, the low-activationenergy route of PCET is not available at the unbuffered nearneutral pH. At acidic pH levels, neighboring oxygen on the surface (likely oxo-bridging two iridium sites, as suggested for RuOx [10,83]) may function as a proton accepter during the CPET at the transition state [36]. In contrast, at unbuffered near-neutral pH, such neighboring oxygen species may be transformed to a species that does not coordinate well with proton, such as hydroxyl groups, thus hindering the transition state configuration required for CPET. The high activation energy of the ET process should be associated with the large thermal contribution upon reorganization at the transition state, which lacks O-H bond breaking. The possibility of dimerization of the surface metal-oxo species as the rds can be ruled out because of the substantially large Tafel slope value of approximately 100 mV dec
1, which is otherwise expected to be 15, 30, or larger than 120 mV dec
1 [60]. Overall, our study revealed that the OER under unbuffered near-neutral pH conditions is kinetically limited by the oxidation of the water molecule likely forming O-O bond at the Ir(V) site, which proceeds via an ET contrasting with the PCET route prevailing in the acidic water oxidation; giving another example of decoupling CPET driven not by the catalyst identity but by the electrolyte identity.

Fig. 6. Arrhenius’ plot of the OER over IrOx. The exchange current density j0 values were obtained from the Tafel plot shown in Figure S14b for pH 1 and in Figure S16 for pH 7.

4. Conclusions as HER or hydrogen oxidation reaction (HOR) ranging from 10 to 50 kJ mol
1 [77]. This observation may indicate a large entropic contribution, as has recently been discussed [78–80]. Such an entropic contribution likely originates from the reorganization of adsorbed water layer [81] claimed to impact the pre-exponential factor of the Arrhenius equation for the electrode reaction [82].

This study investigated the kinetics of the OER under unbuffered conditions over a model IrOx disk electrode in a wide pH range of 1–13. Firstly, by studying the LSVs and the onset potential-pH diagram, it was revealed that the OER in aqueous media can be divided into three parts according to the pH levels 443

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Journal of Catalysis 391 (2020) 435–445

of the electrolyte bulk at appreciable current densities on the order of mA cm
2; namely, above pH 10, below pH 5, and the pH levels in between. At the weakly alkaline pH level of 11, it was observed in the LSVs that with larger overpotentials, the current density first increased but became constant at larger overpotentials. This observation was accounted for by the reactant switching event: the OER first proceeded via the oxidation of the hydroxide ion at lower overpotentials, and then it switched to the oxidation of the water molecules at larger overpotentials achieving current densities greater than the diffusion flux of the hydroxide ion. This rationale points to the reactant switching event triggered by the pH, meaning that the OER proceeds via the oxidation of the hydroxide ion above pH 10, while the oxidation of the water molecules occurs below this threshold pH. Subsequently, by analyzing the Tafel slope, kinetic order, and isotope effect, the formation of O-O bond was identified as the rds, occurring on a clean or OH-covered surface at acidic pH levels and an OH-covered surface in the alkaline pH regimes, when the single-site mechanism is assumed. At near-neutral pH ranging from pH 5 to pH 10, the OER also involved the oxidation of the water molecule, but the picture of the OER differed from that at acidic pH. Our microkinetic analysis revealed that the rds at near-neutral pH levels was an ET likely forming O-O bonds decoupled from a PT step. The elucidated reaction mechanism in the unbuffered near-neutral pH conditions indicates that facilitating the proton transfer can improve the apparent OER performance, which can be achieved by engineering the electrode material and/or electrolyte. These insights could help in the design of OER catalysts, resulting in the improvement of (photo-) electrochemical water splitting at near-neutral pH levels.

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