Combining Experiment and Theory to Unravel the Mechanism of Two

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Combining Experiment and Theory to Unravel the Mechanism of TwoElectron Oxygen Reduction at a Selective and Active Co-Catalyst Vanessa Jane Bukas, Hyo Won Kim, Robert Sengpiel, Kristian B. Knudsen, Johannes Voss, Bryan D. McCloskey, and Alan C. Luntz ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.8b02813 • Publication Date (Web): 09 Nov 2018 Downloaded from http://pubs.acs.org on November 9, 2018

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Combining Experiment and Theory to Unravel the Mechanism of Two-Electron Oxygen Reduction at a Selective and Active CoCatalyst Vanessa J. Bukas†‡, Hyo Won Kim§, Robert Sengpiel§#, Kristian Knudsen§, Johannes Voss†‡, Bryan D. McCloskey§¶* and Alan C. Luntz†‡* †SUNCAT

Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, Stanford, California 94305-5025, United States ‡SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, 2575 Sand Hill Road, Menlo Park, California 94025, United States §Department of Chemical and Biomolecular Engineering, University of California, Berkeley, California 94720, United States ¶Energy Storage and Distributed Resources Division, Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States ABSTRACT: We present a combination of comprehensive experimental and theoretical evidence to unravel the mechanism of two-electron oxygen reduction reaction (ORR) on a catalyst composed of mildly reduced graphene oxide supported on P50 carbon paper (mrGO/P50). This catalyst is unique in that it shows > 99% selectivity towards H2O2, the highest mass activity to date, and essentially zero overpotential in base. Furthermore, the mrGO catalytically active site is unambiguously identified and presents a unique opportunity to investigate mechanisms of carbon-based catalysis in atomistic detail. A wide range of experiments at varying pH is reported; ORR onset potential, Tafel slopes, H/D kinetic isotope effects and O2 reaction order. With DFT reaction energies and known thermodynamic parameters, we calculate the potential and pH dependent free energies of all possible intermediates in this ORR and propose simple kinetic models that give semi-quantitative agreement with all experiments. Our results show that mrGO is semiconducting and cannot support the conventional mechanism of coherently coupled proton-electron transfers. The conducting P50 provides electrons for initiating the ORR via outer-sphere electron transfer to O2(aq), while the semiconducting mrGO provides the active catalytic sites for adsorption of O2-(aq) or HO2(aq), depending upon electrolyte pH. Due to this unique synergistic effect, we describe the mrGO/P50 as a co-catalyst. This concept implies departure from the traditional picture of predicting catalytic activity trends based on a single descriptor, and the co-catalyst design strategy may generally enable other semiconductors to function as electrocatalysts as well. Oxygen reduction reaction, mechanism, graphene, hydrogen peroxide, superoxide, pH, kinetic isotope effect, reaction order.

Oxygen reduction reaction, mechanism, graphene, hydrogen peroxide, superoxide, pH, kinetic isotope effect, reaction order

INTRODUCTION The oxygen reduction reaction (ORR) is one of the technologically most important electrocatalytic reactions. Aqueous ORR occurs through both two-electron (2e-) and four-electron (4e-) reductions. In acidic environments these are given as

(1) O2 + 4(H+ + e-)  2H2O U0 = 1.23 VRHE (2) O2 + 2(H+ + e-)  H2O2 U0 = 0.70 VRHE where U0 is the standard equilibrium potential for the reactions1 and VRHE stands for the potential relative to the reversible hydrogen electrode (RHE). VRHE is related to that of the standard hydrogen electrode (SHE) via VRHE = VSHE + 0.0592 pH so that the thermodynamics of the reaction is independent of pH on the VRHE scale. In strong basic solutions,

H2O becomes the proton donor and at pH > 11.7 these equations become1

(3) O2 + 2 H2O + 4e-  4OH- U0 = 1.23 VRHE (4) O2 + H2O + 2e-  HO2- + OH- U0 = 0.76 VRHE The 4e- ORR is the reaction that generally limits fuel cell efficiency while the 2e- ORR provides an electrochemical path for production of hydrogen peroxide, an important industrial chemical. Since an energy intensive centralized anthraquinone oxidation process now produces H2O2, there has been much recent interest in finding a simple decentralized electrochemical synthesis of H2O2 that could allow small-scale local production for applications such as cleaning of local water supplies.2, 3

