Continuous Representation of the Proton and Electron Kinetic

Feb 26, 2015 - Continuous Representation of the Proton and Electron Kinetic Parameters in the pH–Potential Space for Water Oxidation on Hematite...
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Continuous Representation of the Proton and Electron Kinetic Parameters in the pH−Potential Space for Water Oxidation on Hematite Shima Haghighat and Jahan M. Dawlaty* Department of Chemistry, University of Southern California, Los Angeles, California 90089-1062, United States S Supporting Information *

ABSTRACT: Understanding the mechanisms of multielectron and multiproton electrochemical reactions, particularly in the context of solar-to-fuel water splitting, is an outstanding challenge. Historically, Pourbaix diagrams are used to show the influence of potential and pH on the thermodynamic stability of electrode−electrolyte systems. These diagrams do not carry kinetic or mechanistic information, which often restricts their use to cases in which the thermodynamic limit can be assumed. We introduce and construct from experimental data two new types of diagrams that demonstrate the kinetic variations of electrochemical reactions as a function of pH and potential. These diagrams show the variation of the electron-transfer parameter (α) and the proton reaction order (ρ) in a wide range of potential and pH. We present α(pH, E) and ρ(pH, E) for water electrolysis on an iron oxide electrode in the range of pH 7 to 13. In these plots, regions of acidic and basic mechanisms, relationship to surface protonation equilibria, and switching between acidic and basic mechanisms due to electrochemical production of protons can be easily identified. The proton reaction order is zero in the acidic side, while it is nonzero in the basic limit. A larger empirical electron-transfer parameter is observed in the basic compared to the acidic region. These observations are related to the differences in oxidation mechanism between the two regions. We propose the use of such diagrams to gain an expanded and enhanced view of the kinetics of multielectron and multiproton electrochemical reactions.

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INTRODUCTION Understanding the kinetics of the natural and artificial water splitting reaction is important for future sunlight-to-fuel lightharvesting schemes.1−5 In the net reaction, four electrons and protons are removed from two water molecules to make one oxygen molecule. Despite years of research, the kinetic details of this reaction are poorly understood and remain an area of active study. Metal oxide electrodes are promising catalysts for electrochemical and photoelectrochemical oxidation of water.6−10 Purely inorganic metal oxide films offer substantial practical advantages over synthetic organometallic complexes because they can be manufactured in large scale, at lower cost, and with a smaller environmental footprint. The challenges of understanding the kinetics of water oxidation on metal oxide surfaces, at least in part, are due to their simultaneous dependence on pH and potential, the surface chemistry, and surface electrostatics.11−17 The thermodynamics of electrode−electrolyte systems are commonly depicted in the widely used Pourbaix diagrams.18 They are interpreted as thermodynamic phase diagrams and identify regions of stability of various oxidation forms of a substance in the potential−pH parameter space. It needs no emphasis that Pourbaix diagrams do not carry kinetic information. To convey the kinetic information, the electrochemical current can be plotted as a function of pH and potential I(pH, E) in 2-dimensional contour plots, which have © 2015 American Chemical Society

been recently named dynamic potential−pH diagrams (DPPDs).19 In the current work, we demonstrate continuous representation of two fundamental kinetic quantities, the empirical electron-transfer coefficient (α) and the empirical reaction order with respect to protons (ρ), in a large range of pH and potential.20−22 The electrochemical current is a function of two experimentally controlled parameters, the activity of protons in the electrolyte (aH+) and the applied electrode potential (E): I ∝ a Hρ + exp(αFE /RT )

(1)

where ρ and α are known as the reaction order with respect to protons and the empirical electron-transfer parameter, respectively, and can be expressed as ⎛ ∂log I ⎞ ⎛ ∂log I ⎞ ρ = −⎜ ⎟ = −⎜ ⎟ ⎝ ∂( −log a H+) ⎠ E ⎝ ∂pH ⎠ E α = 2.3

RT ⎛ ∂log I ⎞ ⎜ ⎟ F ⎝ ∂E ⎠ pH

(2)

(3)

