Unravelling the effects of active site densities and energetics on the water oxidation activity of iridium oxides

doi:10.21203/rs.3.rs-2605628/v1

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Unravelling the effects of active site densities and energetics on the water oxidation activity of iridium oxides

https://doi.org/10.21203/rs.3.rs-2605628/v1

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published 07 Jun, 2024

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Understanding what controls the reaction rate on iridium-based catalysts is central to designing more active and stable electrocatalysts for the water oxidation reaction in proton exchange membrane (PEM) electrolysers. Here, we quantify the densities of redox active centres and probe their binding strengths on amorphous IrO_x and rutile IrO₂ using a combination of operando time-resolved optical spectroscopy, X-ray absorption spectroscopy (XAS) and time of flight secondary ion mass spectrometry (TOF-SIMs). Firstly, our results show that although IrO_x exhibits an order of magnitude higher geometry current density compared to IrO_2,the intrinsic rates of reaction per active state, on IrO_x and IrO₂ are comparable at a given potential. Secondly, we establish a quantitative experimental correlation between the intrinsic rate of water oxidation and the energetics of the active states. We use density functional theory (DFT) based models to provide a molecular scale interpretation of our data. We find that the *O species formed at water oxidation potentials have repulsive adsorbate-adsorbate interactions, and thus increasing their coverage weakens their binding and promotes the rate-determining O-O bond formation. Finally, we provide insights into how the intrinsic water oxidation kinetics can be increased by optimising both the binding energy and the interaction strength of the catalytically active states.

Physical sciences/Chemistry/Catalysis/Electrocatalysis

Physical sciences/Materials science/Materials for energy and catalysis/Electrocatalysis

Proton-exchange-membrane (PEM) water electrolysis is a key technology for converting renewable energy into green hydrogen^{1, 2}. However, the highly acidic and oxidizing conditions at the anode have thus far limited the choice of catalysts for the kinetically sluggish water oxidation reaction to oxides of iridium.^{3, 4} Iridium is one of the scarcest elements on the planet, with a production potential of < 1kt/yr.⁵ Consequently, one of the key challenges for the terawatt scale uptake of PEM electrolysers is to minimise the iridium loading at the anode. ^6–10

A number of iridium-based oxides have been studied ranging from crystalline rutile oxides, amorphous oxides, alloyed oxides^11–14 to perovskites.^{15, 16} Activity is usually compared by normalizing the current density to the geometric area or the mass of iridium. Amorphous IrO_x is more active, but less stable, than crystalline rutile IrO₂, showing a 220 mV reduction in overpotential, and approximately 2 orders of magnitude lower life-time at 1 mA cm^− 2_geo.¹⁰ Recent studies on highly active IrNiO_x^{12, 13}, Li-IrO_x¹¹ and SrIrO_x^{10, 15, 17} demonstrate that the surface of these materials is converted to an amorphous phase on exposure to OER potentials. Amorphous IrO_x is easier to oxidize and experiences greater Ir-O bond shortening with increased potential compared to crystalline IrO₂^{11, 12, 18, 19}. These differences in coordination environment of the Ir centre have been suggested to influence the binding energetics of key OER intermediates, possibly resulting in the differences in activity.^{13, 18, 20–23} However, the detailed mechanistic origins of the differences in activity between amorphous IrO_x and rutile IrO₂ remain poorly understood.

Rossmeisl and Norskov have suggested, on the basis of DFT calculations, that the potential-determining step for OER on iridium oxide is the reaction of adsorbed oxygen (*O) with water to form a peroxo intermediate (*OOH): *O + H₂O→*OOH + H⁺ + e⁻ (i.e., O-O bond formation).^{24, 25} Recent work by Nong et al combining XAS and DFT has also implied that this step is the rate-determining step (RDS).²⁶ Crucially however, Nong and co-workers suggest that the activation energy of the RDS decreases linearly with *O coverage, attributed to long-range interactions between *O species. Our recent work using operando optical spectroscopy agrees with this conclusion: an increase in OER kinetics correlates with the density of accumulated oxidized species (*O) on amorphous IrO_x. Conversely, in molecular iridium oxide, where these cooperative interactions are absent, the water oxidation kinetics were independent of *O coverage.^27–29 Although the importance of adsorbate-adsorbate interactions has been raised by Nong et al., it has been unclear how adsorbate-adsorbate interactions influence trends in activity from one catalyst to another. This lack of quantitative understanding is derived from the difficulties in theoretically and experimentally determining electroadsorption isotherms and turnover frequencies on catalysts, with the exception of single crystals.³⁰