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As the U0 for equations (1-4) indicate, the 4e- ORR is always more exoergic than the 2e- ORR. Therefore, thermodynamics always favors the 4e- ORR and a selective catalyst is required to form H2O2 electrochemically. Several metallic and alloy catalysts are partially selective for H2O2 formation in basic conditions, e. g. Au(111)4, Ag(111)5 and PtHg6. Various forms of carbon (glassy carbon7, graphite8, meso/micro porous carbons9) are also partially selective for H2O2 formation, albeit with modest activities. Doped graphenes are generally more active catalysts and give dominantly either 2e- or 4e- ORR, depending on synthesis conditions and/or dopant combinations.10, 11 At present, there is considerable debate as to the identity of the catalytically active sites for these doped graphenes.12 Kim et al.13 have recently shown that a catalyst of mildly reduced graphene oxide supported on P50 Avcarb carbon paper (mrGO/P50) showed selectivity of ≥ 99% to H2O2, the highest mass activity of any H2O2 catalysts previously reported, and essentially zero overpotential in strong basic solution (pH = 13). Extensive spectroscopic evidence identified the catalytically active sites in mrGO as sp2 carbons adjacent to epoxy groups.13 Because the catalytically active site for this H2O2 catalyst is so well established experimentally, this presents a unique opportunity to investigate the mechanism of this 2e- ORR in detail and to understand the origin of its high catalytic activity. ORR electrocatalytic activity on metals and alloys has traditionally been described in terms of the thermodynamics of all of the surface adsorbed intermediates, with the potential dependence of the free energies of the reduction steps given by the computational hydrogen electrode H+ + e- = ½ H2.14 This defines a limiting overpotential for the ORR and plots of this limiting overpotential versus a descriptor such as OH* (where * refers to surface adsorbed species) binding energy have been very successful in describing and predicting ORR activity volcanoes for metals and alloys, both for 2e- and 4e- ORR.15 Related analyses on perovskites using the d-electron occupancy of the transition metal as a descriptor also predict activity volcanoes.16 These descriptions implicitly assume that the mechanism for ORR is based on a series of coherently coupled proton-electron transfers (CPETs). The thermodynamics does not, however, describe the selectivity of ORR since the weak binding legs of the 2e- and 4e- volcanoes are identical and it is only in this weak binding region that 2eORR is observed.17 This same CPET mechanism has also recently been proposed to describe both 4e- ORR and 2 e- ORR on carbonbased catalysts.18, 19, 9 On the other hand, earlier electrochemical studies of the 2e- ORR on carbon7, 8 proposed mechanisms initiated by pure electron transfers, i. e. O2(aq) +e-  [O2-]*, similar to the decoupled proton-electron transfers that typically occur in molecular electrocatalysis.20, 21 Thus, there is considerable disagreement as to the exact mechanism responsible even for the simplest 2e- ORR on these promising non-metal catalysts. However, such understanding is essential for going beyond the catalysis of precious metals/alloys and for guiding new future strategies in electrocatalyst design. This paper describes both detailed experiments and first principles theory to unravel the mechanism behind the active and selective mrGO/P50 2eORR catalyst. We report on a wide range of experiments at varying pH that find semi-quantitative agreement with simple

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kinetic models proposed from theory. A detailed understanding of the mechanism thus emerges and reveals that the high activity of mrGO/P50 is based on its unique behavior as a co-catalyst that solely supports decoupled proton-electron transfers.

