Deciphering the high overpotential of the oxygen reduction reaction via comprehensively elucidating the open circuit potential

Abstract

The onset potential of the oxygen reduction reaction (ORR) maintains an activation overpotential exceeding ∼0.20 V, even with the most efficient catalysts. Despite decades of devotion, the origin of such high overpotential has been a long-standing topic of controversy. As the critical point for generating an apparent current in the ORR below the open circuit potential (OCP), understanding the reasons behind the OCP contributes to decrypting the inherent essence of the overpotential. Unfortunately, the implications represented by time-dependent OCP have not been clarified. In this work, the time-dependent OCP has been categorized into three stages: an initial rapid transition stage, a quasi-steady state stage, and a steady-state stage, wherein the quasi-steady state stage has been substantiated as the mixed potential arising from the oxidation of Pt and the ORR, accompanied with the dissolution of Pt. The quasi-steady OCP persists for an extended period due to the sluggish kinetics of the ORR on the oxidized Pt surface, the specific value of which was predominantly determined by the kinetics of Pt oxidation. It has been verified that the concept of the quasi-steady OCP is rooted in kinetics rather than equilibrium thermodynamics. Therefore, elevating the energy barrier of Pt oxidation through the substitution of H₂O by D₂O reduced the overpotential of the ORR significantly, evidencing that the bottle-neck of reducing the overpotential of the ORR is the inhibition of the oxidation of electrocatalysts kinetically. Hence, the reduction peak potential of Pt-based catalysts during the cathodic scan in linear sweeping voltammetry can be selected as a novel and straightforward indicator for predicting ORR performance and expediting the screening of highly efficient catalysts.

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Broader context

Rising global population and energy demand drive interest in fuel cells for direct electricity conversion. However, the oxygen reduction reaction (ORR) at the cathode faces a persistent activation overpotential of approximately 0.20 V, even with advanced catalysts. Despite decades of research, the origin of this high overpotential remains a contentious issue. The prevailing theory advocates reducing energy barriers for the rate-determining steps on catalysts to enhance ORR efficiency. Conventional studies prioritize identifying intermediate species, largely overlooking how catalyst surface oxidation intertwines with the impact on the ORR. This study delves into the time-dependent open circuit potential (OCP) during the ORR, categorizing it into three stages: initial rapid transition, quasi-steady state, and steady-state. The quasi-steady state, validated as a mixed potential from the ORR and Pt oxidation, persists due to sluggish ORR kinetics on the oxidized Pt surface. The concept is rooted in kinetics rather than equilibrium thermodynamics. By inhibiting Pt oxidation kinetics, the overpotential is reduced. The reduction peak potential of Pt-based catalysts during cathodic scans emerges as a novel indicator for predicting ORR performance, proposing a strategy to enhance OCP and reduce overpotential by kinetically inhibiting Pt oxidation – an innovative approach promising improved catalytic efficiency in electrochemical systems.

Introduction

Proton exchange membrane fuel cells (PEMFCs) have been considered as the alternative to traditional heat engines in the future.¹ However, the cost of anodic and cathodic catalysts used in PEMFCs is high due to the limited resources of platinum.^2,3 Especially, the oxygen reduction reaction (ORR) at the cathode with sluggish apparent reaction kinetics requires much more Pt loading than anodic hydrogen oxidation, hindering the commercial application.^4,5 In addition, unfortunately, the onset potentials of most Pt-based catalysts are still lower than 1.0 V vs. RHE despite much effort devoted over the past few decades.^6,7 These are far away from the thermodynamic standard potential of 1.23 V vs. RHE.

O₂ + 4H⁺ + 4e⁻ ⇄ 2H₂O E^θ = 1.23 V (R1)

The origin of the high overpotential has been explained by the site-blocking theory that Pt sites were poisoned by the strong adsorbed oxygen-pertaining species generated at high potentials.^7–9 Consequently, O₂ molecules cannot be accessible to the free active Pt sites, hindering the initial step of the ORR. Meanwhile, Li et al. stated that the competition between the protonation and hydrolysis of the adsorbed *O from O₂ dissociation determines the onset potential of the ORR.¹⁰ Assisted by the ab initio molecular dynamics, they found that *O protonation is bypassed by hydrolysis, resulting in a high onset potential. Koper et al. proposed that the transition between *O and *OH can be considered as the kinetic descriptor determining the onset potential.¹¹ Apart from the aforementioned argument, the flawed presumption in theoretical calculations lies in assuming that the surface is metallic across a broad potential range. However, in reality, the state of the Pt surface varies from the applied potential.^12–14