Herein, we use operando optical spectroscopy to compare the water oxidation energetics and kinetics of amorphous IrO_x versus crystalline rutile IrO₂. We quantitatively determine the density of redox active centres as a function of potential, and quantify the interaction between adsorbate species by experimentally determining electroadsorption isotherms for both amorphous IrO_x and rutile IrO₂. The location of redox active centres is spatially resolved using isotope labelling and time of flight secondary ion mass spectrometry (TOF-SIMS). The structure and local environment of the catalysts is probed by X-ray absorption spectroscopy (XAS). Finally, time resolved optical spectroscopy is employed to determine intrinsic turn over frequencies as a function of potential. These results are corroborated with density functional theory (DFT) calculations. By experimentally probing the potential-dependent density of active states, adsorbate-adsorbate interactions and turn over frequencies in unprecedented detail, this study elucidates the key roles played by (1) the electronic structure of the active site in governing the binding energy at zero coverage and (2) the adsorbate-adsorbate interactions in controlling the weakening of the binding energy with increasing coverage, in determining the kinetics of water oxidation. Taken together these effects reveal mechanistic origins of the differences in activity between amorphous IrO_x and rutile IrO₂ for water oxidation and provide new rational design strategies for the discovery of water oxidation catalysts

Ir local structure and its influence on proton penetration for IrO _x and IrO₂.

In this study, hydrous, amorphous iridium oxide films were prepared by electrodeposition on FTO substrates, following a well-established procedure.^{27, 31} (see Supplementary Method). Rutile IrO₂ films were obtained by thermal annealing of the electrodeposited IrO_x films at 550 ℃ for 8 hours, as confirmed by XRD (Supplementary Fig. 1C). Scanning electron microscopy (SEM) images show that both amorphous and rutile IrO₂ films consist of 100–200 nm size nanoparticles (Supplementary Fig. 1A and 1B). However, in the as-deposited state, IrO_x shows a combination of Ir³⁺ and Ir⁴⁺ redox states, while the IrO₂ is dominated by Ir⁴⁺, as determined by X-ray photoelectron spectroscopy (XPS) (Supplementary Fig. 1D), consistent with previous work ^{32, 33}. This is further supported by ex-situ X-ray absorption near-edge structure (XANES) measurements at the Ir L₃ edge. The average oxidation state is determined by the position of the white line edge, which primarily corresponds to transition from occupied 2p to empty 5d states.¹² As shown in Supplementary Fig. 2, the d-band holes for amorphous IrO_x and rutile IrO₂ are around 4.4 and 5.0, corresponding to average oxidation states of around 3.4 and 4.0, respectively (see Supplementary Methods for XANES measurements and analysis). The fitted extended X-ray absorption fine structure (EXAFS) results (Fig. 1C and Supplementary Fig. 3, Table 1) show that both samples have a characteristic Ir–O interatomic distance of ~ 2.0 Å, with comparable first shell coordination number of 5.8 (± 0.5) and 6.0 (± 0.7) for amorphous IrO_x and rutile IrO₂, respectively. However, rutile IrO₂ shows a stronger signal from Ir-Ir₁ and Ir-Ir₂ (between 3 and 4 Å interatomic distance) while these signals are very weak in amorphous IrO_x. Thus, these data suggest short range ordered [IrO₆] units are dominant for amorphous IrO_x whilst a longer range ordered structure is present in rutile IrO₂, as schematically shown in Fig. 1A and 1B.