EXPERIMENTAL & METHODS Electrochemical measurements

COMPUTATIONAL

The electrochemical techniques have been described in detail elsewhere so that only novel aspects will be emphasized here.13 Electrochemical measurements were performed using a Biologic SP-300 potentiostat and a custom-built modified hermetically sealed electrochemical H-cell. The voltages of both the counter electrode (Pt wire) and working electrode (WE, mrGO coated P50 Vulcan C, 1.0 cm2 exposed to electrolyte) relative to a reference electrode were recorded simultaneously. A Hg/HgO reference electrode immersed in 1M NaOH was utilized for basic solutions (pH = 9, 13) and a Hg/Hg2SO4 reference electrode immersed in concentrated K2SO4 for acidic solutions (pH = 1, 5). All potentials are reported against the reversible hydrogen electrode (RHE) or standard hydrogen electrode (SHE) and are corrected for ohmic losses. All electrochemical measurements were performed saturated with pure O2, typically at ~1.07 bar unless otherwise noted. Linear sweep voltammetry (LSV) was performed by a cathodic sweep of the working electrode potential from open circuit potential (OCV) at a scan rate of 2.0 mV/s. Simultaneous measurements of the O2 pressure change in a calibrated head-space volume during the LSV gave measurements of e-/O2, the moles of electrons consumed relative to the moles of O2 consumed. A detailed discussion of this setup and its advantages is found in previous work.13 Chronopotentiometry (CP) following an ORR current step from open circuit was conducted to measure the Tafel slopes. pH = 1 and pH = 13 electrolytes used 0.1 M of H2SO4 and KOH solutions, respectively. The pH = 5 electrolyte was created by adding enough 0.1 M H2SO4 solution to 0.1 M K2SO4 solution to reach pH = 5, and pH = 9 electrolyte was also prepared by adding enough 0.1 M of KOH solution to 0.1M K2SO4 solution to reach pH = 9. Kinetic isotope effect experiments were performed in deuterated aqueous solutions (D2O). pD = 1 and pD = 13 electrolytes used 0.1M of D2SO4 and KOD solutions in D2O, respectively. To minimize mass transport limitations during LSV, vigorous stirring of the electrolyte was utilized. During the CP measurements (30 sec), there was no active stirring. However, between CP measurements (2 min), stirring was active. For all data presented (Figs. 1-4), the same electrode was used to collect each profile in that figure, such that the geometric surface area in each was consistent (~1 cm2).

Computational methods We model the mrGO catalyst and corresponding ORR thermodynamics using density functional theory (DFT) of a periodic supercell within the generalized gradient approximation (GGA). DFT calculations rely on the BEEFvdW exchange-correlation (xc) functional22 and the Quantum ESPRESSO plane-wave code23 with ultra-soft pseudopotentials24 and an energy cut-off of 450 eV. The mrGO Brillouin zone is sampled at the -point and periodic images are separated by 12 Å of vacuum along the surface normal. In order to evaluate charged systems such as those


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discussed later, we use non-periodic DFT calculations within the all-electron basis set of the FHI-aims code.25 These are based on an H-terminated, cluster representation of the mrGO structure that was rigorously tested to ensure that it reproduces the electronic structure of the periodic model (see Fig. S11 in the Supplementary Information (SI)). These calculations utilize the PBE xc functional26 in combination with the Tkatchenko-Scheffler van der Waals correction scheme.27 We calculate free energies for all species adsorbed on mrGO in the absence of the aqueous environment (G) by correcting the DFT electronic energy differences (EDFT) by zero-point-energy (ZPE) and entropy (TS) as described by Nørskov et al.14 and further detailed in the SI. The formation free energies of OH*, H*, OOH* adsorbates are calculated as a function of VSHE and pH by using the computational hydrogen electrode.14 Free energy differences for the full ORR reactions (1)-(4) are taken from experiment and all states are referenced against H2(g) and H2O(l) to avoid an explicit DFT calculation for O2(g). This methodology allows for conveniently evaluating the thermodynamics and forms also the basis for our Pourbaix analysis, as explained in detail in the SI. Finally, we estimate the stabilization of OH*, H*, OOH* adsorbates on mrGO due to H-bonding from the local aqueous environment. In a first step, the solvent H2O structure is efficiently sampled using molecular dynamics (MD) simulations at the level of classical interatomic potentials.28, 29 As described in detail in the SI, the solvent is thermally equilibrated at 300 K using a Nose-Hoover thermostat within the LAMMPS simulation package.30 In a second step, we evaluate the 300 K ensemble via single-point DFT calculations on two-hundred uncorrelated structures sampled from the classical MD. The mean of the resulting Gaussian distributions defines the average thermodynamic energy. We perform this procedure for both the adsorbate-covered and clean mrGO surfaces and form the corresponding energy difference to estimate the (average) adsorption energy in solution. Comparing this value to the equivalent vacuum result yields a solvation correction for each OH*, H*, OOH* adsorbate type that is used to correct the vacuum free energies of adsorption.