Experimentally, the parameters involved in the rotating disk electrode tests such as temperature, oxygen partial pressure, rotating rate, and pH value of electrolytes have been thoroughly optimized. However, the overpotential of the ORR is hardly lower than 0.2 V. For example, the onset potential was increased a little even though the partial pressure of O₂ gas was increased to 100 atm from 1 atm, suggesting that the rate of the ORR is not determined by thermodynamics.¹⁵ Hence, Feliu et al. explored the early stage of the ORR on Pt, uncovering that the initial reaction at elevated potentials involves a complex sequence of chemical–electrochemical–chemical–electrochemical reactions coupled with a disproportionation reaction. This intricate process is the primary cause behind the significant overpotentials observed in the ORR.^16,17 Beyond the realm of reaction kinetics, researchers have focused on understanding the impact of the Pt surface state on the ORR process. This attention is warranted because the Pt surface undergoes a transition from metallic in the potentials approaching a thermodynamic equilibrium potential of 1.23 V.^18,19 In this case, as the oxygen reduction reaction is kinetically inhibited on the oxidized surface, an overpotential is needed to reduce the surface for starting the ORR. From reported experimental results, the process of the ORR on the Pt surface entangled with the oxidation of the Pt surface has been sketched roughly. Nevertheless, the interplay between the ORR and the oxidation of the Pt surface has not been fully unravelled, especially at the potentials marking the onset of the ORR.

Undoubtedly, the equilibrium potential can be calculated by the thermodynamic parameters based on the reaction formula. The assessment of the maximum voltage achievable by an electrochemical cell is often determined by the open circuit potential (OCP). This value is conveniently measured using an electrochemical workstation under zero-current conditions when the potential difference remains unchanged over a specific period. Although the OCP is an old concept of more than sixty years,^20,21 the inherent significance of the value directly obtained for the OCP remains elusive, sparking ongoing debates, particularly regarding whether it reflects thermodynamic information or kinetic information. According to the definition, the OCP is different from the equilibrium potential. For a system with a single reaction of the ORR consisting of a series of reversible faradaic elemental steps, if all elementary steps have fast reaction kinetics, a stable OCP can be obtained within a short period, which is identical to the thermodynamic potential of the ORR numerically. Otherwise, it takes a longer time to reach a stable OCP value for the ORR consisting of steps of slow kinetics. At the quasi-steady state, the OCP is determined by the rate-limiting step, which can be used to describe the kinetics of the ORR on the catalyst of interest. The higher the quasi-steady OCP, the lower the overpotential relative to 1.23 V of the ORR. Ultimately, the OCP will reach the thermodynamic potential limit of the ORR over a super long time.

When the system involves additional parallel reactions, such as electrode activation and passivation, elucidating the OCP value becomes challenging. Unfortunately, the Pt surface is partially oxidized during the OCP test and it is difficult to reach the steady state.²² Therefore, Pt oxidation and the ORR take place concurrently at the OCP. In such instances, the OCP at the quasi-steady state fails to accurately depict the overpotential on the pristine Pt surface, given the continuous changes occurring on the Pt surface upon immersion into an O₂-containing electrolyte. Ideally, the ORR can take place once the applied potential is lower than the equilibrium potential of the ORR. However, in fact, the substantial reduction current can be observed only when the applied potential is far lower than an equilibrium potential of 1.23 V. Experimentally, the onset potential E_onset of the ORR located very closely to the OCP recorded by an electrochemical workstation in the lab. Therefore, it is crucial to unveil the OCP for a comprehensive understanding of the ORR.

We state that the following essential questions should be recognized and emphasized, and are extremely important in revealing the formation of the high overpotential of the ORR. First, is it true that the ORR is always accompanied by the oxidation of the Pt surface in the high potential range? It will be beneficial to solve the issue of discrepancy in the reported Tafel slopes with varied values from 60 to 120 mV dec⁻¹. Furthermore, what significance does the time-dependent OCP hold in the presence of both Pt and O₂ in an acidic medium? It will be helpful to gain the real number of electron transfers of the limiting step in the ORR. Finally, how can one logically verify and demonstrate that the kinetic inhibition of the ORR occurs on the oxidized Pt surface through experimental tests? This exploration is crucial for identifying intrinsic parameters that enhance the ORR.

