or Two-Electron Water Oxidation, Hydroxyl Radical, or H2O2 Evolution

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One- or Two-electron Water Oxidation, Hydroxyl Radical or H2O2 Evolution Samira Siahrostami, Guo-Ling Li, Venkatasubramanian Viswanathan, and Jens K. Norskov J. Phys. Chem. Lett., Just Accepted Manuscript • DOI: 10.1021/acs.jpclett.6b02924 • Publication Date (Web): 23 Feb 2017 Downloaded from http://pubs.acs.org on February 24, 2017

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The Journal of Physical Chemistry Letters

One- or Two-electron Water Oxidation, Hydroxyl Radical or H2O2 Evolution Samira Siahrostami,[1] Guo-Ling Li,[2,3] Venkatasubramanian Viswanathan,[4] Jens K. Nørskov*[1,2]

[1] SUNCAT Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, 443 Via Ortega, Stanford, California 94305, United States [2] SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, 2575 Sand Hill Road, Menlo Park, California 94025, United States [3] School of Physics and Engineering, Henan University of Science and Technology, Luoyang 471023, China [4] Department of Mechanical Engineering, Carnegie Mellon University, 5000 Forbes Ave, Pittsburgh, PA 15213, USA

* email: [email protected] Abstract Electrochemical or photo-electrochemcial oxidation of water to form hydrogen peroxide (H2O2) or hydroxide radicals (•OH) offers a very attractive route to water disinfection, and the first process could be the basis for a clean way to produce hydrogen peroxide. A major obstacle in the development of effective catalysts for these reactions is that the electrocatalyst must suppress the thermodynamically favored four-electron pathway leading to O2 evolution. We develop a thermochemical picture of the catalyst properties that determine selectivity towards the one, two, and four electron processes leading to •OH, H2O2, and O2.

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Water can be oxidized electrochemically or photo-electrochemically to form OH radicals, H2O2 or O2. The number of electrons transferred characterizes the three half-cell reactions leading to these products: The one-electron process: → • ( ) + ( + )

° = 2.73 1

(1)

° = 1.76 2

(2)

° = 1.23 V 2

(3).

The two-electron process: 2 → + 2( + ) The four-electron process: → + 4( + )

Where the standard reduction potentials are reported reference to the reversible hydrogen electrode. In the present letter we present a thermodynamic analysis of the free energy of intermediates adsorbed on different catalyst surfaces for the three reactions. This leads us to a set of criteria characterizing catalysts with a tendency for selectivity towards the different products. The approach rationalizes a number of observations in the literature, and provides a set of catalyst selection rules. The background for the present study is the desire to find electrocatalysts for direct production of hydrogen peroxide (H2O2) in a process that is cleaner and more sustainable than the presently used anthraquinone process.3 Hydrogen peroxide is an important chemical with a wide range of applications in industry including paper and textile manufacturing. It is also a very clean oxidant in water treatment.3 Direct synthesis of H2O2 from its elements hydrogen (H2) and oxygen (O2) has been studied extensively in order to reduce the danger of explosion and simultaneously find selective and active catalysts.4–9 An electrochemical route based on solar or wind energies that could be implemented at the point of use would be very desirable for applications such as water cleaning. Electrochemical reduction of molecular oxygen to hydrogen peroxide has shown considerable promise, the main challenges being to 2


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find a cheap, efficient, and selective catalyst.10–16 The oxidation of water to hydrogen peroxide (eq. 2), is the simplest possible process. It requires only water as the reactant and directly produces H2O2 in addition to H2.17,18,19 Unfortunately, it has proven difficult to find good catalysts for this reaction. There are experimental reports on the electrochemical oxidation of water to hydrogen peroxide over manganese oxide.17,18 Hydroxyl radicals, hydrogen peroxide, and superoxide anions have been detected at the surface of titanium dioxide under UV irradiation.20–24 Small amounts of hydroxyl radicals have also been detected over bismuth vanadate (BiVO4) exposed to visible light.

25–28

Most recently, Fuku et al. have reported a

successful example of a photoelectrode system made of tungsten trioxides (WO3) and bismuth vanadate capable of producing and accumulating hydrogen peroxide effectively.29 Tin oxide (SnO2) has been reported for electrochemical wastewater treatment30,31 and has been proposed as a possible electrocatalyst for H2O2 generation.19 In the following we use potential dependent free energy diagrams for the three reactions (1)-(3) to analyze the selectivity trends. Figure 1 shows examples of the free energy diagrams for three representative catalysts, TiO2, WO3 and IrO2 and with the three different major products included, the OH radical, H2O2 and O2, respectively. The free energy diagrams have been constructed for the adsorption free energies of relevant intermediates of the one- (eq. 1), two- (eq. 2) and fourelectron (eq. (3)) water oxidation reactions, i.e. OH*, O* and OOH*. We use the computational hydrogen electrode (CHE) model, which exploits that the chemical potential of a proton-electron pair is equal to that of gas-phase H2, at U=0.0 V vs. the reversible hydrogen electrode. The effect of the electrode potential on the free energy of the intermediates is taken into account by shifting the electron energy by –eU where e and U are the elementary charge and the electrode potential.32 For each catalyst, the adsorption free energy of the different key intermediates has been taken from reported density functional theory calculations using the RPBE functional to describe exchange and correlation effects including corrections for zero point energies and entropy contributions based on the harmonic approximation.16,19,33,34 The energies of the final states, solvated •OH (details in 3