RESULTS AND DISCUSSION

of the carbon D’ band in in-situ Raman spectroelectrochemistry during ORR at pH = 13 inferred that the catalytically active sites were sp2 hybridized carbons on the basal planes of mrGO.13 Figure S4 shows that similar D’ broadening occurs at pH = 1 as well and suggests that the same sp2 basal plane carbon sites are the catalytically active ones at both pH = 1 and pH = 13. Table 1. Ratio of e- to O2 consumption measured by linear sweep voltammetry at different pH. The ideal ratio for forming H2O2 is 2.00. Details of the experiments are provided in Fig. S2 of the SI. pH

1

5

9

13

e-/O2

1.96

2.01

1.98

1.99

Figure 1 shows the pH dependence of the ORR LSV on mrGO/P50, both as a function of VRHE (a) and VSHE (b). It is clear from this figure that the onset of the ORR is independent of VSHE and shows a shift of ~ 59 mV/pH unit on the VRHE scale. Since the VRHE scale defines the thermodynamics of the ORR, there is no apparent overpotential at pH = 13. However, there is a large pH dependent overpotential as the electrolyte becomes more acidic. This also suggests that the potentiallimiting step in the ORR is not a CPET, as this would produce an onset that was constant on the VRHE scale, especially at low pH where protons are certainly the dominant reducing species. This conclusion is strengthened by measuring the H/D isotope effect for ORR at pH = 1 and pH = 13 as shown in Fig. 2 and S8. The lack of any isotope effect in the onset potential for ORR again suggests that proton transfers from either H3O+ or H2O are not involved in the potential-limiting step. If a H+ were involved, its replacement with D+ would effectively change the free energy difference (G) for this step and appear as a shift in the ORR onset potential (G=eU), as explained in detail around Fig. S8 in the SI. We do note, however, that although there is no measurable isotope effect for pH = 1, there is a small inverse isotope effect in the ORR current at pH = 13 as given in Fig. 2 and S8.

In this section, we first describe experimental results for ORR on mrGO/P50 that span a range of pH and that were specifically designed to probe the catalytic mechanism(s). Some qualitative mechanistic conclusions are obvious directly from the experiments, and hence are discussed alongside the latter. We then develop a first-principles theoretical model of the ORR that gives good semi-quantitative agreement with all experiments. Throughout the discussion of the mechanism, we distinguish between potential-limiting steps that define the onset of ORR and rate-limiting steps that define the current dependence with potential during ORR.

Experimental results It was previously shown that mrGO supported on P50 carbon paper is ≥ 99% selective at pH = 13 for 2e- ORR to produce H2O2.13 The extreme H2O2 selectivity of this catalyst is independent of pH as shown in Table 1 and Fig. S2 of the SI. Figure S5 demonstrates that mrGO/P50 is much more active as an ORR catalyst than P50 alone at both pH = 1 and pH = 13. However, the onset potential for ORR is nominally the same for both P50 and mrGO/P50 at both pH. Broadening


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Because both the cathode double layer capacitance Cdl and the Faradaic current depend on U through the ORR regime, we use chronopotentiometry following current steps of varying magnitudes from open circuit to separate the two dependencies in order to obtain the best Tafel slopes. In the chronopotentiometry, charging Cdl causes a transient potential response, while the Faradaic reaction represents the steady state U at the given current i as shown in Fig. S3. The Faradaic Tafel behavior at pH = 1 and 13 from the chronopotentiometry are given in Fig. 3, with a Tafel slope of b = 59 mV/decade at pH = 13 and b = 179 mV/decade at pH = 1. This implies that the rate-limiting step, in contrast to the potential-limiting step, is different at the two pH and the reason for this will be discussed below when the theoretical model is presented.

Figure 1. ORR cathodic linear sweep voltammetry (current i in mA vs applied potential U in volts) at the pH indicated in the figure to form H2O2, where the applied potential is given on the scale of a) the reversible hydrogen electrode (RHE), and b) the standard hydrogen electrode (SHE).

Figure 3. Tafel plots (log of the ORR current i in Amps vs potential U in volts relative to the standard hydrogen electrode) for both pH = 1 and pH = 13 as obtained from chronopotentiometry.