Therefore, we comprehensively investigated and elaborated the underlying time-dependent OCP of the ORR on a Pt/C covered rotating-disk electrode based on the well-established experimental design and precise experimental control. For the first time, the Pt oxidation followed by its dissolution has been identified as a predominant parameter of determining the quasi-steady OCP. Inspired by the perspectives, the couple reactions occurring at the onset potentials of the ORR are disentangled. Furthermore, we attributed the quasi-steady OCP to characterize the non-equilibrium properties of an electrocatalyst for the ORR, rather than focusing on its equilibrium properties. As a result, we can clarify the origin of the increased overpotential in the ORR, accompanied by the introduction of a convenient criterion for evaluating ORR performance.

Results and discussion

Pt surface oxidation at the OCP in an acidic medium

Fig. 1(a) shows that the OCP in an O₂-free electrolyte was 0.503 V, which was elevated sharply after entering O₂, suggesting that a new state tends to be established on Pt nanoparticles supported on carbon (Fig. S1, ESI†). The OCP was increased with P_O2 in line with the results that high P_O2 minimized the overpotential of the ORR (η_ORR). In the O₂-saturated electrolyte, the OCP was held at 1.024 V, which is approximately 0.2 V lower than the standard equilibrium potential of 1.23 V for the 4-electron ORR, consistent with the other experimental results.^21,23,24 By applying the Nernst equation (eqn (1)) corresponding to (R2), the number of electron transfers (n) derived from the slope of the plot of against E^eq should be 1 for the ORR.where R is the gas constant, T is the temperature, and F is the Faraday constant. E^eq and E^Θ are the equilibrium and standard potentials of (R2), respectively. a_H⁺ and a_H2O are the activities of protons and water, respectively. P_O2 and P₀ are the partial pressures of O₂ gas and ambient pressure, respectively.

Fig. 1 (a) O₂ partial pressure effect on the OCP. The insets in (a) are the number of electron transfers at the OCP calculated by the Nernst equation of eqn (1) of the reaction (R2). (b) pH effect on the OCP. (c) Effect of H₂Pt(OH)₆ and O₂ on the OCP. (d) OCP test of Pt/C in O₂-saturated 0.1 M HClO₄ and the concentration of the dissolved Pt ion during the OCP test.

However, n was calculated to be 0.54, much smaller than 1 as shown in the inset of Fig. 1(a). According to the Marcus theory, the elementary step most likely involves a single-electron transfer on the energy scale.^25,26 Therefore, the non-integer value of n implies that the OCPs under different P_O2 cannot be simply regarded as the equilibrium potentials of the ORR, but rather there may exist some oxidation reactions that make n less than 1 on a macroscopic scale. If the OCP was just determined by the ORR, pH will not change the OCP versus RHE. In contrast, the OCP increased with pH as shown in Fig. 1(b), inferring that additional reactions beyond the ORR were proceeding during the OCP test. Furthermore, the OCP had a rapid increase till a plateau of 0.875 V in the presence of H₂Pt(OH)₆ without O₂ as shown in Fig. 1(c). After the OCP plateaued in the presence of O₂, the introduction of H₂Pt(OH)₆ caused a sudden increase in the OCP to 1.041 V, implying that Pt ions should be involved while establishing OCP. Experiments have confirmed that the Pt surface was partially oxidized around 1.0 V, accompanying the dissolution of Pt ions into the electrolyte.^12,14 The dissolution of Pt into the electrolyte in the presence of O₂ occurred as well at the OCP testing as shown in Fig. 1(d). The Pt dissolution rate was rapid initially on the introduction of O₂, corresponding to the first sharp transition stage of the time-dependent OCP curve. The concentration of the dissolved Pt ions increased slowly, corresponding to the quasi-steady state of the time-dependent OCP. During a span of 48 hours, merely 0.28% of platinum (Pt) underwent dissolution, indicating the OCP for an exceptionally prolonged quasi-steady state. The detectable Pt ions in the solution confirmed the result that Pt ions influenced the OCP as shown in Fig. 1(c), and verified the oxidation reaction in addition to the ORR proposed according to Fig. 1(a) and (b). Therefore, it is confirmed that the OCP was influenced by both the ORR and Pt oxidation.