Received: January 3, 2015 Revised: February 22, 2015 Published: February 26, 2015 6619

DOI: 10.1021/acs.jpcc.5b00053 J. Phys. Chem. C 2015, 119, 6619−6625

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The Journal of Physical Chemistry C Both α and ρ depend on the mechanism of the reaction.20,23,24 The central emphasis of this work is to calculate from experimental data the variation of α and ρ in a large range of pH and potential. In contrast to the Pourbaix diagrams, plots of α(pH, E) and ρ(pH, E) give global and insightful maps of electrochemical reaction kinetics, rather than just thermodynamic variations. The reaction order ρ(pH, E) is a measure of the sensitivity of the reaction rate on pH and can, in principle, be obtained by measuring current while varying pH and holding potential constant.20,21 Experimentally, it is more convenient to scan the potential E at constant pH values and acquire a number of I−E curves, each for a fixed pH, and evaluate ρ(pH, E) from the data according to eq 2. The parameter α(pH, E) is related to the commonly measured Tafel slope b as α = 2.3(RT/F)b−1.21 For reference, the value of α = 0.5 corresponds to the Tefel slope of 119 mV/decade at T = 298 K. To our knowledge, such a representation of the water oxidation kinetics has not been performed before.

then was closed to the ambient air. Because reaction order is a derivative, its continuous representation is limited by the number of discrete points along the pH axis. In our experiment, current−voltage curves were obtained for 14 values of pH between pH 7 and pH 13. The indicated potentials are always referenced to the normal hydrogen electrode. Linear sweep voltammetry experiments were carried out at a 10 mV/s scan rate, between 0 and 2.4 V versus Ag/AgCl electrode. The resistance of the electrolyte solution and the electrode film was evaluated using a common method as outlined by Krstajic and Trasatti.28 When current I flows through the cell, this uncompenstated resistance (Ru) contributes to the measured voltage by an amount IRu. Hence, at large currents, the measured voltage is influenced by the uncompensated resistance giving rise to linear I−E curves. The uncompensated resistance appears as the intercept of a plot of dE/dI versus 1/I, where dE and dI are the change in potentail and current for two consecutive data points, respectively, and I is the mean current for the same two data points. The value of Ru was determined to be ∼70 Ω and changed little for various pH values. Small variations of the ionic strength did not yield any significant change in the I−E curves, pointing to the fact that the uncompensated resistance mostly arises from the film. To account for this resistance, for each pair of I−E data, the value IRu was subtracted from the voltage V, giving rise to the corrected voltage. For completeness, all of the raw data and the kinetic parameters calculated based on them are presented in the Supporting Information. We emphasize, however, that it is best to infer chemical mechanisms from the resistancecompensated I−E curves and the corresponding kientic parameters. Only regions of low current (up to ∼200 mV above the onset potential) in the raw data carry information about the chemical kinetics at the surface. Because of resistance compensation, the range of the corrected voltage axis is determined by the measured current and hence varies for different pH values, as shown in Figure 1.