We obtain the order of the reaction in O2 by measuring the ORR LSV as a function of O2(g) pressure (𝑃𝑂2) at both pH=1 and pH=13 as shown in Fig. S6. The reaction order is defined as 𝜌 ≡

[

∂log 𝑖

] , i.e. reflecting the change in overall current i

∂log 𝑃𝑂2

𝜂

with 𝑃𝑂2 at constant overpotential conditions in the Tafel regime. The results are shown as a function of 𝜂 in Fig. 4a). Both pH=1 and pH=13 show 𝜌 < 1, but with a qualitatively different dependence on 𝜂. Unfortunately, the experimental Ueq for pH = 1 deviated considerably from the expected Nernstian shift with 𝑃𝑂2 (likely due to build-up of H2O2 during sequential experiments) and this made the absolute values of 𝜌 (𝜂) at pH = 1 less accurate than for pH = 13.

Figure 2. ORR cathodic linear sweep voltammetry (current i in mA vs applied potential U in volts on the reversible hydrogen electrode scale) to form H2O2 (solid lines) and D2O2 (dashed lines) at pH = 1 and 13.

Although the ORR onset Ueq occurs at U ~ 0 VSHE at all pH, there is a difference with pH in how the ORR current depends on U or overpotential 𝜂 = |𝑈 ― 𝑈𝑒𝑞|. At low 𝜂, before transport significantly limits the charge transfer kinetics, i. e. in the linear Tafel regime, pH = 1 and 5 appear to have higher Tafel slopes than pH = 9 and 13. At higher ORR currents, the intermediate pH = 5 and 9 have a more complicated Tafel dependence showing evidence of either more extensive transport limitations or a transition between rate-limiting steps.


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Figure 4. ORR reaction order 𝜌 of O2 gas as a function of overpotential 𝜂 in volts for both pH =1 and pH =13. a) from experiment as discussed in the text and b) from simple kinetic models based on the theoretical mechanisms.

In the following, we discuss pH dependent mechanisms that give semi-quantitative agreement with these experiments. We first develop a DFT-based computational model of mrGO and its equilibrium composition in aqueous solution. Based on theoretical free energies of the intermediates in the ORR at the surface and in solution, we propose mechanisms for ORR at the different pH. We then use qualitative kinetics based on the derived free energy diagrams to discuss the experimental results.

DFT model of the mrGO active catalyst Dispersed graphene oxide (GO) typically consists of micron-sized flakes that are dominated by hydroxyl and epoxy groups on the basal planes, but with a concentration of various functional groups (OH, C=O, COOH, etc) on the edges of the planes.31 The reduction of graphene oxide (rGO) has been extensively studied spectroscopically and theoretically by many authors.32 In the previous study of mrGO13, a combination of infrared spectroscopy (IR), near edge x-ray absorption spectroscopy (NEXAFS), X-ray Photoelectron Spectroscopy (XPS) and Raman in-situ spectroelectrochemistry indicate that thermal reduction in H2O at 100°C forms a mildly reduced graphene oxide mrGO, with C/O ratio of ~ 2.85. Here the basal plane is dominated by loss of some of the OH and epoxy groups, with aggregation of the remaining groups. Mildly reduced GO thus consists of nmscale patches of graphene interspersed with similar sized patches of epoxy-rich GO.33, 34 With further reduction, some of the epoxy and OH groups remaining are reduced resulting in a shrinking of the GO islands while the graphene patches grow and weakly conducting links are established between them.35 This allows a reasonable characterization of the rGO in terms of the fraction of sp2 carbons (or equivalently the C/O ratio). Given this background, we discuss a simple theoretical model of the mrGO given in Fig. 5. It corresponds to a C/O = 2.67, close to that measured for mrGO. This model is equivalent to one of those proposed by Lundie et al. who studied the band gaps and optical properties of rGO in terms of the C/O ratio.36

represent C atoms and red spheres represent the epoxy O atoms. The breaking of -bonds gives rise to two unpaired electrons that are distributed at the boundaries of the two regions. The yellow spheres represent the DFT derived spin charge density of the unpaired electrons.