Evidencing the slower rate of the ORR on oxidized Pt sites than that on Pt sites

To study the effect of Pt oxidation on the ORR, the initial oxidation state of Pt was controlled by adjusting the upper potential limit (E_upl) and the quiet time at E_upl. As shown in Fig. 2(a), the half-wave potential (E_1/2) of the ORR decreased continuously as E_upl increased during the cathodic scan, while the limiting current density did not change. Fig. 2(b) shows that the cathodic reduction peak current density was increased with E_upl in the O₂-free electrolyte, resulting from the reduction of more Pt surface oxides at higher E_upl. Fig. 2(c) shows that the E_1/2 of the ORR and the cathodic reduction peak current density had an opposite trend as E_upl increased. Moreover, the initial oxidation state of Pt can be regulated by adjusting the quiet time at E_upl as well. Fig. 2(d) shows that as the quiet time increased, the oxidized degree of Pt increased, leading to a decrease in E_1/2 of the ORR and an increase in the cathodic peak current density in the O₂-free electrolyte. For an E_upl of 1.07 V, the mass-transfer corrected Tafel slopes were 67 and 114 mV dec⁻¹ in the high and low potential ranges as shown in Fig. S3 (ESI†), respectively, indicating that the reaction kinetics changed with potentials. The results comprehensively confirm that the highly oxidized Pt surface severely affects the ORR.

Fig. 2 E _upl effect on the LSV curves obtained in (a) O₂-saturated and (b) O₂-free 0.1 M HClO₄ electrolytes at 1600 rpm at a scan rate of 10 mV s⁻¹. (c) E_upl effect on E_1/2 of LSV curves obtained in the O₂-saturated electrolyte (blue) and the peak current density in the cathodic scan in the O₂-free electrolyte (red). (d) Effect of quiet time at an E_upl of 1.17 V on E_1/2 of the LSV curves in the O₂-saturated electrolyte (blue) and the peak current density in the cathodic scan in the O₂-free electrolyte (red). The values in (d) are extracted from Fig. S2 (ESI†).

The effect of Pt oxidation on the ORR was quantitatively analyzed by testing the influence of different scan directions and scan rates. Unlike the phenomena that cathodic scans varied from E_upl as shown in Fig. 2(a), the anodic scans overlap very well as shown in Fig. 3(a). Moreover, E_1/2 and current density during the anodic scan were higher than those during the cathodic scan at the same potential, agreeing well with the ORR on the Pd electrode.²⁷ Interestingly, the cathodic scan curves overlapped very well at different scan rates while the anodic scan curves varied as shown in Fig. 3(b). Moreover, a high scan rate corresponded to high E_1/2 during the anodic scan. To further understand the intrinsic reason for the hysteresis between the anodic and cathodic scans as shown in Fig. 3(a) and (b), the coverages of Pt (θ_Pt) and oxidized Pt (θ_O) sites were extracted from the cyclic voltametric curves of Pt in the O₂-free electrolyte (Fig. S5, ESI†) and plotted as the x-axis as shown in Fig. 3(c) and (d), respectively. It is assumed that both types of sites are occupied by the reactants involved in the ORR. Hence, the total current j_T without considering the capacitive current (eqn (2)) is the summation of j₁ (eqn (3)) generated on Pt sites and j₂ (eqn (4)) on Pt oxide sites.

j_T = j₁ + j₂ (2)

where α and η are the symmetry factor and the overpotential of the ORR, respectively. j₀₁ and j₀₂ are the exchange current densities of the ORR on metallic Pt and oxidized Pt sites, respectively. The current densities in the cathodic and anodic scans are labeled as j_Tc and j_Ta, respectively. Fig. 3(a) and (d) show that j_Ta during the anodic scan was always higher than that (j_Tc) in the cathodic scan at the same η, namely j_Ta > j_Tc.