Cyclic voltammograms of rutile IrO_x show one order of magnitude higher redox current densities, when normalised to geometric area than IrO₂. IrO_x also shows higher activity than IrO₂ as reflected in the overpotential required to achieve 0.5 mA/cm_geom² (~ 210 mV and 330 mV for IrO_x and IrO₂ respectively, Fig. 1D). IrO_x shows two distinct broad redox peaks at around 0.8 V and 1.2 V vs RHE. In contrast, no distinct redox peak can be observed for IrO₂. In order to determine whether the increased current density on amorphous IrO_x is only a consequence of higher permeability of electrolyte within the porous structure (resulting in a higher density of catalytically active sites) or is also a result of higher proton penetration through the bulk of the oxide (resulting in participation of subsurface iridium sites), a combination of TOF-SIMS and deuterium isotope labeling was used. The film was impregnated and covered with paraffin wax, using an adapted method from Hadden and co-workers.³⁴ This method eliminates the effects of varying porosity on the ion incorporation (See Fig. 1E and Supplementary Figs. 4 and 5 for details) - and thus probes only the effect of surface exchange and bulk diffusion directly. Samples were cycled in 0.1 M DClO₄ (99.8% D) from 0.66 V to 1.5 V. The depth of proton penetration during water oxidation was assessed by the difference in deuterium signal intensity of the cycled samples and the samples exposed to electrolyte but without electrochemical cycling. Figure 1F shows that amorphous IrO_x cycled in DClO₄ has a higher D/H ratio than the sample without cycling, for the first 100 seconds of sputtering, corresponding to a distance of 20 nm. This signal then drops to the same level after 100 s (see Supplementary Figs. 6 and 7 and Table 2 for details). In contrast, no obvious difference in D/H ratio can be observed before and after cycling for rutile IrO₂. These results indicate that protons penetrate far deeper in amorphous IrO_x than rutile IrO₂ under CV cycling, thus offering more iridium sites for redox reactions and possibly for water oxidation. We conjecture that the short range ordered structure in amorphous IrO_x is more flexible and hence allows more proton intercalation than the long range ordered – and hence presumably more rigid – structure in rutile IrO₂. This is also consistent with the observation from atom probe tomography by Mayrhofer, Cherevko, Gault, Kasian and co-workers that indicate that hydrous IrO_x undergoes a greater degree of oxygen exchange with the electrolyte during OER than rutile IrO₂ where exchanges is limited to within 2.5 nm of the surface.^{35, 36}

Identification And Quantification Of Redox Transitions On Amorphous And Rutile Iridium Oxides

Next, we probe the redox chemistry of both materials using time-resolved operando UV-Vis spectroscopy and stepped potential spectroelectrochemistry (see Supplementary method for details). On increasing the potential, amorphous IrO_x shows broad absorption bands at 600 nm, 800 nm and 500 nm in potential ranges of ~ 0.66 V to 1.0 V vs RHE, ~ 1.1 V to 1.32 V vs RHE and > 1.32 V vs RHE respectively (Fig. 2a). Three distinct differential absorption spectral shapes can be obtained by a subtraction of adjacent spectra (Supplementary Fig. 9), and are denoted as redox 1, redox 2 and redox 3, in accordance with our previous work.^{27, 28} The differential absorption of the redox transitions in IrO₂ are very similar to those for IrO_x, with only a slight shift of the dominant absorption band for redox 1 to 700 nm (Supplementary Fig. 10). However, the intensity of the optical absorption is 6 times lower in the case of IrO₂ (Fig. 2b), consistent with this more crystalline and rigid structure having a lower density of redox active Ir centres, as discussed above. These results show that although IrO_x and IrO₂ have significantly different CV curves, they undergo the same redox processes, with both generating three similar redox states within the applied potential range.

To analyse the contribution of each redox species to the observed spectra as a function of potential, we deconvoluted the absorption spectra through a linear combination fitting process, which gives the absorption intensity from each individual redox transition as a function of potential (see Supplementary Figs. 11 and 12 for details). The obtained absorption intensities were then converted to an area density of redox species by applying a potential step and correlating the charge passed to the absorption intensity (see Supplementary Figs. 13 and 14 for details). The potential limits for the stepped potential spectroelectrochemistry measurements were selected to cover ranges in which the differential spectral shape is not changing (Supplementary Figs. 9 and 10), thereby allowing us to assume that only one redox transition takes place during the potential step. The redox peak positions obtained from these optical signals for the amorphous and crystalline iridium oxide were found to be at about 0.82 and 0.9 V for redox 1, 1.22 and 1.26 V for redox 2, and 1.47 and 1.57 V for redox 3, respectively (Supplementary Table 3, Fig. 2d and 2e). These optically determined redox transition waves match well with the less well resolved redox transitions observed in the linear sweep voltammetry curves for amorphous IrO_x (grey line in Fig. 2d, top panel). The positions of redox 1 and 2 are also in agreement with the redox peak positions on amorphous IrO_x and perovskite SrIrO_x determined by Geiger et al (both are at around 0.8 and 1.2 V)¹⁰. Notably, the redox current for redox 3 corresponding to the optical data at OER relevant potentials is masked by the OER current measured electrochemically in the CV. A redox peak at similar potential (1.63 V) has also been observed by Suntivich and co-workers on a IrO₂ (110) surface.³⁰ It is clear that the amorphous IrO_x undergoes the redox 3 process (circa 1.3 V) at approximately 100 mV lower potential than rutile IrO₂ (circa 1.4 V, see Supplementary Fig. 15 for direct absorption signal comparison). This suggests the differences in coordination between Ir centres in amorphous and crystalline iridium oxides result in stronger binding of oxygenated intermediates in IrO_x compared to IrO₂.

