Solvation Effects for Oxygen Evolution Reaction Catalysis on IrO

Article pubs.acs.org/JPCC

Solvation Effects for Oxygen Evolution Reaction Catalysis on IrO2(110) Joseph A. Gauthier,†,‡ Colin F. Dickens,†,‡ Leanne D. Chen,†,‡ Andrew D. Doyle,†,‡ and Jens K. Nørskov*,†,‡ †

SUNCAT Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, Stanford, California 94305, United States ‡ SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, 2575 Sand Hill Road, Menlo Park, California 94025, United States ABSTRACT: We study the electrochemical interface between rutile IrO2(110) and water to investigate how the inclusion of an explicit solvent influences the stabilities of adsorbed intermediates in the oxygen evolution reaction. Solvent is modeled by explicit nondissociated water molecules, and their structure is determined by a global optimization method. We find that the inclusion of an explicit solvent can significantly affect the geometry of adsorbed intermediates, changing from an interaction with the surface to an interaction with the water bilayer. These water structures consist of stacked octagonal sheets in an ordered network. Solvent stabilization is pronounced for adsorbed *OH and *OOH, which are capable of donating hydrogen bonds. We find little to no change in adsorbate binding energy as the number of layers of solvent is increased from 1 to 3, suggesting a single water bilayer is sufficient to describe the system. With either *O or *OH coadsorbates, the energetics of the reaction pathway are relatively unchanged with the inclusion of explicit solvent.

1. INTRODUCTION

H 2O(l) + * → *OH + (H+ + e−)

(1)

Solvent can play a very important role in chemistry, both energetically and mechanistically. For example, in biology, aqueous solvation can significantly stabilize zwitterionic species and consequently affect the predicted mechanism from a concerted gas phase process to a solvated stepwise process.1 In electrochemical reactions, water also frequently acts as an electrolyte, mediating ion transportation between the anode and the cathode. There is much interest in understanding how water interacts with solid oxide surfaces. In particular, many studies have sought to elucidate the water−metal interface for close-packed metal surfaces,2−5 and some studies have investigated the water−solid interface for transition metal oxides such as TiO2 and RuO2 due to their application in solar cell and fuel cell technologies.6−10 Solvation effects are especially important for electrochemical reactions that take place at the electrode−electrolyte interface, such as the oxygen evolution reaction (OER). Previous theoretical studies have been successful in recreating OER electrocatalyst activity trends using a reaction mechanism consisting of the four proton-coupled electron transfer (PCET) steps11−13 shown in eqs 1−4. This mechanism proceeds through a step that produces an adsorbed hydroperoxy species, which has been observed experimentally under OER conditions on some metal surfaces.14,15

*OH → *O + (H+ + e−)

(2)

*O + H 2O(l) → *OOH + (H+ + e−)

(3)

*OOH → O2 (g) + (H+ + e−)

(4)

© XXXX American Chemical Society

The effect of solvation on the stability of OER intermediates has been studied in the context of the oxygen reduction reaction (ORR) on transition metal surfaces. By calculating explicit solvent stabilizations of 0.1−0.3 eV for *OH on Pt(111) depending on coverage,16 it was suggested early on that solvent is critical in understanding the energetics of the ORR. Tripkovic et al. find *OOH to be stabilized by 0.5 eV using a half dissociated water layer network on Pt(111),17 and Karlberg et al. and Rossmeisl et al. find *OH to be stabilized by about 0.4 eV on several close-packed metal surfaces.18,19 On the basis of thermodynamic reasoning, Calle-Vallejo et al. estimated the solvation stabilization of *OH and *OOH on metal surfaces to be about 0.30 eV.20 In 2016, Liu et al. found that, on Pt(111), *O and *OO receive little stabilization, while *OH is stabilized by about 0.6 eV, and *OOH is stabilized by about 0.7 Received: March 13, 2017 Revised: May 9, 2017 Published: May 10, 2017 A

DOI: 10.1021/acs.jpcc.7b02383 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C eV21 when using 2−3 bilayers of explicit solvent. In a more recent study, Briquet et al. used an incomplete water bilayer over IrO2(110) to determine that ΔGO − ΔGOH decreases by about 0.2 eV, whereas ΔGOOH − ΔGOH increases by about 0.2 eV with solvent, depending on the exchange-correlation functional chosen.22 The structure of solvent at the electrode interface is much less studied on transition-metal oxides, which make up the majority of OER catalysts. Many recent studies have also investigated solvation effects by solving a linearized Poisson−Boltzmann equation, and treating the solvent implicitly as a polarizable dielectric continuum that ramps from one in the surface, to a bulk dielectric constant away from the surface. For example, the Adaptive Poisson−Boltzmann Solver (APBS)23 and VASPsol24,25 packages both solve the Poisson−Boltzmann equation to calculate a solvation energy provided by an implicit solvent. Using VASPsol, Sakong and Groß illustrate that, in this model, inclusion of an aqueous environment can significantly affect selectivity in the electro-oxidation of methanol.26 Using the APBS package, Sha et al. find that, on Pt(111), *O is stabilized by about 0.7 eV, *OH is stabilized by about 0.4 eV, *OOH is stabilized by about 0.5 eV, and *OO is stabilized by about 0.4 eV.27 Implicit solvation calculations have the distinct advantage of being much cheaper computationally relative to explicit solvent calculations. However, they predict stabilizations that are very different from the results of a more explicit inclusion of solvent, possibly due to neglecting local effects such as hydrogen bonding. This study will focus on specific interactions using an explicit solvent model. In this work, we use Density Functional Theory (DFT) and a global optimization technique (minima hopping) to investigate the structure of the electrode−electrolyte interface and how it affects the thermodynamics of the OER. Both RuO2(110) and IrO2(110) are known to be low overpotential OER catalysts. IrO2 is reasonably stable in acidic environments, an important prerequisite for their use in proton exchange membrane electrolyzer technologies.28,29 However, concerns have been raised over the corrosion of RuO2(110),30,31 with one study claiming that up to 10% of the current generated in the OER on RuO2(110) goes to corroding the surface compared to 1% for IrO2(110).32 For this reason, we investigate the IrO2−water interface. Siahrostami et al. have previously investigated the influence of water on the OER,33 which involved optimization of the water structure. The method for water placement in the present work involves a global minimization scheme called minima hopping, described in the next section, which we believe yields the most stable solvent structure on IrO2(110) and represents a more systematic approach to determining water structures on surfaces. In the following sections, we report only the most stable water configurations on the IrO2 surfaces.