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EXPERIMENTAL SECTION Thin Film Preparation. The iron oxide thin films were prepared by the photochemical metal organic deposition (PMOD) thechnique.25,26 In this method, a photosensitive precursor solution is spin-coated onto a substrate and then exposed to UV radiation. A 15% w/w solution of Iron(III)ethylhexanoate (50% in mineral spirits, Alfa Aesar) in hexane was prepared using 0.058 g of Iron(III)ethylhexanoate and 0.142 g of hexane. Flourin-doped tin oxide (FTO) coated glass (MTI Corporation, TEC70) substrates were cleaned ultrasonically in acetone followed by methanol then dried under air streaming. A few drops of the solution were dispensed on the FTO by spin coating at a rate of 3000 rpm for 1 min. The film was irradiated under an ultraviolet (UV) lamp (254 nm) for 10 h until removal of organic ligands was more than 99% complete, leaving an amorphous metal oxide film on the substrate. Fourier transfrom infrared (FT-IR) spectroscopy was performed on the thin film deposited on glass substrates to ensure elimination of the 2-ethylhexanoate signature vibrational peaks (see Figure 1 in Supporting Information). The thin film was annealed at T = 600° C for 1 h in air using a Vulcan A-550 oven. The film was characterized by X-ray diffraction (XRD) and energy dispersive X-ray spectroscopy (EDX), as shown in the Supporting Information (Figures 2 and 3 in Supporting Information), and two characteristic peaks for α-hematite were observed in conformity with previous literature.25,27 The film thickness was measured by profilometry and was determined to be ∼50 nm. The electrode was prepared by contacting the FTO layer with a copper wire and silver paint in a region that was deliberately not covered by the hematite film. The contact region was fully covered and secured using epoxy glue. Electrochemical Measurements. Electrochemical data were obtained using a Gamry Reference 3000 potentiostat and a three-electrode cell. A Ag/AgCl reference electrode (Ag/ AgCl vs normal hydrogen electrode (NHE) = +197 mV) and Pt mesh counter electrode were used in all the experiments. The electrolyte solution was 0.1 M KOH, and its pH was adjusted by adding a dilute HCl solution. The change in the ionic strength of the solution was estimated from the number of added drops of HCl and varied from 0.1 M to ∼0.2 M over the pH range of interest. To avoid influences from the specific adsorption of multivalent anions, no agent for pH buffering was added. The cell was purged with N2 gas prior to each scan and

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RESULTS AND DISCUSSION The resistance-compensated I−E curves for pH values ranging from 7 to 13 are shown in Figure 1. Although differences in the

Figure 1. Measured linear sweep voltammograms after compensation for film resistance at different pH values from pH 7 to pH 13. Although differences in the behavior of I−E curves as a function of pH exist, it is hard to visualize the influence of pH using this representation. This is better addressed in the two-dimensional representation in Figure 2. 6620

DOI: 10.1021/acs.jpcc.5b00053 J. Phys. Chem. C 2015, 119, 6619−6625

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The Journal of Physical Chemistry C behavior of I−E curves between the acidic and basic sides are visible, it is hard to quantify the pH dependence of the I−E curves from this figure. In particular, the differences in the slop of I−E curves, and the relatively flat region between pH ∼ 10− 12, demand explanation. For that reason, we prepare the twodimensional representation of the I−E curves as a function of potential and pH. The plot of log I(pH, E) for water oxidation on a hematite thin film electrode is shown in Figure 2. The

Figure 3. Proton reaction order plot ρ(pH, E) for water splitting on the surface of iron oxide. The reaction order shows the sensitivity of the electrochemical current to changes in pH as defined in eq 2. In the acidic region, the electrochemical current, and correspondingly the water oxidation reaction rate, is practically insensitive to pH, while for the basic mechanism ρ ≠ 0. The largest sensitivity to pH is observed in the region where a switch between the two mechanisms is expected. This is explained in the text using a surface protonation model.

Figure 2. Electrochemical current log I(pH, E) for water splitting on the surface of iron oxide as a function of pH and potential. The Pourbaix line for water stability (solid line) is overlaid for reference. The double arrows indicate the overpotential for water oxidation and demonstrate its variation as a function of pH. The hashed region corresponds to no data because the compensated voltage axis is different for various values of pH (see Experimental Section).

thermodynamic line for water stability (the Pourbaix line) with the Nernstian slope of 59 mV/pH is overlaid on the plot for reference. An immediate insight can be already gleaned from this diagram. While on the basic side the onset potential follows a Nernstian slope, on the acidic side it does not show any dependence on pH. This feature, along with others, will be explained shortly in the context of surface protonation equilibrium and differences in electron-transfer parameter in the basic and acidic sides. We emphasize that in this work the words acidic and basic do not mean below and above pH 7, but rather refer to the two limits of the pH range that we have studied (pH 7 and pH 13). We construct the kinetic diagrams by taking the gradient of the log I along the pH and potential axes and calculate α(pH, E) and ρ(pH, E) according to eqs 2 and 3. In calculating the derivatives, we set the onset current to 0.1 mA and ignore the regions of low overpotential where the current is too low for a derivative to be meaningful. The resulting plots are shown in Figures 3 and 4. Now we will interpret the observations in the diagrams beginning with the proton reaction order plot in Figure 3. The data clearly shows that at low pH values, the electrochemical current has no sensitivity to pH (ρ = 0). At high pH values, a clearly nonzero value of ρ is observed, indicating that the reaction mechanism in the basic region is different from that in the acidic region. The highest magnitude of reaction order is observed in the intermediate region (ρ ≈ −1.2) and corresponds to the region where a switch between the two mechanisms is expected. To understand these observations, first we briefly review a model that accounts for both surface and bulk contributions of protons to the reaction rate. Then we