Fig. 3 (a) The cathodic and anodic scans subtracted by the background in the O₂-free 0.1 M HClO₄ electrolyte with different E_upl in the O₂-saturated 0.1 M HClO₄ electrolyte at 1600 rpm at a scan rate of 10 mV s⁻¹. (b) The cathodic and anodic scans with different scan rates in the O₂-saturated 0.1 M HClO₄ electrolyte subtracted by the background in the O₂-free 0.1 M HClO₄ electrolyte at 1600 rpm with an E_upl of 1.07 V. The original curves are shown in Fig. S4 (ESI†). (c) The potential dependent coverage of Pt during the cathodic and anodic scans with different E_upl in the O₂-free electrolyte, which was extracted from the curves in Fig. S5 (ESI†) (detailed calculations in the ESI†). (d) The potential dependent coverage of Pt during the cathodic and anodic scans with different scan rates in the O₂-free electrolyte.

In this case, eqn (5) can be derived.

where θ_Pt/a + θ_O/a = 1 and θ_Pt/c + θ_O/c =1.

By substitution and simplification, we can get eqn (6) from eqn (5).

Fig. 3(c) and (d) show that θ_Pt in the anodic scan is always larger than that in the cathodic scan, namely, θ_Pt/a − θ_Pt/c > 0. Hence, eqn (7) can be obtained.

Assuming that the symmetry factors α₁ and α₂ are equally close to 0.5, it is deduced that j₀₁ is always larger than j₀₂ at the same η (eqn (7)), demonstrating that the reaction rate of the ORR on metallic Pt sites is faster than that on the oxidized Pt sites. Therefore, it is safe to conclude that the oxidized Pt sites weakened the ORR rate undoubtedly. In summary, the ORR current is significantly influenced by the oxidized degree of the Pt surface, which is investigated by changing the scan direction and scan rate. Fig. 3(a) and (c) show that in the cathodic direction, the scan starts from a region where the Pt surface is covered by oxides and the coverage of Pt (θ_Pt) decreases as the E_upl increases, leading to a suppression of the ORR activity. Conversely, the anodic scan begins with a metallic surface at low potentials, and as the potential increases from reduction to oxidation, θ_Pt decreases. Fig. 3(c) demonstrates that θ_Pt is independent of E_upl in the anodic scan and larger than that in the cathodic scan at the same potential, corresponding to higher ORR current density. Fig. 3(b) and (d) show that for the effect of the scan rate on the oxidized degree of the Pt surface, a higher scan rate in the anodic direction implies that the Pt surface has a shorter time to be oxidized at high potentials, leading to higher θ_Pt and higher ORR current density. However, for the cathodic direction, since the Pt is pre-oxidized at the beginning of the scan, the scan rate above 10 mV s⁻¹ seems to have no significant effect on the reduction of oxidized Pt in the subsequent process. Therefore, the difference in ORR activity is not significant, which is consistent with the results observed by Minteer et al. on the Pd electrode.²⁷ Accordingly, reducing the oxidized degree of the Pt surface will enhance the reduction current of the ORR.

Illustration of the transient process of OCP establishment

Before analyzing the time-dependent OCP, the classical “mixed potential theory” is introduced as depicted in the Evans diagram as shown in Fig. 4(a), which establishes a connection between potentials and currents arising from various components, combining them into a ‘weighted’ potential when there is a zero net current. The Evans diagram necessitates anodic and cathodic linear segments, with the mixed potential (E_mix) and the total current density pinpointed at the intersection of the oxygen reduction line and the metal oxidation line in a metal–water–oxygen system. Also, the slope of each redox reaction correlates to the exchange current. In this case, each reaction is no longer at the equilibrium state. It should be recalled that the OCP exhibited a rapid increase upon the introduction of O₂ into the O₂-free electrolyte, aiming to establish a new stable state as shown in Fig. 1(a). At the newly established state, Pt was dissolved into the electrolyte (Fig. 1(d)). With time increase, the concentration of Pt ions was increased with a decreased dissolution rate, suggesting that the dissolution of Pt occurred always but was mitigated by surface passivation over time or inhibited by the increased Pt ions thermodynamically. Accordingly, the OCP was elevated extremely slowly to reach a quasi-steady state. Based on the experimental results aforementioned, the underlying process can be effectively elucidated by referring to the scheme illustrated in Fig. 4(b)–(e). The initial OCP (ϕ₁) was 0.503 V in the O₂-free electrolyte, resulting from the spontaneously established electrical double layer (EDL) at the Pt electrode–water interface. Once O₂ gas was introduced into the electrolyte, the Pt surface interacted with the dissolved O₂ to elevate the potential of the interface to ϕ₂, where the ORR and Pt surface oxidation coexisted. In this case, the OCP indeed was the E_mix raised by Wagner and Traud in the field of corrosion science as depicted in Fig. 4(a).²⁸Fig. 4(c) shows the coexistence of Pt oxidation and the ORR made the net current zero at E_mix. The slope of Pt oxidation γ is smaller than |β| of the ORR at the onset potential region, evidenced by the positive overall current (oxidation current) in the high potential region in the anodic scan in the O₂-saturated electrolyte (Fig. S6, ESI†). The fast oxidation rate of Pt was confirmed as well.¹⁸