electrons have been expanded as plane waves with a kinetic energy cutoff of 600 eV and a density cutoff of 6000 eV. The bulk rutile IrO2 structure was found by optimizing over lattice constants a and c and internal coordinate u in an 8 × 8 × 12 Monkhorst Pack (MP) k-point grid.40 The optimal structural parameters found were a = 4.58 Å, c = 3.21 Å, and u = 0.308, which are in agreement with previous theoretical findings33,41,42 as well as with experimental results.43 The (110) facet of the catalyst surface was modeled using a 1 × 2 × 4 supercell, with the bottom two layers fixed to the bulk positions and the top two layers allowed to relax, and the Brillouin zone was sampled with a 4 × 4 × 1 MP k-point grid. Vacuum separation was at least 8 Å from the top of the highest water molecule to the bottom of the periodic slab as determined by convergence tests. Geometry optimizations were considered converged when the residual forces were smaller than 0.03 eV/ Å. It was determined by Siahrostrami and Vojvodic that, in a (2 × 1) rutile supercell, water bilayers are most stable with four water molecules in the first bilayer above the surface.33 Explicit solvent structures were optimized using the minima hopping global optimization method44 and implemented in ASE,45 starting with previously reported minima.33 Structures were optimized for the three primary intermediates of OER (*OH, *O, *OOH), at both OH and O coverages, and at four levels of explicit solvent inclusion, namely, 0−3 water bilayers. Additionally, water structures for *OO were optimized similarly to the OER intermediates. OH/O overage are defined as having an *OH/*O coadsorbate on the neighboring coordinatively under-saturated site (CUS). In our model, this means both nearest neighbors to the OER intermediate of interest are determined by the coadsorbate. A Hookean constraint with a cutoff distance of 1.2 Å was placed between O and H in *OOH during geometry optimization to prevent spurious hydrogen transfer to either the neighboring CUS or a facilitated hydrogen transfer to a bridge oxygen during the molecular dynamics step of minima hopping. Following structure optimization by minima hopping, the most stable structure in each combination of (coverage) and OER intermediate was taken and the active adsorbate was swapped with the other two OER intermediates. For example, if OH was adsorbed to the active CUS, the *OH was removed and replaced with *O and *OOH, and geometry optimized. The resulting optimized geometries were then compared in energy to the most energetically stable minima hopping result, and the overall lowest was taken as the most stable energetically overall geometry. This was done to more effectively explore the phase space of stable water structures. If the minima hopping technique found a particularly stable water structure for a given adsorbate, the same water structure might be stable for a different adsorbate as well. It is important to note that it has been shown that, for liquid water, dispersion can be a significant factor in the DFT determined water structure,46 and, furthermore, that RPBE tends to slightly overstructure liquid water.47 It remains unclear whether these effects extend to the ice-like water structures presented here. If dispersion effects are important for ice-like water structures, it is expected that the globally optimized water structure would be less structured. To determine the most stable vacant site at each level of explicit solvent, the solvent structures for each adsorbate determined from minima hopping were considered. Cases where an explicit solvent molecule adsorbed to the vacant CUS were neglected, as this substantially changes the local water

2. COMPUTATIONAL DETAILS The stability of OER intermediates on IrO2(110) was investigated using DFT calculations in conjunction with the computational hydrogen electrode (CHE) model.16 All of the calculations were performed using the spin-polarized revised Perdew−Burke−Ernzerhof (RPBE)34 exchange-correlation functional in Quantum Espresso35 interfaced through the Atomic Simulation Environment (ASE).36 The core electrons of iridium, oxygen, and hydrogen are represented in this work by ultrasoft pseudopotentials,37−39 where their valence B


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The Journal of Physical Chemistry C

Figure 1. Top and side views of the most stable water configurations on the IrO2(110) surface for increasing number of bilayers (BL). The images presented demonstrate the water structures for *OOH on the active CUS, for O covered IrO2(110). Blue spheres represent iridium, red spheres represent oxygen, and white spheres represent hydrogen. Shown below each image is the free energy to adsorb *OOH relative to H2O(l) and H2(g).