Figure 4. Electron-transfer coefficient plot α(pH, E) revealing mechanistically distinct regions. The electron-transfer coefficient (α) is a dimensionless measure of the sensitivity of the reaction rate to applied potential, as shown in eq 3. A noteworthy feature is that α in the basic region is distinctly larger than that in the acidic region. The dotted arrow in the intermediate region shows a reaction that starts with the basic mechanism, but because of electrochemical generation of protons near the interface, switches to acidic mechanism, displaying a characteristic smaller α. These features are further explained in the text.

explain the observed variation of ρ based on this model. It is known that even nonadsorbed molecules within a few monolayers of an interface do not behave like bulk molecules and often assume anisotropic structure. This work is not concerned with distinguishing that type difference, and the word bulk here is used to refer to all nonadsorbed species. The rate of a reaction that consumes an adsorbed species is expected to depend on the coverage function (θ) of the adsorbed species, rather than its bulk concentration. In this case, as in conventional kinetics, a rate law of the form I ∝ θn can be written, where n depends on the stoichiometry of the adsorbed species.21,29,30 For example, if a reaction consumes adsorbed OH− with a stoichiometry of 1, its rate will depend on 6621

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The Journal of Physical Chemistry C surface coverage of OH− to the first power. The coverage θ, in turn, is a function of the bulk concentration of that species via the commonly used adsorption functions, such as the Langmuir or Temkin isotherms.30 An adsorption function θ(pH) that shows the variation of surface coverage of a species with respect to bulk pH can be used to rewrite eq 2 as ⎛ ∂log I ⎞ ⎛ ∂(log θ) ⎞ ρads = −⎜ ⎟ ⎜ ⎟ ⎝ ∂(log θ) ⎠ E ⎝ ∂pH ⎠ E

(4)

The above is best understood in the limiting cases of adsorption. When the surface is completely saturated with an adsorbed species (θ → 1, as in the flat region of the Langmuir isotherm), any small change in the bulk pH will not change the surface coverage significantly. The second factor in the above equation, and consequently ρ, vanishes. For example, in the pH range where the surface is fully covered with OH−, a differential change in the bulk pH will not influence the surface coverage of OH− any further. In the limit of low coverage θ → 0, lateral interactions between adsorbates vanish and adsorption models converge to the low-coverage region of the Langmuir isotherm. In this limit the coverage θ increases linearly with increasing bulk concentration and the second factor in eq 4 becomes unity. For intermediate coverage, a sublinear dependence of coverage on concentration is expected. In that case, the value of the second factor is between 0 and 1 and can, in principle, be determined from the adsorption isotherm model. When the reaction involves both an adsorbed species x and the same species in the bulk (such as the reaction of bulk OH− with adsorbed OH−), the rate law will depend on both the surface coverage θ and the bulk concentration [x] as I ∝ θn[x]m. In this case, the reaction order is ρ = ρads + ρbulk (5)

Figure 5. Hypothesized paths for two-electron, two-proton generation of surface MO, where M represents a surface active site (i.e., an iron ion). At zero applied potential, the pH determines the extent of the equilibrium MOH2 ⇆ MOH− and hence the starting ingredients for oxidation. These paths provide a reasonable explanation of the observed α and ρ in the acidic and basic limits.