Fig. 4 (a) Classical mixed potential theory depicted by the Evans diagram. (b) Schematic of the state of OCP established in the Ar-saturated electrolyte. (c) Schematic Evans diagrams illustrating the OCP transition during the introduction of O₂ into the Ar-saturated electrolyte. E_mix is the corrosion potential that leads to the Pt oxidation due to the presence of the oxidant of dissolved O₂ in an acidic medium. (d) Illustration of the OCP at the quasi-steady state as the increase of the oxidizing degree of the Pt surface. (e) Illustrated OCP is at the steady state.

As shown in Fig. 4(c), when O₂ was introduced into the electrolyte instantaneously, according to the Nernst equation and the collision theory, the increase of P_O2 leads to the increase of E_ORR and j₀₁ (causing the red dashed line shifting upward and right to the red solid line). This also caused the intersection point of the red and blue curves to rise significantly, reflecting the increase of E_mix. This explains why the OCP increased sharply once P_O2 suddenly increased in Fig. 1(a). Over the time, platinum oxidation (denoted by the blue solid line) followed by its dissolution gradually reached the quasi-steady state as shown in Fig. 4(d), and hence, E_mix was increased slowly to ∼ϕ₃ but always smaller than E_ORR. Such mixed potential theory can explain why the electron transfer number n in Fig. 1(b) was less than one. According to the classical mixed potential theory, the E_mix located closer to the thermodynamic potential of the half-reaction with a higher exchange current is depicted in Fig. 4(a). Moreover, the higher exchange current of Pt oxidation caused E_mix to approach E_PtOR (the thermodynamic potentials associated with a series of oxidation reactions involving Pt in Fig. S7, ESI†) more closely at the quasi-steady state of the OCP. However, the Pt corrosion in the presence of O₂ was different from the typical metal corrosion, since Pt was oxidized to solid Pt oxide rather than directly to soluble Pt ions. Therefore, platinum oxidation could not remain at a constant rate. As the degree of oxidation increased, the amount of Pt that could be oxidized decreased, resulting in a decrease of the exchange current, and at the same time, the reaction it underwent also changed constantly, causing E_PtOR to rise continuously, eventually forming a transition from the blue dashed line to the blue solid line as shown in Fig. 4(d). Therefore, the quasi-steady OCP (recorded potential on the electrochemical workstation) is E_mix within the specified time interval, which is close to the equilibrium potential of the reaction with a higher reaction rate. Consequently, the OCP is limited by the equilibrium potential of Pt oxidation or the potential of its potential depending step rather than the ORR. To observe a substantial ORR current, the applied potential should be lower than the quasi-steady OCP, enabling the reduction of Pt oxides and lowering the activation energy of the ORR (Fig. S8, ESI†). In the end, the OCP will attain a value of 1.23 V for the ORR only if the equilibrium potential of Pt oxidation is increased to 1.23 V or if the Pt oxidation ceases.