bridging oxygen sites have formed hydrogen bonds with the bridging oxygen atoms, while the water molecules above the CUS interact more strongly with the *OOH. Similarly, bond conformation changes are observed when considering *OH on the active CUS. With no explicit solvent, the *OH weakly hydrogen bonds with the neighboring CUS oxygen. With the inclusion of explicit solvent, the *OH is stabilized more strongly by hydrogen bonding with the water molecules. With *O on the active CUS, however, there is no observed stabilization with the inclusion of explicit solvent. The adsorbate on the neighboring CUS plays an important role in determining the bond conformation of the active adsorbate. With *O on the neighboring CUS, there is little stabilization to be gained from *OH on the active CUS, though without solvent, the *OH will still donate a weak hydrogen bond to the neighboring *O. *OOH has little to gain from a neighboring *O as well. If directed at the neighboring CUS oxygen, generally the hydrogen will spontaneously transfer, forming *OO and *OH on the two sites. However, if the neighboring CUS has *OH adsorbed, networks of hydrogen bonding can be formed across the CUS site with the periodic image. *OOH can find stabilization by receiving a hydrogen bond from the neighboring *OH, which lessens the stabilization received from explicit solvent. The final resulting solvent structure for each OER intermediate and neighboring CUS combination in the single

density in the innermost layer. The vacant site solvent structures did not undergo the minima hopping global minimization scheme, because the increased kinetic energy during molecular dynamics caused explicit solvent to adsorb to the vacant site in all cases. Hence, it is expected that there may exist a more stable solvent configuration for the vacant site. Furthermore, the same vacant site solvent structure was used for each of the three binding energies. Therefore, the binding energies in this work may underestimate experimental values. It is worth noting that the differences in binding energies, namely, ΔGOOH − ΔGOH and ΔGO − ΔGOH, do not rely on the state of the vacant site.

3. RESULTS AND DISCUSSION 3.1. Bond Conformation Changes with Explicit Solvent, Overall Solvent Structure. Figure 1 shows the most stable water structures found for an IrO2(110) surface with *OOH adsorbed at the active CUS and *O on the neighboring CUS, along with *O at each bridge site (O coverage). This figure also illustrates changes observed in the bond conformation of the adsorbate. Without explicit solvent, the hydrogen on the *OOH species is most stable when directed at the bridging oxygen. In the presence of explicit solvent, a more stable conformation can be found by forming a hydrogen bond with the solvent. The water molecules over the C


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Figure 2. Solvent structure for the single bilayer case over IrO2(110). The oxygen atoms of water molecules are colored in green to differentiate them from the surface oxygen atoms in red. Blue spheres represent iridium atoms, while white spheres represent hydrogen atoms.

liquid water is calculated by calculating the free energy of water vapor at the vapor pressure of water at 300 K. The water vapor is then in equilibrium with liquid water, and hence their chemical potentials are identical. For step 4 of the OER mechanism, the experimental free energy of formation of H2O is used to avoid describing the high spin triplet state of O2(g) which has been shown to give errors in DFT.49 In order to more clearly illustrate the effects of solvation on *OOH, the case of no solvent with *OOH on an oxygen covered IrO2(110) surface is shown with the *OOH pointing up instead of hydrogen bonding at the bridge site. The magnetic moment on the oxygens in OOH changes significantly between the two states, which significantly affects the energy. The resulting binding energies are shown in Figure 4. The addition of explicit solvent has a small effect for *OH and *O. For *OOH on an oxygen covered surface, ΔGOOH decreased by about 0.3 eV in the presence of explicit solvent; at OH coverage, ΔGOOH increases by about 0.2 eV, suggesting that *OOH is destabilized with the introduction of a received hydrogen bond when it is already donating a hydrogen bond. It is worth noting that, if *OOH is allowed to hydrogen bond with the adjacent bridging oxygen, the apparent effects of solvation are much smaller due to the changing magnetic moments. The effect of explicit solvent appears to be converged with 1−2 bilayers of solvent, suggesting treatment of electrochemical reaction energies and barriers is adequate at this level of complexity. It has been found previously that eqs 2 and 3 usually represent the most thermodynamically challenging steps for the OER. Additionally, it has been found that ΔGOOH − ΔGOH is nearly constant across a wide variety of metal and metal-oxide surfaces (3.2 ± 0.2 eV).13 This leads to the use of ΔGO −

bilayer case is an octagonal sheet of water molecules. Figure 2 illustrates the octagonal sheet for an oxygen covered IrO2(110) surface. The octagonal structure persisted when optimizing in a 4 × 1 supercell, suggesting that the structure does not originate from the smaller unit cell and periodic boundary conditions. This solvent network is analogous to the hexagonal water structure found on Pt(111),5,48 considering that the separation between metal atoms in IrO2(110) is significantly larger than on Pt(111). In the case of multiple bilayers, the water network forms stacked octagonal sheets, with each sheet translated slightly in the x and y directions. Between the octagonal sheets, rings of size 5 and 7 water molecules form in an ice-like hydrogen bonding network. Figure 3 shows a heptagonal ring forming between the first and second sheets, and a pentagonal ring forming between the second and third sheets. This solvent network is consistent among the water structure determined for all OER intermediates and surface coverages. 3.2. Effects on the Thermodynamics of the Oxygen Evolution Reaction. Binding energies were calculated relative to liquid water and hydrogen gas. The free energy of each reaction in eqs 1−4) was calculated as ΔG RXN = ΔG0 − eU