water. Transfer of electrons into the electrode renders the surface bound proton acidic, such that the proton is immediately lost to the bulk deprotonating species. A fully concerted pathway means that as the electron transfer approaches the transition state, concurrently the hydrogen bonding of the proton to the bulk species strengthens and the proton transfer also approaches transition state. A second electron transfer is also associated with proton transfer, giving rise to the surface MO species. At low pH range, this path is not sensitive to the initial adsorbed MOH2 coverage because the surface is fully covered and any small change in pH will not change the surface coverage appreciably. The second term in eq 4 is zero in this case, and the reaction rate loses its sensitivity to the bulk pH changes ρads = 0. Because bulk water is the deprotonating agent, no bulk reaction order will arise and an overall proton reaction order of ρ = 0 is expected. This is exactly observed in the acidic region of the plot of ρ(pH, E) in Figure 3. In the basic regime, the surface is fully hydroxylated at zero applied potential. Only the first electron transfer is associated with deprotonation, and the second electron transfer yields MO. In this regime, the deprotonating agent is a bulk hydroxide. Hence, even when the surface is fully saturated with hydroxyls and ρads = 0, a nonzero reaction order with respect to bulk hydroxides or bulk pH is expected. This matches our observations, in which ρ is distinctly nonzero just above the onset potential in Figure 3 in the basic region. We also note that the slope of the onset potential in the basic range is very close to the Nernstian value of 59 mV/pH, suggesting that bulk hydroxides play a role in lowering the kinetic barrier for the basic path. In the intermediate pH values, the surface protonation equilibrium (eq 6) is not favored heavily to either of the two extremes. The surface is covered with both hydroxyls and adsorbed water. Unlike the pH extremes, the second factor in eq 4 is not zero; thus, a large ρads is expected. In fact, a differential change in the pH significantly changes the population of the surface adsorbed species and tips the balance between acidic and basic mechanisms easily. Therefore, a large sensitivity to pH is expected in this range, which is indeed observed in our data as ρ ≈ −1.2 in Figure 3.

where ρads is given by eq 4 and ρbulk = m. For a metal oxide in equilibrium with an aqueous electrolyte, the following surface protonation equilibrium is justified in the absence of potential: MOH 2 ⇆ MOH− + H+

(6)

where M stands for a surface active site, namely an available iron ion.31 At low (high) pH values, the equilibrium is shifted far to the left (right) and the surface is covered completely with adsorbed water (hydroxyls). Thus, the initial ingredients of the oxidation process at low and high pH are adsorbed water and adsorbed hydroxyls, respectively. Figure 5 presents a scheme that follows the fate of the adsorbed water or hydroxyl species as they lose electrons and protons and become surface MO, which is considered an important intermediate of water oxidation. We point out that various proposed schemes for water oxidation mechanisms in the acidic and basic limits have existed for several decades.12,13,32,33 The scheme in Figure 5 is the one that reasonably explains both the α and ρ data observed in this work, both in the acidic and basic limits, as will be shown shortly. In the hypothesized acidic pathway, as the potential is increased and an electron is transferred into the electrode, the adsorbed water (MOH2) is concurrently deprotonated. Thus, the generation of the highly acidic and unstable MOH2+ is avoided. Such a concerted electron−proton transfer (CPET) step has been hypothesized for heterogeneous reactions previously.34−38 Prior to electron transfer, a surface bound proton is likely already hydrogen bonded to a bulk hydroxyl or 6622