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Supporting Information), solvated H2O2, and gas phase O2 are taken from experiment.35–37 The thermodynamic analysis can only be taken as qualitative, since it does not include activation barriers, but is has proven useful in rationalizing trends for a number of electrochemical reactions involving oxygen.33,38,39 Consider first the free energy diagram for TiO2, Figure 1(a). Here the largest free energy step in the process is the first oxidation step to form adsorbed OH (OH*). In fact this step is so close in energy to the solvated hydroxyl that once *OH is formed it is equally favorable to desorb and form solvated hydroxyl, •OH(aq) or recombine at the surface and form H2O2. This analysis suggests that TiO2 should have a propensity to form OH radicals, and that the potential needed to apply to make that possible is quite large, U•OH~2.4 . This is in good agreement with observations for UV illuminated titania.20–23 Several reports suggest that OH radicals are formed, and the energy of the holes in TiO2 during UV radiation corresponds to roughly a potential of 2.5V.20–23,40 We note that the free energy of O* is considerably higher than that of H2O2 suggesting that hydrogen peroxide is a likely additional oxidation product. This is in agreement with some experimental reports on detecting small amount of H2O2 under UV illumination of TiO2.20–23” The analysis above can be generalized to suggest that a criterion for forming hydroxyl radicals is that the free energy of OH* on the catalyst surface is higher (more endergonic relative to water) than the free energy of •OH(aq). eU.OH~2.4 . This sets an upper limit for the free energy of OH* at the surface for H2O2 evolution. In other words, the minimum requirement to avoid the one-electron oxidation reaction (eq. 1) and proceed through the two- electron oxidation reaction (eq. 2) is ∆ ≲ 2.4 . Figure 1(b) shows the free energy diagram for IrO2(110), which is known both experimentally41–43 and theoretically33,38 to be an excellent catalyst for O2 evolution. IrO2 strongly adsorbs OH* 33 and the free energy of OH* is far below the free energy for OH radical formation. In addition the free energy of O* is far below the one for H2O2 (∆ = 3.52 ), suggesting that the IrO2 surface has a strong driving force for the complete four-electron oxidation.

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Overall, based on the thermodynamic analysis we identify two requirements for catalysts in terms of OH* and O* binding energies that would lead the product selectivity towards H2O2 evolution including, ∆ ≳ 3.5 , and ∆ ≲ 2.4 . Figure (1c) shows the free energy diagram for WO3 as an example of a material that satisfies these criteria.

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Figure 1. Free energy diagram for water oxidation reaction on (a) TiO2(110), (b) IrO2(110), and (c) WO3(100) surfaces plotted at zero potential. The OH*, O* and OOH* binding energies of TiO2(110) and IrO2(110) are adapted from Ref. [33]. The relevant binding energies for WO3(100) are adapted from Ref. [16]. The two criteria developed above are included in the two-dimensional selectivity diagram shown in Figure 2. The selectivity diagram shows that there is a large region in (∆ , ∆ ) space where O2 evolution (highlighted in blue) is expected to be dominating (the overpotential may be high, but that is another matter), another window associated with weak O adsorption energies (highlighted in green) that has a driving force towards H2O2 evolution, and region for very weakly interacting catalysts where OH radical formation (highlighted in red) is expected to be dominant. The map also shows why it is hard to find catalysts for selective H2O2 synthesis. The O and OH adsorption energies tend to be correlated as indicated by the dashed scaling line, and only crosses a corner of the region of H2O2 selectivity.

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Figure 2. Phase diagram in terms of the binding energies of O* vs. OH*. The black solid line displays the scaling line between O* and OH* on different oxides adapted from Refs. [16,19,33,34]. The corresponding adsorption energies for IrO2, RhO2 and PtO2 are adapted form Ref. [33]. For MnO2 the energies are from Ref. [44] with selected Hubbard U value of 2 as recommended by Wang et al. 45 BiVO4 adsorption energies are calculated in this work. Blue, green and red highlighted colors indicate regions in which O2, H2O2 or OH radical are expected to be the major product, respectively in terms of purely thermodynamic constraints. A number of different reported oxides have been included in Figure 2. Most of the data are adapted from previous reports.16,19,33,34,44 In addition we calculated the BiVO4, in order to be able to make comparisons with all current the experimental reports. Details of calculations are in the supplementary information. The selectivity diagram rationalizes the experimental observations for OH radical, H2O2 and O2 evolution. It shows that WO3, BiVO4, MnO2, and SnO2, which are known to produce H2O2, are well into the green region, whereas TiO2, which is known to form both OH radical and H2O2, is close to the boarder of weak OH adsorption, red region. Lastly, catalysts such as IrO2, RhO2 and PtO2, which are known to evolve O2 as the major product are thermodynamically located in the selectivity region for O2 evolution. Given the ability to systematize known catalyst trends, we suggest that the selectivity diagram can also be used as a tool to identify leads for new H2O2 production catalysts. Promising leads need to be further investigated to make sure that there are not kinetic barriers preventing their use. We also note that we have not addressed the question of production rates. In summary we developed a cohesive understanding of water oxidation incorporating the formation of hydroxyl radicals, a one-electron water oxidation reaction. This analysis provides a solid theoretical background for rationalizing product selectivity in a number of different experimental reports.

Supporting Information Available: Computational details, description of the BiVO4 structure and adsorption energies, estimated free energy of OH radical are included.

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Acknowledgement We gratefully acknowledge support from the U.S. Department of Energy, Office of Sciences, Office of Basic Energy Sciences, to the SUNCAT Center for Interface Science and Catalysis. S.S acknowledges support from the Global Climate Energy Project (GCEP) at Stanford University (Fund No.52454). Part of the calculations was financially supported by Henan University of Science and Technology (No. 2013ZCX018) and National Natural Science Foundation of China (Nos. U1404212 and 11404098).

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