(5)

where ΔG0 = ΔZPE + ΔE − TΔS0, with E as the calculated DFT electronic energy, and S0 the calculated entropic contributions assuming adsorbed species have only vibrational degrees of freedom. The entropy of gas-phase species was obtained under the ideal-gas limit. ZPE is the zero-point energy contribution, and − eU is the free energy contribution of the electron chemical potential.16 Any solvation effects are included in the calculated DFT electronic energies. The free energy of D


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ΔGOH as a convenient descriptor of OER activity, which is shown alongside ΔGOOH − ΔGOH as a function of explicit solvent amount in Figure 5. As can be seen in Figure 5, the OER binding energy difference ΔGO − ΔGOH does not show a strong dependence on the explicit solvent. For the *O covered IrO2(110) surface, ΔGO − ΔGOH increases by about 0.2 eV with explicit solvent. For *OH coverage, ΔGO − ΔGOH changes by less than 0.1 eV with explicit solvent. However, with *O on the neighboring CUS, ΔGOOH − ΔGOH decreases by about 0.3 eV in the presence of explicit solvent due to the large stabilization of *OOH. For hydroxyl coverage, ΔGOOH − ΔGOH does not significantly changeincreasing by about 0.1 eV. As was the case with the binding energies, the free energy differences seem to have converged with 1−2 bilayers of explicit solvent, indicating the effects of explicit solvent can be reasonably determined without greatly increasing the cost of our computations. To determine the cause of the stabilization of *OOH relative to *OH, we investigated two possible competing factors beyond solvation. First, it is possible that, with varying the adsorbate and coverage, charge may accumulate near the surface, which would affect the stability of the adsorbed species. However, Bader charge analysis on the slab atoms revealed that there is no appreciable difference in charge accumulation in the metal based on coverage or adsorbate. Second, the *OOH species is somewhat strained in the “up” position as it is when solvated relative to bulk hydrogen peroxide, and so, it is possible that the strain could cause some of the stabilization observed. To investigate the impact of this effect, we considered non-spin-polarized calculations to eliminate the observed change in energy as a result of changing magnetic moment. In this case, the *OOH “up” position is only very slightly less stable than when *OOH hydrogen bonds with the bridging oxygen, about 0.05 eV, and so this cannot fully explain the stabilization observed.

Figure 3. Illustration of the solvent network forming over an oxygen covered IrO2(110) surface with multiple bilayers. In dark blue is the heptagonal ring forming between the first and second layers. Yellow shows the pentagonal ring forming between the second and third layers. Green oxygen atoms show the two water molecules in this unit cell that are shared between the pentagonal and heptagonal ring structures. Light blue, red, and white represent iridium, oxygen, and hydrogen atoms, respectively.

Figure 4. Binding energies of OER intermediates on rutile IrO2(110) in the presence of varying amounts of explicit solvent. In this plot, red triangles represent a binding energy with OH on the neighboring CUS, and black triangles represent a binding energy with O on the neighboring CUS. E


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Figure 5. OER descriptors on rutile IrO2(110) in the presence of varying amounts of explicit solvent. Red triangles represent descriptor values with OH on the neighboring CUS, and black triangles represent descriptor values with O on the neighboring CUS.

Figure 6. Illustration of estimation of stabilization due to favorable solvent interaction. Here, Gopt refers to the optimized geometry with explicit solvent, Gsolvent refers to the solvent water molecules, Gslab refers to the slab without solvent, and Gsolvation refers to the estimated stabilization the system gains from the favorable interaction of the solvent with the surface.

by the presence of an electric field on Pt(111), despite the intermediates containing dipoles.50 Because water contains a dipole, it stands to reason that the energetics and structure of the solvent layer near the surface could be sensitive to this imposed electric field. Previous studies found that water structures at the surface of RuO2 and TiO2 experience very little perturbation in moderate strength electric fields.51 We found, similarly, that, when applying an electric field of 0.2 V/Å, the most stable water structure does not change and, furthermore, the binding free energies of OER intermediates change negligibly. It is important to note that the water structures presented here were determined in the 0 K temperature limit. Furthermore, the electric field present near the electrode surface will reduce the degrees of freedom of the water molecules in the inner Helmholtz layer. At ambient temperatures, the water structure will sample various local minima, while here we only report results for what we determine to be global minima. Therefore, these results represent an upper-

To directly estimate the magnitude of solvation’s role in the stabilization of *OOH relative to *OH, we considered a series of single-point calculations detailed in eq 6 below and illustrated in Figure 6. Here, we take the free energy of the optimized structure with explicit solvent, and subtract from it the free energy of the separate components, the metal slab + adsorbate, and the solvent as calculated by single-point DFT calculations (i.e., non-optimized geometries). This estimates the stabilization the adsorbate receives from the solvent. We then compared the stabilization of *OOH to *OH and found that *OOH is stabilized more than *OH, by 0.3−0.5 eV. It hence follows that the stabilization observed is due primarily to the solvent. ΔGsolvation = Gopt − (Gsolvent + Gslab)

(6)

At the surface of an experimental water oxidation electrode, there would exist an electric field as a result of the applied positive potential and ions near the surface. It was found previously that OER intermediates are not significantly affected F