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interface, the interfacial pH gradually changes to acidic, and the intermediate behavior is observed. Finally, after a certain potential, the reaction rate is large enough that the produced protons near the surface enforce a fully acidic mechanism. The above description highlights the importance of measuring both α and ρ over a wide range of potential and pH. We emphasize that mechanistic studies of multielectron and multiproton reactions benefit greatly from measuring both α and ρ. We point out, however, that such an approach should only be taken as complementary to other mechanistic studies, such as in situ measurements of intermediates, to disentangle the challenging multielectron and multiproton oxidation problem. Finally, we comment on the influence of mass transfer and adsorbed anions from electrolyte on kinetic studies. First, in the regions of large current, mass transfer of ions limits the current and complicates kinetic studies. The current no longer obeys an exponential relation with respect to potential, as in eq 1. We have compensated the measured voltage for resistance and have identified the electrode thin film as the main contributor to the resistance. However, the influence of mass transfer is still identifiable in the reported data, especially in the intermediate pH range. The protons that are released as reaction products acidify the interface to pH values far lower than that of the bulk. This enforces an effective acidic mechanism with the characteristic smaller α, as observed for large potentials in Figure 4. As the bulk pH is made more basic, surface acidification due to the reaction is more easily compensated by bulk hydroxides. Thus, at higher pH, a larger electrochemical current will be required to overwhelmingly acidify the interface. This is evident in the boundary between the acidic and the transition region in Figure 4, which shifts to larger potentials as the pH is raised. We have performed experiments under stirring conditions (see Figure 5 in Supporting Information). In that case, we observe a shift of the intermediate region to lower pH by about 0.5 pH, with all of the salient features of other regions remaining the same. This observation is consistent with the fact that any finite current acidifies the region near the interface with respect to the bulk, resulting into an effective shift of the pH axis of the diagram. Thus, the interpretation of the horizontal axis of such diagrams should always be performed with the experimental conditions in mind. The second point is about the influence of the electrolyte and buffer for the experiment. It is well-known that the choice of buffer can influence the mechanism either via adsorbed buffer anions (specific adsorption) or bulk buffer anions acting as proton acceptors or donors.15,19,44 In this study, we have chosen a relatively simple electrolyte composed of KOH and HCl to adjust the pH from high values to lower values, and our discussions apply only to this choice of electrolyte. We have also performed experiments with multivalent anions (phosphate) electrolytes. In that case, some of the essential features of the diagrams remain the same. In particular, we have observed the 59 mV/pH slope for the onset potential in the basic side and that the coefficient α in the basic side is larger that of the acidic side. However, the ranges and shapes of the mechanistic regions differ from the case of the KOH and HCl system (see Figure 6 in Supporting Information). Their interpretation requires accounting for adsorption of the buffer anions on the surface and the buffer protonation equilibrium. It is a matter of future studies.

Now we turn to the empirical electron-transfer plots. For a single-electron process, the electron-transfer parameter is also known as the symmetry factor β. It is interpreted as the proportionality factor that relates the thermodynamic stabilization of the transferred electron in the electrode to the electrontransfer barrier height (i.e., the energy of the transition state). For example, β = 0.5 means that if the electrode potential is raised by 0.1 V, the electron-transfer barrier will reduce by 0.05 V. It has been given a relatively simple geometric interpretation, which is related to the crossing angle of the free energy surfaces for electron transfer.29 While this is reasonable for singleelectron processes, the potential dependence of multielectron processes is more complicated. A multistep kinetic model20,29,39 for multielectron-transfer processes assumes a distinct ratedetermining step (RDS) and pre-equilibrium prior to the RDS. Using standard approaches, the electron-transfer parameter is derived as α = γ/ν + rβ. Here, γ is the number of electrons transferred prior to the RDS, ν the number of times that the RDS is repeated per catalytic cycle, r the number of electrons transferred at the RDS, and β the familiar symmetry factor. Although the above relation for α has been known for several decades, it has not yielded unique and unambiguous identification of the mechanism of complicated reactions such as water oxidation. The above relation still holds qualitative value, and we will use it in that capacity. If we assume that the RDS is the oxygen−oxygen bond formation, as is suggested in several works,40,41 then both the acidic and basic pathways will have the same number of steps prior to the RDS and will have the same γ and the same ν. Then the observed differences in α may be attributed to differences in the symmetry factor β for the acidic and basic regions. Several previous works38,42,43 suggest that the electrontransfer parameter for a concerted electron−proton transfer reaction is different from that of single electron transfer. In particular, smaller values of α were measured and justified for the CPET process.42 Although it is quite desirable, the exact simulation of the α parameter for the concerted and stepwise mechanisms is beyond the scope of this work. It is reasonable to postulate that when a single step involves both electron and proton transfer, the transition state will correspondingly have partial electron- and proton-transfer character. Variation of the electrode potential has a larger influence on the electrontransfer character of the transition state. Thus, the transition state of one electron transfer is expected to be more sensitive to electrode potential variation, compared to the transition state of a step that involves both electron and proton transfer. Therefore, a smaller β is expected for the steps that involve both electron and proton transfer, compared to the steps that are only electron transfer. In the acidic regime, two electron− proton transfer steps are required, while in the basic regime only one of the steps is combined proton and electron transfer. Correspondingly, a smaller α is expected in the acidic side compared to the basic side. Our observations are in conformity with this explanation, as shown in Figure 4. A smaller value of α ≈ 0.1−0.2 is observed in the acidic region, while in the basic region α is as high as α ≈ 0.7. A valuable feature of the continuous representation of α and ρ in the two-dimensional plot is easy identification of the transition in mechanism between the two regimes. This is most prominently seen in the intermediate pH values near pH 10.5. In this region, as the potential is scanned to higher values (see dotted line in 4), the reaction starts out with the basic mechanism. Because the reaction produces protons near the 6623