The concentration of the redox transitions on amorphous IrO_x was obtained by integration of the above redox waves; they show typical sigmoid shapes with increasing potential (bottom panel of Fig. 2d), with similar concentrations of redox active states participating in all three redox processes. These electroadsorption isotherms can be modelled with a Frumkin isotherm with R² values as high as 0.99, as opposed to a simple Langmuir isotherm (dashed line in Fig. 2d, bottom panel, see supplementary Fig. 16–17 for details). These excellent fits indicate the existence of lateral interactions between the redox states, as assumed in the Frumkin model. Mathematically, in a Frumkin electroadsorption model, the redox transition free energy ∆G⁰_redox is a function of the redox state coverage, i.e. ∆G⁰_redox(θ)= ∆G⁰_{redox (θ=0)} + rθ, where r (in eV) is the interaction energy of the absorbates at full coverage, θ is the coverage of the absorbates (0 < θ < 1). θ was calculated by assuming that saturating concentrations observed in our spectroelectrochemical analyses indicates full coverage, and r was obtained by fitting the θ vs U data (see supplementary Fig. 16–17 for details). The same analysis procedure was also applied to rutile IrO₂. The saturated redox concentration in IrO_x (~ 3.5x10¹⁶ cm^− 2) is more than 6 times higher than that for IrO₂ (~ 5.5x10¹⁵ cm^− 2), which we attribute to the higher proton penetration in addition to the higher porosity within the structure (Fig. 2d). The fitted interaction energy r and the half-coverage potential U_(θ=1/2) for both amorphous and rutile iridium oxides are labelled in Fig. 2D and 2E. Interestingly, we found that the adsorbate-adsorbate interaction energy r for redox 3 on amorphous IrO_x and rutile IrO₂ are similar, both at about 0.15 eV, indicating the interaction strength between the absorbates involved in this redox transition are similar despite the changes in the Ir coordination environment between amorphous and crystalline iridium oxides.

The UV-vis absorption changes in Fig. 2 can be assigned to changes in intervalence charge transfer within the Ir d orbitals, which correlates to the oxidation of Ir and the surface absorbates coordinated to the Ir centre.^37–39 To verify the nature of the surface absorbates corresponding to these redox transitions, we use density functional theory (DFT) calculations (see Computational Methods in the supplementary information). We model the rutile IrO₂ (110) surface, which consists of rows of Ir atoms coordinated by bridging oxygens separated by rows of coordinatively unsaturated (CUS) sites (inset in Fig. 2G, see details in supplementary Fig. 18). The adsorption of *H, *H₂O, *OH and *O on the CUS site and *H on the bridging oxygen (*H_b) was considered at full coverage in order to determine the most stable surface adsorbate at different potentials (Supplementary Fig. 19). The interaction between the absorbates was explored by calculating the free energy of the stepwise transition between the most stable adsorbates in a unit cell with three CUS sites, i.e at different coverages of the species. For selected surfaces the OER was modelled, two different reaction pathways were considered as described in previous work⁴⁰(Supplementary Fig. 20). Based on these energies, the redox transition potential at half coverage U_(θ=1/2), and the interaction energy, r, were calculated and plotted as adsorption isotherms in Fig. 2G. The calculated energies at half coverage for *OH (1.05 V_DFT−RHE) and *O (1.52 V_DFT−RHE) on the CUS site of rutile IrO₂ match well with experimental half-wave redox potentials for redox 2 (1.26 V) and redox 3 (1.57 V) (Supplementary Fig. 20, supplementary table 5). Our calculations are also in agreement with previous work which has attributed the redox transitions at around 1.2 V and 1.63 V on IrO₂ (110) surfaces in 0.1 M HClO₄ to *OH and *O adsorption on the CUS site, respectively.³⁰ The surface transition for redox 1 was considered as *H desorbing from the CUS site or *H₂O dissociating to form *OH on the CUS site, but the calculated potential values for both transitions (< 0.19 V_DFT−RHE) do not match closely with experimental results of 0.90V, even if the possibility of an intermediate structure with some CUS sites covered by *H₂O is considered (see supplementary Fig. 19 for details). It is therefore likely that our DFT calculations do not fully capture the nature of redox transition 1. Based on the calculations and the charge observed in the experiment, we assign the redox transitions to (also see schematic in Fig. 2C):