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Figure 7. Free energy diagram of the OER on the ideal OER catalyst (dashed line), on rutile IrO2(110) with no solvent (solid line), and on rutile IrO2(110) with explicit solvent (blue line).

unlikely to be the rate-limiting step in the reaction. This seems plausible based on a calculated desorption free energy change of 0.7 eV from the free energy diagram above, especially considering that DFT tends to overestimate the stability of the O−O bond (we calculate the error on the O2 molecule relative to H2O and H2 is approximately −0.7 eV using RPBE). An alternative, but similar, pathway first involves the transfer of hydrogen to a neighboring bridge or CUS *O and has been proposed previously to occur concertedly as part of the third step, completely bypassing *OOH as an intermediate.52 These uncertainties surrounding the stability of *OOH on the surface highlight the need for calculation of kinetic barriers and microkinetic modeling for this reaction, which are outside the scope of this study As a result of the method used to determine optimal solvent structure for each adsorbate and coverage combination, the free energy diagram above is based on global energetic minima. For nearly all cases, the positions of the oxygens in the solvent water molecules from step to step are very similar, with only rotations in the hydrogen bonds relative to each other, meaning that transitioning from one step to the next involves some reorganization of the solvent. Vartia et al. found the characteristic reorientational time for bulk water at 298 K to be 2.6 ± 0.1 ps.53 We can relate this to the rate of water reorientation and roughly estimate a free energy barrier for reorienting a water molecule in a hydrogen bond network by eq 7 below. Using this expression, we calculate a free energy barrier of about 0.05−0.1 eV. However, it has been found that water near a metal surface can form an ice-like bilayer on strongly interacting surfaces.54,55 Indeed, the reorientational time for water reorientation is significantly higher than in bulk water, on the order of nanoseconds.56 This corresponds to a reorganization barrier of 0.2−0.3 eV, as was used in previous

bound for adsorbate stabilization. Hence, the idealized hydrogen bond network might be disrupted, and so, the stabilizations presented here are an upper bound on the stabilization of OER intermediates resulting from solvation. 3.3. Free Energy Diagram and Mechanistic Implications. The changed binding energies of OER intermediates lead to a modified free energy diagram for the OER on IrO2(110) as shown in Figure 7. Because the binding energies of adsorbates are calculated relative to H2 and H2O, the free energy diagram shown is effectively at 0 V vs the Reversible Hydrogen Electrode (RHE), where OER is 4.92 eV uphill in free energy. The ideal catalyst is one that has four equal-sized steps of 1.23 eV, such that, at the OER equilibrium potential (1.23 V vs RHE), the free energy landscape is flat. The free energy diagram shown in Figure 7 shows that IrO2(110) binds *OH and *OOH too strongly, and hence requires a higher potential than 1.23 V to drive the OER to completion. As a result of the stabilization of *OOH, steps 2 and 4 are left as the most thermodynamically difficult in the proposed mechanism. The fact that the fourth step, *OOH removal, is suggested to be limiting in this case raises important questions about the stability of adsorbed *OOH. In previous theoretical studies that utilize the same reaction mechanism as the one considered here, scaling relationships between OER intermediates result in either steps 2 or 3 being thermodynamically limiting, and the fourth step has not been carefully considered as a limiting step.13 Rather than a concerted removal of a proton from *OOH with the desorption of *OO, it is possible that these steps happen in series. In fact, the electrochemical removal of such a proton has a much more favorable driving force (1.0 eV) compared to concerted *OOH removal (1.8 eV). The desorption of *OO must then happen chemically, and based on experimental observation of Tafel behavior for the OER, it is G


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The Journal of Physical Chemistry C electrochemical oxygen reduction kinetics studies.57 As such, the energetics presented here are more useful for thermodynamic rather than kinetic analyses. rate =