Kinetically inhibiting Pt oxidation reduced the overpotential of the ORR

It has been widely accepted that the H₂O molecule rather than the OH⁻ anion takes part in the electrooxidation of Pt in the acidic media. To investigate the role of water in Pt oxidation, therefore, H₂O/H₂SO₄ was replaced by D₂O/D₂SO₄. Fig. 5(a) shows that n is 0.56 in the deuterium electrolyte in the average, which was larger than that in the H₂O/H₂SO₄ electrolyte, indicating that more proportion of the ORR was contributed in the deuterium system at E_mix even though the absolute value of the total current was quite small. In other words, Pt oxidation became slower in D₂O/D₂SO₄ during establishing the quasi-steady OCP compared with that in the H₂O/H₂SO₄ electrolyte, making OCP rise more in D₂O/D₂SO₄ under the same P_O2. Extra information can be extracted that the potential depending step in the ORR in the acidic medium was not likely to involve H-pertaining intermediates. Otherwise, n should be smaller in the deuterated electrolyte. By adding H₂O/H₂SO₄ into the D₂O/D₂SO₄ electrolyte, a continuous decrease of the quasi-steady OCP from 1.027 V in D₂O/D₂SO₄ to 1.004 V in H₂O/H₂SO₄ was observed within the interval, confirming that H₂O decreased OCP as shown in Fig. 5(b). This was due to the faster oxidation rate of Pt in the H₂O/H₂SO₄ electrolyte. The CVs in O₂-free electrolytes as shown in Fig. 5(c) show that the onset potential of Pt oxidation or the surface *OH formation shifted positively and the cathodic peak potential had a positive shift of 9 mV, indicating that Pt oxidation in D₂O/D₂SO₄ became slower. Namely, the activation energy of Pt oxidation was increased in the deuterium system due to the isotopic kinetic effect.

Fig. 5 (a) The partial pressure of the O₂ effect on the OCP in 0.5 M H₂SO₄/H₂O and 0.5 M D₂SO₄/D₂O. The original curves are shown in Fig. S9 (ESI†). (b) The volume ratio of D₂SO₄/D₂O to the H₂SO₄/H₂O effect on the OCP in the O₂-saturated electrolyte. The original curves are shown in Fig. S9 (ESI†). (c) CV curves of Pt/C in O₂-free 0.5 M H₂SO₄/H₂O and 0.5 M D₂SO₄/D₂O electrolytes at a scan rate of 50 mV s⁻¹. The insets in (c) are the scheme of D₂O increasing the activation barrier of Pt oxidation and the shift of the reduction peak potential. (d) Cathodic LSV curves of Pt/C in O₂-saturated 0.5 M H₂SO₄/H₂O and 0.5 M D₂SO₄/D₂O electrolytes at a scan rate of 10 mV s⁻¹. The inset in (d) is the H/D kinetic inverse effect (KIE) based on the kinetic current density extracted from the Koutecký–Levich equation.

In terms of the proposed mechanism, E_1/2 exhibited positive shifts of 16 and 14 mV of the ORR for the cathodic and anodic scans in D₂O/D₂SO₄, respectively, as shown in Fig. 5(d) and Fig. S10 (ESI†). The positive shift in the millivolt level can be reproduced very well (Fig. S11, ESI†). Moreover, the mass-transfer corrected Tafel slope was decreased from 146 mV dec⁻¹ in the H₂O/H₂SO₄ electrolyte to 130 mV dec⁻¹ in the D₂O/D₂SO₄ electrolyte in the low potentials as shown in Fig. S12 (ESI†), suggesting the accelerated kinetics of the ORR in the deuterated system. The inset in Fig. 5(d) shows that higher current density was obtained in the deuterated electrolyte, agreeing well with the isotope kinetic effect by the Feliu group recently.²⁹ Meanwhile, upon the analysis, we contributed the improved ORR performance to the high activation energy of Pt oxidation due to the low activity of D₂O compared to H₂O (the inset in Fig. 5(c)). Additionally, the Pt surface state can be affected by the dissolved Pt ions. In the case of dissolved Pt ions in the electrolyte, the reduction of Pt surface oxide can be inhibited during the negative scan, resulting in a negative shift of 48 mV for E_1/2 of the ORR (Fig. S13, ESI†).