Redox 1: * H₂O _(cus) + *H_b → *OH_(cus) + *H_b + H⁺ + e⁻ (1)

Redox 2: *OH_(cus) + *H_b→ *OH_(cus) + H⁺ + e⁻ (2)

Redox 3: *OH_(cus) → *O_(cus) + H⁺ + e⁻ (3)

The surface of amorphous IrO_x cannot be modeled by periodic DFT calculations. Instead, as a representative of a more open IrO_x structure with larger separation between the adsorbates, the (001) surface of hollandite IrO₂ is modelled and calculations are performed for the transitions (OH_(cus) + *H_b → *OH_(cus) + H⁺ + e⁻ ) and (*OH_(cus) →*O_(cus) + H⁺+ e⁻).⁴¹ The energy of the redox transitions 2 (0.97 V) and 3 (1.34 V) on hollandite IrO_x are lower than rutile IrO₂ (Supplementary Fig. 21), in agreement with our experimental observations.

Our DFT calculations also capture the interaction strength between *OH absorbates on both the rutile and the hollandite surfaces, showing that the calculated redox potential increase from 0.89 V at 1/3 coverage to 1.20 V at full coverage of *OH for the rutile structure, and from 0.80 to 1.13 V for the hollandite structure. (see details of the interaction parameter calculation in Supplementary Fig. 22–23 and supplementary table 5). However, the interaction of *O cannot be fully captured in this model. The calculated interaction energies of *O are only about 0.02 eV and 0.08 eV for rutile and hollandite IrO₂, respectively, considerably lower than the experimentally observed value of 0.15 eV. This observation is in agreement with recent work suggesting that the repulsive interactions for *O species can only be captured by including solvent effects in the DFT model as the *O interactions might propagate through the water layer ²⁶. The above experimental and theoretical results indicate the key role of repulsive interactions between the adsorbates, which give rise to a Frumkin electroadsorption isotherms (Fig. 2D, E), and change the free energy of elementary steps in OER as a function of coverage.

Determining The Intrinsic Water Oxidation Kinetics

Next, we investigated how the free energy of the elementary steps controls the reaction rate on amorphous IrO_x and rutile IrO₂. We used two methods to estimate the intrinsic rate of the rate-determining step as a function of potential (and thus also as a function of *O coverage). Firstly, the rate of O₂ generation (obtained from the current) was divided by the concentration of redox state 3. This gives the active-state-normalized rate of O₂ release, which is equivalent to the intrinsic rate of the RDS (denoted as TOF (O₂ redox3^− 1 s^− 1)). Secondly, we measured the decay kinetics of the accumulated states when the potential was released from an oxygen-evolving potential to open circuit using transient time-resolved UV-Vis absorption spectroscopy.^{28, 42, 43} This experiment is shown in Fig. 3A, where two spectrally distinct decay phases were observed during open circuit decay – a fast decay component (initial 1 s) and a slow decay component (100–200 s) (see supplementary Fig. 24 for spectra of this two decay components). The differential absorption spectra for these two phases are in excellent agreement with the differential absorption spectra of redox 3 and redox 2, respectively (Supplementary 25), indicating the fast decay of absorption primarily arises from the decay of redox state 3 species (Fig. 3B), while slow decay mainly arises from the decay of redox state 2. Thus, this experiment can be interpreted as described in Fig. 3B: a site in the active state proceeds along the reaction path by nucleophilic attack from water molecular forming O-O bond followed by release of oxygen. Here, charge is balanced by other active states by accepting a proton and becoming reduced to regenerate *OH on the surface (since no charge is passed through the external circuit during open circuit). If the time constant of this active state decay (i.e.: fast phase) is τ, it implies an intrinsic rate of the RDS of 1/ 4τ of oxygen per second (see supplementary information for further discussion). By varying the applied potential and probing the decay when the cell is switched to open circuit, we are able to control the concentration of active states and investigate its effect on the intrinsic rate of the rate-determining step (denoted as TOF(1/4τ)) as a function of potential (detailed measurement of τ and the calculation of TOF(1/4τ) can be seen in Supplementary Figs. 26 and 27). Figure 3C shows the active state (*O) coverage and the intrinsic RDS rate of the RDS for amorphous IrO_x and rutile IrO₂. We find that while rutile IrO₂ has fewer active states at a given potential, it has the same intrinsic rate of the RDS as measured by O₂ evolution rate and slightly higher intrinsic rate of the RDS as measured by open-circuit decay.