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(5) Walter, M. G.; Warren, E. L.; McKone, J. R.; Boettcher, S. W.; Mi, Q.; Santori, E. A.; Lewis, N. S. Solar Water Splitting Cells. Chem. Rev. (Washington, DC, U.S.) 2010, 110, 6446−6473. PMID: 21062097. (6) Trasatti, S. Transition Metal Oxides: Versatile Materials for Electrocatalysis. In The Electrochemistry of Novel Materials; Lipkowski, J., Ross, P., Eds.; Frontiers of Electrochemistry; VCH: New York, 1994. (7) Liao, P.; Keith, J. A.; Carter, E. A. Water Oxidation on Pure and Doped Hematite (0001) Surfaces: Prediction of Co and Ni as Effective Dopants for Electrocatalysis. J. Am. Chem. Soc. 2012, 134, 13296− 13309. (8) Zhao, P.; Kronawitter, C. X.; Yang, X.; Fu, J.; Koel, B. E. WO3-αFe2O3 Composite Photoelectrodes with Low Onset Potential for Solar Water Oxidation. Phys. Chem. Chem. Phys. 2014, 16, 1327−1332. (9) Matsumoto, Y.; Sato, E. Electrocatalytic Properties of Transition Metal Oxides for Oxygen Evolution Reaction. Mater. Chem. Phys. 1986, 14, 397−426. (10) Klepser, B. M.; Bartlett, B. M. Anchoring a Molecular Iron Catalyst to Solar-Responsive WO3 Improves the Rate and Selectivity of Photoelectrochemical Water Oxidation. J. Am. Chem. Soc. 2014, 136, 1694−1697. PMID: 24437445. (11) Conway, B. E.; Gileadi, E. Kinetic Theory of PseudoCapacitance and Electrode Reactions at Appreciable Surface Coverage. Trans. Faraday Soc. 1962, 58, 2493−2509. (12) Damjanovic, A.; Dey, A.; Bockris, J. Kinetics of Oxygen Evolution and Dissolution on Platinum Electrodes. Electrochim. Acta 1966, 11, 791−814. (13) Bockris, J. O. Kinetics of Activation Controlled Consecutive Electrochemical Reactions: Anodic Evolution of Oxygen. J. Chem. Phys. 1956, 24, 817−827. (14) Takashima, T.; Hashimoto, K.; Nakamura, R. Mechanisms of pH-Dependent Activity for Water Oxidation to Molecular Oxygen by MnO2 Electrocatalysts. J. Am. Chem. Soc. 2012, 134, 1519−1527. (15) Kushner-Lenhoff, M. N.; Blakemore, J. D.; Schley, N. D.; Crabtree, R. H.; Brudvig, G. W. Effects of Aqueous Buffers on Electrocatalytic Water Oxidation with an Iridium Oxide Material Electrodeposited in Thin Layers from an Organometallic Precursor. Dalton Trans. 2013, 42, 3617−3622. (16) Klahr, B. M.; Hamann, T. W. Current and Voltage Limiting Processes in Thin Film Hematite Electrodes. J. Phys. Chem. C 2011, 115, 8393−8399. (17) Jordan, D. S.; Hull, C. J.; Troiano, J. M.; Riha, S. C.; Martinson, A. B. F.; Rosso, K. M.; Geiger, F. M. Second Harmonic Generation Studies of Fe(II) Interactions with Hematite (α-Fe2O3). J. Phys. Chem. C 2013, 117, 4040−4047. (18) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions; National Association of Corrosion Engineers: Houston, TX, 1974. (19) Minguzzi, A.; Fan, F.-R. F.; Vertova, A.; Rondinini, S.; Bard, A. J. Dynamic Potential-pH Diagrams Application to Electrocatalysts for Water Oxidation. Chem. Sci. 2012, 3, 217−229. (20) Bockris, J., Reddy, A. Modern Electrochemistry: An Introduction to an Interdisciplinary Area; Springer: Boston, MA, 1973; Vol. 2. (21) Gileadi, E. Electrode Kinetics for Chemists, Chemical Engineers, and Materials Scientists; VCH: New York, 1993. (22) Bamford, C., Tipper, C., Compton, R. Electrode Kinetics: Principles and Methodology; Comprehensive Chemical Kinetics Series; Elsevier Science: New York, 1986. (23) Trasatti, S., Lodi, G. Oxygen and Chlorine Evolution at Conductive Metallic Oxide Anodes; Studies in Physical and Theoretical Chemistry; Elsevier: Amsterdam, 1981. (24) Angelinetta, C.; Falciola, M.; Trasatti, S. Heterogenous AcidBase Equilibria and Reaction Order of Oxygen Evolution on Oxide Electrodes. J. Electroanal. Chem. Interfacial Electrochem. 1986, 205, 347−353. (25) Smith, R. D. L.; Prévot, M. S.; Fagan, R. D.; Zhang, Z.; Sedach, P. A.; Siu, M. K. J.; Trudel, S.; Berlinguette, C. P. Photochemical Route for Accessing Amorphous Metal Oxide Materials for Water Oxidation Catalysis. Science 2013, 340, 60−63.