Depending on the concentration of functionalizing epoxies and OH groups, the rGO electronic structure varies between that of pure graphene (C-sp2 = 1) and pure GO (C-sp2 = 0). With decreasing C/O ratios, graphene’s semi-metallic character is gradually lifted as sp2-hybridized carbons are converted to sp3 carbons and this gives rise to a significant band gap as shown by the calculated density of states (DOS) in Fig. S13. As a result, the electrical conductivity is also a strong function of the C-sp2 fraction as outlined in Fig. S9.35 The mrGO catalyst has a C-sp2 fraction of ~ 0.35 (C/O ~2.85) so is predicted to have a very low inherent electrical conductivity of < 10-10 S/cm (see Fig. S9). Our model of mrGO only includes epoxy regions on the basal plane. However, this simplification maintains the basic feature of the electronic structure based on C-sp2 regions interspersed with C-sp3 regions. Calculation of the density of states (DOS) for this mrGO model structure is shown in Fig. 6(a), consistent with earlier calculations by Lundie et al.36 The DOS for this model shows a number of localized C defect states that emerge within the semiconducting energy gap. These are associated with the breaking of -bonds within the sp2-hybridized graphene domains forming local unsaturated ‘dangling’ -bonds as indicated in Fig. 5. These form half-filled states that appear around the Fermi level in the DOS of Fig. 6(a). Spin-polarized GGA calculations (see Fig. S14 in the SI) confirm this interpretation of the dangling -bonds. Integrating the occupied pz DOS in these dangling bonds reveals a total of two unpaired electrons for this model of mrGO. These are not localized over single surface sites but are generally distributed, with varying weights, as given by the spin density among different C-sp2 sites (size of yellow spheres in Fig. 5). The latter are predominantly found at the boundary of the graphene-GO regions as illustrated by the spin charge density in Fig. 5.

Figure 5. Model of the mildly reduced graphene oxide (mrGO) catalyst with C/O ratio of ~ 2.67 showing nm-sized patches of graphene interspersed with nm-sized patches of GO. Gray spheres


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Fig. 6(c). Hybrid-DFT calculations confirm this picture and show even an increased band gap as compared to the GGA band gap (see Fig. S16 in the SI). Clearly, the number of dangling bonds correlates with both the size and shape of the embedded graphene patches.36 We find that quenching of all dangling bonds by OH- or H+ adsorption as defined by the Pourbaix diagrams is true for less symmetric mrGO models or those with different sized graphene patches as well. It is this that justifies using a simple single model to define the mechanism on the complicated mrGO catalyst structure.

eV-1

Figure 6. Projected density of states (pDOS) in vs energy relative to the Fermi energy of carbon atoms in (a) the uncovered mrGO model structure, and after adsorption of (b) one and (c) two OH- species due to alkaline conditions. Colored lines show the angular-resolved px, py, pz components as denoted in the figure’s legend. The pDOS were obtained with a 0.1 eV Gaussian smearing.

These dominant dangling -bonds sites are naturally most active towards adsorption and therefore likely covered by OHor H+ species in alkaline or acidic solutions. Figure 6(b) shows that adsorption modifies the -bonding network and quenches the existing dangling bonds. Since adsorption of 2 OH- (Fig. 6(c)) quenches all dangling bonds, further adsorption is less likely and the mrGO becomes fully semiconducting. This is consistent with the prediction of two unpaired electrons in mrGO from integrating the DOS occupation at the clean surface. Adsorbing H+ species in acid leads to similar conclusions (see Fig. S15 in the SI). Whether OH- or H+ adsorption occurs, however, depends upon whether the thermodynamics of adsorption is downhill relative to OH-/H+ stabilities in solution, i e. on the surface Pourbaix diagrams. This depends both on vacuum calculated DFT adsorption energies and the stabilization of these adsorbates by the surrounding H2O solvent. The latter is included after efficient ensemble sampling of the H2O structure as described briefly in the Methods section and further elaborated in the SI. We estimate the following solvent stabilization energies for each adsorbate type: -0.4 eV for OH*, -0.2 eV for H*, and -0.4 eV for OOH*. Fig. 7 shows the surface Pourbaix diagram for sequential adsorption of a first, second, and third OH- species in the graphene regions of mrGO at the dangling bond sites (marked in Fig. S17 of the SI). Figure 7 suggests that only adsorption of the first two OHspecies is spontaneous (GOH* 0.2 V) that are incompatible with η < 10mV in base (see Fig. S24). If the surface were not fully OH/H-saturated, the thermodynamics of CPET forms too strong-binding in OOH*, with η > 0.3 V which again is inconsistent with experiments (see Fig. S24). We thus exclude CPET and investigate a decoupled proton-electron transfer potential limiting step as indicated by the experimental results.