k ∗T − ΔGa 1 = B ∗ e kB ∗ T time h

Transition and Noble Metal Surfaces. Phys. Rev. Lett. 2003, 90, 216102. (3) Michaelides, A. Density Functional Theory Simulations of WaterMetal Interfaces: Waltzing Waters, a Novel 2D Ice Phase, and More. Appl. Phys. A: Mater. Sci. Process. 2006, 85, 415−425. (4) Carrasco, J.; Hodgson, A.; Michaelides, A. A Molecular Perspective of Water at Metal Interfaces. Nat. Mater. 2012, 11, 667−674. (5) Ogasawara, H.; Brena, B.; Nordlund, D.; Nyberg, M.; Pelmenschikov, A.; Pettersson, L. G. M.; Nilsson, A. Structure and Bonding of Water on Pt(111). Phys. Rev. Lett. 2002, 89, 276102. (6) Cheng, J.; Sprik, M. Acidity of the Aqueous Rutile TiO2(110) Surface from Density Functional Theory Based Molecular Dynamics. J. Chem. Theory Comput. 2010, 6, 880−889. (7) Cheng, J.; Sulpizi, M.; Vandevondele, J.; Sprik, M. Hole Localization and Thermochemistry of Oxidative Dehydrogenation of Aqueous Rutile TiO2(110). ChemCatChem 2012, 4, 636−640. (8) Lister, T. E.; Chu, Y.; Cullen, W.; You, H.; Yonco, R. M.; Mitchell, J. F.; Nagy, Z. Electrochemical and X-Ray Scattering Study of Well Defined RuO2 Single Crystal Surfaces. J. Electroanal. Chem. 2002, 524−525, 201−218. (9) Sun, C.; Liu, L.-M.; Selloni, A.; Lu, G. Q.; Smith, S. C. TitaniaWater Interactions: A Review of Theoretical Studies. J. Mater. Chem. 2010, 20, 10319. (10) Watanabe, E.; Rossmeisl, J.; Bjorketun, M. E.; Ushiyama, H.; Yamashita, K. Atomic-Scale Analysis of the RuO2/water Interface under Electrochemical Conditions. J. Phys. Chem. C 2016, 120, 8096. (11) Rossmeisl, J.; Logadottir, A.; Nørskov, J. K. Electrolysis of Water on (Oxidized) Metal Surfaces. Chem. Phys. 2005, 319, 178−184. (12) Rossmeisl, J.; Qu, Z. W.; Zhu, H.; Kroes, G. J.; Nørskov, J. K. Electrolysis of Water on Oxide Surfaces. J. Electroanal. Chem. 2007, 607, 83−89. (13) Man, I. C.; Su, H. Y.; Calle-Vallejo, F.; Hansen, H. A.; Martínez, J. I.; Inoglu, N. G.; Kitchin, J.; Jaramillo, T. F.; Nørskov, J. K.; Rossmeisl, J. Universality in Oxygen Evolution Electrocatalysis on Oxide Surfaces. ChemCatChem 2011, 3, 1159−1165. (14) Sanchez Casalongue, H. G.; Ng, M. L.; Kaya, S.; Friebel, D.; Ogasawara, H.; Nilsson, A. In Situ Observation of Surface Species on Iridium Oxide Nanoparticles during the Oxygen Evolution Reaction. Angew. Chem., Int. Ed. 2014, 53, 7169−7172. (15) Yeo, B. S.; Klaus, S. L.; Ross, P. N.; Mathies, R. A.; Bell, A. T. Identification of Hydroperoxy Species as Reaction Intermediates in the Electrochemical Evolution of Oxygen on Gold. ChemPhysChem 2010, 11, 1854. (16) Nørskov, J. K.; Rossmeisl, J.; Logadottir, A.; Lindqvist, L.; Kitchin, J. R.; Bligaard, T.; Jónsson, H. Origin of the Overpotential for Oxygen Reduction at a Fuel-Cell Cathode. J. Phys. Chem. B 2004, 108, 17886−17892. (17) Tripkovic, V.; Skulason, E.; Siahrostami, S.; Norskov, J. K.; Rossmeisl, J. The Oxygen Reduction Reaction Mechanism on Pt(111) from Density Functional Theory Calculations. Electrochim. Acta 2010, 55, 7975−7981. (18) Karlberg, G. S. Adsorption Trends for Water, Hydroxyl, Oxygen, and Hydrogen on Transition-Metal and Platinum-Skin Surfaces. Phys. Rev. B 2006, 74, 153414. (19) Rossmeisl, J.; Greeley, J.; Karlberg, G. S. Electrocatalysis and Catalyst Screening from Density Functional Theory Calculations. In Fuel Cell Catalysis: A Surface Science Approach; John Wiley & Sons, Inc.: Hoboken, NJ, 2009; pp 57−92. (20) Calle-Vallejo, F.; Martínez, J. I.; Rossmeisl, J. Density Functional Studies of Functionalized Graphitic Materials with Late Transition Metals for Oxygen Reduction Reactions. Phys. Chem. Chem. Phys. 2011, 13, 15639−15643. (21) Liu, S.; White, M. G.; Liu, P. Mechanism of Oxygen Reduction Reaction on Pt(111) in Alkaline Solution: Importance of Chemisorbed Water on Surface. J. Phys. Chem. C 2016, 120, 15288. (22) Briquet, L. G. V.; Sarwar, M.; Mugo, J.; Jones, G.; Calle-Vallejo, F. A New Type of Scaling Relations to Assess the Accuracy of

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4. CONCLUSIONS In this work, we have identified and studied the most stable nondissociated water structures for different amounts of explicit solvent on rutile IrO2(110). We utilized the global optimization scheme minima hopping, in addition to adsorbate swapping, to determine the optimal solvent structure for each level of adsorbate, coverage, and solvent. We find that the nondissociated explicit solvent has the most pronounced effect on binding energies when the adsorbates are capable of donating hydrogen bonds to the solvent, and less so when the adsorbate can only receive hydrogen bonds. Binding energies and descriptor energies seem to be converged with 1−2 bilayers of solvent, mitigating the computational cost of including explicit solvent in one’s analyses, and give a reasonable pathway for determining the effect of solvent on new surfaces. We find that the magnitude of the changes in binding energies is small for all cases except for *OOH on an oxygen covered IrO2(110) surface, which decreases by about 0.4 eV with solvent. This leads to a decrease in the ΔGOOH − ΔGOH descriptor by about 0.3 eV with solvent. The solvent structures determined are consistent between steps in the OER, with slight adjustments to the hydrogen bonding of some of the water molecules to adjust for the changing adsorbate. We estimate a free energy barrier for reorientation of water molecules at 298 K using a calculated reorientation time, suggesting that solvent reorientation is not negligible in this study and hence the energetics presented are more useful for thermodynamic analyses.