Given that the coverage of the absorbates in the Frumkin isotherm plays a key role in the binding energetics of the adsorbates and thus controls the reactivity of the states, we next compare the TOF₃ (O₂ redox3^− 1 s^− 1) and TOF(1/4τ) for amorphous IrO_x and rutile IrO₂ as a function of the coverage of *O (Fig. 3D). We note that for both iridium oxides, log[TOF (1/4τ)] increases roughly linearly with the *O coverage in the observed coverage range. This result is similar to the observation in a recent work by Nong et al, where they found the log(current) linearly increase with the total charge and suggested from DFT that the activation energy of the RDS decreases linearly with the increase of *O coverage because of the long-range interaction among the *O species²⁶. Here, by experimentally tracking the intrinsic rate of RDS as a function of *O coverage, we observe similar results. We also note that the slopes of log[TOF (1/4τ)] vs *O coverage in Fig. 3D are similar in amorphous IrO_x and rutile IrO₂, consistent with our observation that they have similar absorbate-absorbate interaction energies r. The repulsive interaction observed here between oxo species causes a destabilization of adsorbed *O species (ΔG_*O-ΔG_*OH becomes larger), thus reducing the free energy of the RDS (O-O bond formation from *O).

From Fig. 3D, it is also apparent that at the same *O coverage, the intrinsic activity of the RDS on rutile IrO₂ is around 1 order of magnitude higher (as measured by TOF (O₂ redox3^− 1 s^− 1)) and factor of 5 higher (as measured by TOF (1/4τ)) than in amorphous IrO_x. We rationalize this difference in activity to the potential dependent binding energetics of *O. From the experimentally observed energetics of *O, the values of ΔG_*O-ΔG_*OH at *O (θ_*O = 0) for rutile IrO₂ is around 0.1 eV larger than that of amorphous IrO_x (requiring an extra 0.1 V to form *O on rutile), and both increase with *O coverage at a slope of roughly 0.15 eV/coverage due to the similar Frumkin interaction strengths (Fig. 3E). As a result, at a constant coverage of *O, compared with amorphous IrO_x, rutile IrO₂ binds oxygen more weakly, which facilitates O-O bond formation from *O in the rate-determining step, leading to higher intrinsic activity per active state. On the other hand, at a constant potential, for example at 1.48 V, this binding energy difference decreases to around 0.05 eV because amorphous IrO_x has a higher coverage of *O at this potential, thus weakening the *O binding more significantly. Therefore, the intrinsic activity at a constant potential is similar, and the order of magnitude greater geometric-area-normalized current density observed for the amorphous is mainly because it has a higher concentration of active states, as shown in Fig. 2D,E.