Thermodynamic analysis of 2e- ORR on mrGO The standard aqueous electrochemical reduction of O2 is

(5)

where

𝑂2(𝑔) + 𝑒 ― = 𝑂2(𝑎𝑞) + 𝑒 ― ⇌ 𝑂2― (𝑎𝑞) ∆𝐺5 = 𝑒 (𝑈 ― 𝑈𝑒𝑞 5 )

(

𝑃𝑂 2

0 𝑈𝑒𝑞 5 = 𝑈5 ―0.0592 𝑉 ∙ log [𝑂 ― (𝑎𝑞)] 2

)

is

the

equilibrium potential that is related to the standard-state


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O2(aq) reduction potential 𝑈05 = ― 0.33 𝑉SHE1 via the Nernst equation. 𝑃𝑂2is the O2 gas pressure in bar and [𝑂2― (𝑎𝑞)] is the molar concentration of the superoxide. The left equality in reaction (5) simply reflects that O2(aq) is in equilibrium with O2 gas. When considering a likely experimental O2-(aq) concentration of ~ 10-6 M near the onset of the ORR, 𝑈eq 5 ~ 0 VSHE for this reaction, as shown in the H2O/H2O2 Pourbaix diagram in Fig. S1. Since VSHE ~ 0 is the experimental onset of ORR at all pH (Fig. 1), we contend that an outer-sphere reduction of O2(aq) to O2-(aq) is the potential-limiting step of ORR on mrGO/P50 at all pH. Since mrGO is a semiconductor with no electrons available at EF, we argue that reaction (5) is initiated by electron transfer from the conducting P50 carbon support. However, without the mrGO catalyst, the 2e- ORR activity is low (but with similar onset potentials) as shown in Fig. S5. Therefore, we refer to this throughout the paper as a mrGO/P50 co-catalyst. Once formed, O2-(aq) behaves differently in acidic and alkaline media as defined by the pKa of HO2(aq). The thermochemistry of the O2-/HO2 couple is given by:

(6)

𝑂2― (𝑎𝑞) + 𝐻 + ⇌𝐻𝑂2(𝑎𝑞) ∆𝐺6 = ∆𝐺06 +0.0592 eV ∙ pH

The O2-(aq) at pH > 4.8 or HO2(aq) at pH < 4.8 will adsorb on the OH/H covered mrGO surface if ΔG < 0 for the adsorption reactions. We use DFT as described in the Methods section to calculate the ΔG for each of these processes. The ΔG for the adsorption reaction in acid is

𝐻𝑂2(𝑎𝑞) + ∗ →𝑂𝑂𝐻 ∗ ∆𝐺7 = ―0.22 eV

where ∆𝐺7 is obtained by combining the free energy of HO2(aq) as determined above (Eqs. 5, 6) with that of a surface-bound OOH* at the most stable adsorption site of the H-covered mrGO surface. The free energy of OOH* is calculated from the thermodynamics using the H2O reference for O2 (see details in the SI) and including also the -0.4 eV solvation correction for OOH*. The ΔG for the adsorption reaction in base is

(8)

(9)

𝑂2― (𝑎𝑞) + ∗ →[𝑂2― ] ∗ ∆𝐺8  ―0.18 eV

where ∆𝐺8 is obtained by non-periodic DFT calculations on the OH-covered mrGO cluster model discussed in the Methods section. We describe O2- solvation by explicitly modeling the ion’s first solvation shell of four H2O molecules38 and this accounts for the majority of the total solvation energy.39 The large uncertainty in ΔG8 arises principally from the uncertainty in whether the adsorption of [O2(H2O)4]-1 increases the water density locally at the surface. While this is possible, we believe the increase in water density should be minimal and the dashed line for O2-* represents the case with no increase in water density. The grey shaded area below this line in Fig. 8(b) represents the range of values possible including an

[𝑂2― ] ∗ + 𝐻2𝑂 ⟶𝑂𝑂𝐻 ∗ + 𝑂𝐻 ― (𝑎𝑞) ∆𝐺9 = ∆𝐺09 +59 meV ∙ pH

where ∆𝐺09 = +0.28 eV is obtained by combining the free energy of [O2-]* as determined above (Eqs. 5, 8) with that of a surface-bound OOH* that is evaluated from thermodynamics using the H2O reference for O2 (see discussion in SI). Once OOH* exists, it can undergo a proton-electron transfer to form H2O2 (pH < 11.7) or an electron transfer to form HO2(pH > 11.7) as determined by the pKa of H2O2 and given in eqs. (10) and (11).