Proposing Pt reduction peak potential as the descriptor

Pourbaix diagrams of Pt–Pt oxide and O₂–H₂O systems in Fig. 6(a) show that the formation of Pt oxide can occur at lower potentials relative to the ORR, qualitatively suggesting that at the onset potential of the ORR the redox of Pt and O₂ shall be entangled together. Pt oxidation is fast kinetically and reaches the equilibrium state readily.¹⁸ As illustrated in Fig. 6(a), the quasi-steady OCP was located between the two equilibrium potentials and closer to the equilibrium potential of the reaction with a faster reaction rate.²⁸ Therefore, the quasi-steady OCP was determined by Pt oxidation kinetics. Therefore, lowering the activation energy of the ORR can increase E_1/2 as shown in Fig. 6(b). Oppositely, decreasing the activation energy of Pt oxidation can reduce E_1/2. Fig. 5(d) proves that lowering the rate of Pt oxidation promoted the rate of the ORR due to exposing more Pt sites to O₂. In nature, the overpotential of the ORR can stem from either the activation energy of the ORR or the number of accessible Pt sites. The former is determined by the interaction between the intermediates and active sites, whereas the latter is determined by the oxidation kinetics of the active sites, suggesting that the low overpotential will be obtained on the surface which can hold the metallic state at high potentials. Increasing the number of exposed metallic Pt sites by adjusting the initial oxidized state of Pt can significantly shift E_1/2 positively (Fig. 2), indicating that a high concentration of metallic Pt sites is crucial for maintaining the high ORR performance. Conventionally, the presence of D₂O does not enhance the reaction rate of the ORR due to the isotopic effect on certain elementary steps involving the cleavage or formation of H-bonds. In contrast, ORR performance was enhanced (Fig. 5(d)) in the deuterated electrolyte where Pt oxidation was kinetically inhibited (Fig. 5(c)) to guarantee the higher fraction of metallic Pt sites. Hence, it is reasonable to conclude that the kinetics of Pt oxidation play a major role in determining the overpotential of the ORR. Based on the discussions above, therefore, the reduction peak potential in the cathodic scan can be considered as a practical descriptor to correlate with E_1/2 of the ORR. Fig. 6(c) and Table S1 (ESI†) summarize the differences between E_1/2 (ΔE_1/2) and cathodic peak potential E_cp (ΔE_cp) of some ORR target materials relative to the reference materials reported in the high-impact journals, respectively. E_1/2 can be increased by alloying,³⁰ constructing,^31,32 decorating with organic molecules,³³ and exposing to different electrolytes.³⁴ The summarization shows that no matter how E_cp was increased, E_1/2 was elevated, keeping the same change direction with E_cp. Therefore, the cathodic reduction peak potential of the LSV in the O₂-free electrolyte can be considered as the electrochemical indicator for the ORR in addition to the indicators based on the material structure.^35–37

Fig. 6 (a) Pourbaix diagram of Pt and O₂ in water solution. (b) Scheme of the competing role of Pt redox and the ORR in determining the E_1/2 in LSV. (c) The shift of the peak potential (ΔE_cp) in the cathodic scan in the O₂-free electrolyte between the reference and target materials (dark color). The shift of E_1/2 (ΔE_1/2) in the O₂-saturated electrolyte between the reference and target materials (light color). The inset is the positive relationship between ΔE_cp and ΔE_1/2.

Conclusions

Assisted by the fact of Pt dissolution, the OCP on the Pt/C electrocatalyst in the O₂-saturated acidic electrolyte was comprehensively elucidated in nature, which indeed is a mixed potential of the ORR and Pt oxidation. According to the mixed potential theory and the continuous decay of the rate of Pt oxidation, it is concluded that the process of reaching the steady E_mix is continuously establishing a new state of zero overall current since the current of the ORR and Pt oxidation decayed continuously. Within the time scale of laboratory research, the OCP reached the quasi-steady state, where each reaction proceeded with an extremely slow reaction rate. In addition, the ORR rate on Pt oxides is lower than that on metallic Pt and has been confirmed quantitatively. The quasi-steady OCP was close to the potential of Pt oxidation, which has a higher exchange current than the ORR in the onset potential region of the ORR. Therefore, enhanced capability in inhibiting the kinetics of Pt oxidation makes the surface expose more Pt sites to O₂, leading to a small overpotential. Finally, we proposed that the cathodic reduction peak potential can be the novel descriptor in predicting the performance of the ORR. Therefore, this work emphasizes the negative effect of Pt oxides on the ORR and provides new insight into designing highly efficient ORR catalysts by enhancing the capability in antioxidation.

Data availability

Characterization, electrochemistry, and calculations can be found in the ESI.†

Author contributions

Z. W. and G. Y. performed electrochemical measurements. G. Y., F. P., and Z. L. conceived and supervised the project and wrote the manuscript. Q. Z. and Z. L. revised the manuscript. All the authors discussed the results and approved the manuscript.

Conflicts of interest

There are no conflicts to declare.

Acknowledgements

This work was supported by the National Natural Science Foundation of China (22008076 and 22272038). We extend our sincere gratitude to Dr Jun Huang at Forschungszentrum Jülich GmbH in Germany for his illuminating theoretical discussion on the concept of the open circuit potential.

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Footnote

† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d3ee04368a

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