Design Principles For Highly Active Catalytic Materials For Water Oxidation

The potential (and coverage) dependence of the intrinsic water oxidation kinetics determined herein show the importance of the absorbate-absorbate interactions in controlling the kinetics of the OER reaction. The conventional approach for catalyst design has focussed on developing catalysts with optimal binding energetics (ΔG_*O-ΔG_*OH) corresponding to approximately 1.6 eV ^{24, 25, 44}. This implicitly assumes that this value is coverage independent. Herein, we have clearly experimentally demonstrated the weakening of the *O binding energy with increasing coverage. This indicates that optimal binding energetics can also be obtained on surface sites that are strongly binding at low coverage ((ΔG_*O-ΔG_*OH <1.6 eV at θ_*O ~ 0), but have strong repulsive interactions, which weaken *O binding energies at high coverage, thereby enabling these sites to achieve the optimal ΔG_*O-ΔG_*OH values under reaction conditions corresponding to higher coverage.

To include the influence of absorbate-absorbate interactions on ΔG_*O-ΔG_*OH values, and thus the intrinsic activity per state, we construct a new 3-dimensional volcano plot. This volcano plot can be seen as a deconvolution of the previous descriptor ∆G_O - ∆G_OH into two components (1) binding energetics on a surface with zero *O coverage, ∆G_O - ∆G_OH(θ_*O = 0), and (2) interaction strength between *O species, following a Frumkin-isotherm-related equation [∆G_O - ∆G_OH = ∆G_O - ∆G_OH(θ_*O = 0) + r*θ_*O]. We first construct this plot assuming near saturated coverage of θ_*O = 0.9. As shown in Fig. 4A, the optimal activity of the states can be achieved either (1) via an optimal *O binding at low coverage without *O interaction (i.e. ΔG_*O-ΔG_*OH ~1.6 eV at θ_*O ~ 0, r = 0 eV), which is the top of conventional volcano plot; or (2) via a stronger *O binding than optimal and strong interaction (e.g. ΔG_*O-ΔG_*OH=1 eV, at θ_*O ~ 0, r ~ 0.7 eV). In the latter case, although each state binds *O much stronger than optimal at zero coverage, the strong interaction strength results in weakening binding and optimum ΔG_*O-ΔG_*OH at high coverage.

We note that the applied potentials to achieve the same coverage are different for catalysts with different ∆G_O - ∆G_OH(θ_*O = 0) and r values. To include both potential and coverage effects, we construct a volcano plot at a fixed overpotential of 250 mV (Fig. 4B, and supplementary Fig. 28). The steep part of the left side of the volcano (for example, when ∆G_O - ∆G_OH(θ_*O = 0) < 1.18 eV in the curve of r = 0.15 eV) is where *O coverage is saturated, whereas the less steep part represents where ∆G_O - ∆G_OH(θ_*O) is changing due to the Frumkin repulsion term. The Frumkin repulsion term has a higher impact on increasing the activity for a material with stronger *O binding energy, for a given interaction strength. This is evident from our data on iridium-based catalysts, where the absorbate-absorbate interaction increases the relative activity of the two materials by different degrees, although they have similar interaction parameter (~ 0.15 eV). This is because the stronger binding strength of *O on the IrO_x surface results in a higher coverage for IrO_x compared to IrO₂ at this potential (Fig. 2D,E), and consequently higher repulsion between adsorbates. However, both show higher activity than the molecular Ir catalyst analysed previously, which have no adsorbate-adsorbate interaction.²⁸ This observation is consistent with our previous suggestion that although the studied molecular Ir showed faster kinetics than amorphous IrO_x at low overpotentials, it is less active than amorphous IrO_x at higher overpotentials due to a lack of cooperative interaction between active states²⁸. Similar analyses at overpotentials of 200 and 400 mV are also given in Supplementary Fig. 29. These results indicate that potential can alter the coverage of *O species, which in turn can weaken the binding energies for catalysts that bind oxygen too strongly in a conventional volcano plot, hence improving OER activity. This analysis elucidates the fundamental origin of the activity discrepancies between different Ir-based catalysts, including amorphous and crystalline rutile iridium oxides.