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AUTHOR INFORMATION

ORCID

Joseph A. Gauthier: 0000-0001-9542-0988 Colin F. Dickens: 0000-0002-6151-0755 Leanne D. Chen: 0000-0001-9700-972X Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Support from the U.S. Department of Energy, Office of Basic Energy Science, to the SUNCAT Center for Interface Science and Catalysis is gratefully acknowledged. L.D.C. acknowledges financial support from the Natural Sciences and Engineering Research Council of Canada for the CGS-D3 fellowship. C.F.D. acknowledges fellowship support from the National Science Foundation Graduate Research Fellowship (Grant No. DGE114747). A.D.D. would like to acknowledge funding from a Lieberman fellowship through the Vice Provost for Graduate Education at Stanford University.

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REFERENCES

(1) Truong, T. N. Solvent Effects on Structure and Reaction Mechanism: A Theoretical Study of [2 + 2] Polar Cycloaddition between Ketene and Imine. J. Phys. Chem. B 1998, 102, 7877. (2) Michaelides, A.; Ranea, V. A.; de Andres, P. L.; King, D. A. General Model for Water Monomer Adsorption on Close-Packed H


Article

The Journal of Physical Chemistry C Computational Predictions of Catalytic Activities Applied to the Oxygen Evolution Reaction. ChemCatChem 2017, 9, 1261. (23) Baker, N. A.; Sept, D.; Joseph, S.; Holst, M. J.; Mccammon, J. A.; Wolynes, P. G. Electrostatics of Nanosystems: Application to Microtubules and the Ribosome. Proc. Natl. Acad. Sci. U. S. A. 2001, 98, 10037. (24) Mathew, K.; Sundararaman, R.; Letchworth-weaver, K.; Arias, T. A.; Hennig, R. G. Implicit Solvation Model for Density-Functional Study of Nanocrystal Surfaces and Reaction Pathways. J. Chem. Phys. 2014, 140, 084106. (25) Mathew, K.; Hennig, R. G. Implicit Self-Consistent Description of Electrolyte in Plane-Wave Density-Functional Theory. 2016, arXiv:1601.03346 [cond-mat.mtrl-sci]. arXiv.org e-Print archive. https://arxiv.org/abs/1601.03346. (26) Sakong, S.; Groß, A. The Importance of the Electrochemical Environment in the Electro-Oxidation of Methanol on Pt(111). ACS Catal. 2016, 6, 5575−5586. (27) Sha, Y.; Yu, T. H.; Liu, Y.; Merinov, B. V.; Goddard, W. A. Theoretical Study of Solvent Effects on the Platinum-Catalyzed Oxygen Reduction Reaction. J. Phys. Chem. Lett. 2010, 1, 856−861. (28) McCrory, C. C. L.; Jung, S.; Ferrer, I. M.; Chatman, S. M.; Peters, J. C.; Jaramillo, T. F. Benchmarking Hydrogen Evolving Reaction and Oxygen Evolving Reaction Electrocatalysts for Solar Water Splitting Devices. J. Am. Chem. Soc. 2015, 137, 4347−4357. (29) McCrory, C. C. L.; Jung, S.; Peters, J. C.; Jaramillo, T. F. Benchmarking Heterogeneous Electrocatalysts for the Oxygen Evolution Reaction. J. Am. Chem. Soc. 2013, 135, 16977−16987. (30) Kötz, R.; Stucki, S.; Scherson, D.; Kolb, D. M. In-Situ Identification of RuO4 as the Corrosion Product during Oxygen Evolution on Ruthenium in Acid Media. J. Electroanal. Chem. Interfacial Electrochem. 1984, 172, 211−219. (31) Park, S.; Shao, Y.; Liu, J.; Wang, Y. Oxygen Electrocatalysts for Water Electrolyzers and Reversible Fuel Cells: Status and Perspective. Energy Environ. Sci. 2012, 5, 9331. (32) Danilovic, N.; Subbaraman, R.; Chang, K. C.; Chang, S. H.; Kang, Y. J.; Snyder, J.; Paulikas, A. P.; Strmcnik, D.; Kim, Y. T.; Myers, D.; et al. Activity-Stability Trends for the Oxygen Evolution Reaction on Monometallic Oxides in Acidic Environments. J. Phys. Chem. Lett. 2014, 5, 2474−2478. (33) Siahrostami, S.; Vojvodic, A. Influence of Adsorbed Water on the Oxygen Evolution Reaction on Oxides. J. Phys. Chem. C 2015, 119, 1032−1037. (34) Hammer, B.; Hansen, L. B.; No, J. K. Improved Adsorption Energetics within Density-Functional Theory Using Revised PerdewBurke-Ernzerhof Functionals. Phys. Rev. B: Condens. Matter 1999, 59, 7413. (35) Giannozzi, P.; Baroni, S.; Bonini, N.; Calandra, M.; Car, R.; Cavazzoni, C.; Ceresoli, D.; Chiarotti, G. L.; Cococcioni, M.; Dabo, I.; et al. QUANTUM ESPRESSO: A Modular and Open-Source Software Project for Quantum Simulations of Materials. J. Phys.: Condens. Matter 2009, 21, 395502. (36) Bahn, S. R.; Jacobsen, K. W. An Object-Oriented Scripting Interface to a Legacy Electronic Structure Code. Comput. Sci. Eng. 2002, 4, 56−66. (37) Vanderbilt, D. Soft Self-Consistent Pseudopotentials in a Generalized Eigenvalue Formalism. Phys. Rev. B: Condens. Matter Mater. Phys. 1990, 41, 7892−7895. (38) Laasonen, K.; Car, R.; Lee, C.; Vanderbilt, D. Implementation of Ultrasoft Pseudopotentials in Ab Initio Molecular Dynamics. Phys. Rev. B: Condens. Matter Mater. Phys. 1991, 43, 6796−6799. (39) Laasonen, K.; Pasquarello, A.; Car, R.; Lee, C.; Vanderbilt, D. Car-Parrinello Molecular Dynamics with Vanderbilt Ultrasoft Pseudopotentials. Phys. Rev. B: Condens. Matter Mater. Phys. 1993, 47, 10142−10153. (40) Pack, J. D.; Monkhorst, H. J. Special Points for Brillouin-Zone Integrations”A Reply. Phys. Rev. B 1977, 16, 1748−1749. (41) García-Melchor, M.; Vilella, L.; López, N.; Vojvodic, A. Computationally Probing the Performance of Hybrid, Heterogeneous,