CONCLUSION We have evaluated the electron-transfer parameter (α) and the proton reaction order (ρ) as continuous variables in the pH and potential space for water oxidation on hematite. This treatment of α and ρ is useful and insightful because it gives a global view of the sensitivity of the reaction to potential and pH. Simultaneous consideration of ρ and α sheds light on the basic and acidic pathways of formation of surface MO species, which is considered a precursor for oxygen bond formation. The evaluated ρ and α for the acidic and basic regions are reasonably justified for the two pathways shown in Figure 5. We hope that the hypothesized pathways inspire future experimental studies such as surface-specific sum-frequency generation, in situ extended X-ray absorption fine structure analysis, and transient FTIR measurements for identification of intermediates. We anticipate that our work will inspire further investigation in two directions. First, we hope that continuous representation of kinetic parameters, especially in the potential−pH space, will be used more extensively in investigations of heterogeneous catalysis reactions as complementary tools. Putative mechanisms should be checked to consistently explain observed values of α and ρ over a large range of potential and pH. Second, we hope that theoretical models of surface equilibria, such as adsorption isotherms, can benefit from the insight provided by these diagrams.

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ASSOCIATED CONTENT

S Supporting Information *

Infrared spectra of a thin film of Fe(III) 2-ethylhexanoate subjected to UV irradiation; X-ray diffraction pattern of FTO substrate and iron oxide thin film on FTO; EDX spectrum of hematite thin film on silicon substrate; linear sweep voltammograms of hematite thin film at different pH values; kinetic diagrams calculated from the raw data; kinetic diagrams under stirring condition; and kinetic diagrams for an experiment in which phosphoric acid, instead of hydrochloric acid, was used for pH adjustment of the electrolyte. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Tel: 213-740-9337. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge support from the University of Southern California start up grant and the AFOSR YIP Award (FA9550-13-1-0128). S.H. was supported by the Mork Family Scholarship for part of the duration of this work.

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REFERENCES

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