In conclusion, our work combines experimental and theoretical techniques to re-evaluate the design principles for water oxidation catalysts by unravelling the role of the density of active sites, their ordering and the interaction between them on the reaction rate. Here, we have resolved the formation of the redox centres in amorphous IrO_x and crystalline IrO₂ as a function of potential, using time resolved UV-vis spectroelectrochemistry and as a function of depth using TOF-SIMS combined with deuterium labelling. We have used DFT to provide a molecular-level interpretation to our results. We assign the third spectral redox wave to the transition of *OH to *O on Ir_cus site in both catalysts from DFT; the experimentally obtained O₂ evolution kinetics show an exponential dependence on the coverage of this species *O, corroborating predictions from earlier DFT based models that O-O bond formation controls the reaction rate.^24–26

Our potential decay results reveal that the intrinsic kinetics of O₂ evolution are comparable, or slightly faster, on amorphous IrO_x relative to crystalline IrO₂ at a given potential; however, the greater penetration and hence density of sites per unit geometric area of redox centres in amorphous IrO_x, as measured both by optical spectroscopy and TOF-SIMS, account for the significantly higher geometric current density observed in this catalyst, relative to crystalline IrO₂. IrO_x binds more strongly to *O relative to crystalline IrO₂ at low coverage; nonetheless under the same potential at reaction condition, the coverage of *O is much higher than that of crystalline IrO₂, leading to greater weakening of *O as a result of repulsive adsorbate-adsorbate interactions. These two opposing effects lead to comparable kinetic rates on the two types of IrO₂ under study. Based on these insights, we might expect that nanocrystalline Ir-based catalysts, taking on the high stability^{10, 45} and high intrinsic rate of rutile IrO₂ but with the high site density of amorphous IrO_x, would provide optimal performance for O₂ evolution.

Beyond iridium oxides, our work highlights the role of adsorbate-adsorbate interactions on water oxidation catalysis; they are particularly important for catalysts on the strong binding side of the oxygen evolution volcano, such as non-precious metal oxide MnO_x²⁵. Our work leads to a more accurate description of water oxidation kinetics than conventional models based solely on the binding energetics at zero coverage; by doing so we also provide an additional, unexplored lever to aid the discovery of more active catalysts.

Data availability

The Spectroelectrochemistry data and Frumkin isotherm fitting data were analysed by Matlab, the deconvolution of absorption spectra was analysed using Matlab and Julia programming language. The data and analysis scripts will be openly available at https://github.com/Caiwu-L/Amorphous-IrOx-vs-Rutile-IrO2. The details of home-built spectroelectrochemistry set-up and the Labview based software is openly available at https://opensourcespectroscopy.com/. DFT optimised structures and python scripts used to analyse the data are available from https://nano.ku.dk/english/research/theoretical-electrocatalysis/katladb/water-oxidation-activity-of-iridium-oxides/. All data in the study are available from the corresponding authors upon reasonable request.

Acknowledgements

I.E.L.S., J.R.D., R.R.R, and C.L. would like to acknowledge the funding and technical support from BP through the BP International Centre for Advanced Materials (bp-ICAM) as well as the EPSRC grant EP/W033232/1 which made this research possible. C.L. acknowledges the Imperial College London and China Scholarship Council for the IC-CSC joint scholarship. K.L.S, A.M.F and J.R acknowledges the Danish National Research Foundation Center for High-Entropy Alloy Catalysis (CHEAC) DNRF-149. We acknowledge Diamond Light Source for time on Beamline B18 under Proposal SP30396.

Author contributions

C.L, R.R.R., I.E.L.S. and J.R.D. conceived the idea. C.L., R.R.R. and B.M. performed the XAS experiments. C.L. performed the optical spectroscopy measurements; B.M. custom-built the optical spectroscopy set-up. The code for linear combination fitting of the optical spectroscopy data was developed by M.S. and adapted by C.L., who analysed the data with the help of R.R.R., J.R.D., I.E.L.S, S.B.S and B.M.. K.L.S, A.M.S and J.R. performed and analysed the DFT results. J.H.L.H performed the TOF-SIMS experiment; M.P.R and D.J.R. supervised the experiments and contributed to the data analysis. The manuscript was written by C.L.; all authors reviewed and edited the manuscript.

Competing interests

The authors declare no competing interests.

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IrOxvsIrO2supportingCaiwu20230130.docx
Unravelling the effects of active site densities and energetics on the water oxidation activity of iridium oxides

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Unravelling the effects of active site densities and energetics on the water oxidation activity of iridium oxides

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Introduction

Results

Identification And Quantification Of Redox Transitions On Amorphous And Rutile Iridium Oxides

Determining The Intrinsic Water Oxidation Kinetics

Design Principles For Highly Active Catalytic Materials For Water Oxidation

Conclusion

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References

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