and Homogeneous Iridium-Based Catalysts for Water Oxidation. ChemCatChem 2016, 8, 1792−1798. (42) Zhang, H.; Soon, A.; Delley, B.; Stampfl, C. Stability, Structure, and Electronic Properties of Chemisorbed Oxygen and Thin Surface Oxides on Ir(111). Phys. Rev. B 2008, 78, 045436. (43) Bolzan, A. A.; Fong, C.; Kennedy, B. J.; Howard, C. J. Structural Studies of Rutile-Type Metal Dioxides. Acta Crystallogr., Sect. B: Struct. Sci. 1997, B53, 373−380. (44) Goedecker, S. Minima Hopping: An Efficient Search Method for the Global Minimum of the Potential Energy Surface of Complex Molecular Systems. J. Chem. Phys. 2004, 120, 9911. (45) Peterson, A. A. Global Optimization of Adsorbate-Surface Structures While Preserving Molecular Identity. Top. Catal. 2014, 57, 40−53. (46) Gillan, M. J.; Alfè, D.; Michaelides, A. Perspective: How Good Is DFT for Water? J. Chem. Phys. 2016, 144, 130901. (47) Distasio, R. A.; Santra, B.; Li, Z.; Wu, X.; Car, R. The Individual and Collective Effects of Exact Exchange and Dispersion Interactions on the Ab Initio Structure of Liquid Water. J. Chem. Phys. 2014, 141, 084502. (48) Schiros, T.; Naslund, L.-A.; Andersson, K.; Gyllenpalm, J.; Karlberg, G. S.; Odelius, M.; Ogasawara, H.; Pettersson, L. G. M.; Nilsson, A. Structure and Bonding of the Water-Hydroxyl Mixed Phase on Pt(111). J. Phys. Chem. C 2007, 111, 15003−15012. (49) Hummelshøj, J. S.; Luntz, A. C.; Nørskov, J. K. Theoretical Evidence for Low Kinetic Overpotentials in Li-O2 Electrochemistry. J. Chem. Phys. 2013, 138, 034703. (50) Karlberg, G. S.; Rossmeisl, J.; Nørskov, J. K. Estimations of Electric Field Effects on the Oxygen Reduction Reaction Based on the Density Functional Theory. Phys. Chem. Chem. Phys. 2007, 9, 5158. (51) Futera, Z.; English, N. J. Electric-Field Effects on AdsorbedWater Structural and Dynamical Properties at Rutile- and AnataseTiO2 Surfaces. J. Phys. Chem. C 2016, 120, 19603. (52) Halck, N. B.; Petrykin, V.; Krtil, P.; Rossmeisl, J. Beyond the Volcano Limitations in Electrocatalysis - Oxygen Evolution Reaction. Phys. Chem. Chem. Phys. 2014, 16, 13682−13688. (53) Vartia, A. A.; Mitchell-Koch, K. R.; Stirnemann, G.; Laage, D.; Thompson, W. H. On the Reorientation and Hydrogen-Bond Dynamics of Alcohols. J. Phys. Chem. B 2011, 115, 12173−12178. (54) Groß, A.; Gossenberger, F.; Lin, X.; Naderian, M.; Sakong, S.; Roman, T. Water Structures at Metal Electrodes Studied by Ab Initio Molecular Dynamics Simulations. J. Electrochem. Soc. 2014, 161, E3015−E3020. (55) Roman, T.; Groß, A. Structure of Water Layers on HydrogenCovered Pt Electrodes. Catal. Today 2013, 202, 183−190. (56) Limmer, D. T.; Willard, A. P.; Madden, P.; Chandler, D. Hydration of Metal Surfaces Can Be Dynamically Heterogeneous and Hydrophobic. Proc. Natl. Acad. Sci. U. S. A. 2013, 110, 4200−4205. (57) Hansen, H. A.; Viswanathan, V.; Nørskov, J. K. Unifying Kinetic and Thermodynamic Analysis of 2 E − and 4 E − Reduction of Oxygen on Metal Surfaces. J. Phys. Chem. C 2014, 118, 6706−6718.

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Solvation Effects for Oxygen Evolution Reaction Catalysis on IrO

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