(10)

where ∆𝐺06 = ―0.28 eV is the standard-state reaction free energy based on the experimental pKa = 4.8 of HO2(aq).37 This relates the free energy of HO2(aq) to O2-(aq), with the latter in equilibrium with O2 gas at 𝑈𝑒𝑞 5 .

(7)

increase in water density at the surface and is discussed in detail in the SI. We note that the most stable sites for both the OOH* and [ 𝑂2― ] ∗ adsorption is a C-sp2 site next to an epoxy. This is exactly the catalytically active site proposed earlier by Kim et al. in their spectroscopic studies of ORR on mrGO.13 Having formed an [O2-]* species in base, this can protonate via

𝑂𝑂𝐻 ∗ + 𝐻 + + 𝑒 ― →𝐻2𝑂2(𝑎𝑞) + ∗ ∆𝐺10 = ∆𝐺010 +59 meV ∙ pH + 𝑒𝑈

where ∆𝐺010 = ―1.22 eV is obtained by combining the free energies of OOH* (eq. (13) in the SI) and H2O2(aq) (from eq. (1)), with a common H2O reference. Note that we do not assume that this step is a CPET as will be discussed in more detail later.

(11)

𝑂𝑂𝐻 ∗ + 𝑒 ― →𝐻𝑂2― (𝑎𝑞) + ∗ ∆𝐺11 = ∆𝐺011 + 𝑒𝑈

where ∆𝐺011 = ―0.47 eV is obtained by combining again the free energies of OOH* (eq. (13) in the SI) and H2O2(aq) (from eq. (1)), along with the experimental pKa=11.7 of H2O2(aq).

Mechanism and kinetics of 2e- ORR Given this thermodynamic information on reaction intermediates, we propose the following mechanisms for ORR on mrGO/P50 to H2O2(aq) or HO2-(aq). As discussed previously, the first step in the ORR is the outer-sphere reduction of O2(aq) to O2-(aq) initiated by the P50 rather than the mrGO. Since the mrGO is supported on P50, this solutionmediated electron transfer may be initiated by a bare region of P50 that is next to the mrGO, followed by diffusion of the superoxide to the active catalyst. Alternatively, it could also occur via tunneling from the P50 through the semiconducting mrGO at points of contact. While both methods of initiating the outer-sphere electron transfer will lead to the same kinetics, we believe the former is more likely. This is because i) the P50 has much more surface area than the mrGO and because ii) tunneling induced ORR/OER through a semiconductor is limited to lengths of < 4 nm, while the mrGO stack of a few layers may be thicker.40 Therefore we distinguish between intermediate species that are formed in the aqueous electrolyte adjacent to the P50 support, [𝑂2― (𝑎𝑞)]𝑃50 in base or [𝐻𝑂2(𝑎𝑞)]𝑃50 in acid, from those which diffuse adjacent to the mrGO active site, [𝑂2― (𝑎𝑞)]𝑚𝑟𝐺𝑂 or [𝐻𝑂2(𝑎𝑞)]𝑚𝑟𝐺𝑂. We treat these diffusions as simple rate


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𝑘13

expressions that should be a good approximation at steady state.41 Therefore, to discuss the kinetics we add the following diffusion reactions:

(13)

𝑘12

(12)

[𝑂2― (𝑎𝑞)]𝑃50

𝑘 ―12

𝑘5

𝑂2 + ∗ +2(𝐻 pH=1

𝑘13 𝑘 ―13

+𝑒 ) ―

𝑘 ―5

[𝑂2― (𝑎𝑞)]𝑃50 +

[𝐻𝑂2(𝑎𝑞)]𝑚𝑟𝐺𝑂 + ∗ + (𝐻 + + 𝑒 ― )

𝑂2 + ∗ + (𝐻2𝑂 + 2𝑒 pH=13

𝑘12

𝑘 ―12

[𝑂2― (𝑎𝑞)]𝑚𝑟𝐺𝑂 + 𝑘9

[𝐻𝑂2(𝑎𝑞)]𝑃50

𝑘 ―13

[𝐻𝑂2(𝑎𝑞)]𝑚𝑟𝐺𝑂

Assuming k6 >> k12, the mechanism at pH = 1 is summarized in Scheme 1.

[𝑂2― (𝑎𝑞)]𝑚𝑟𝐺𝑂

+

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―

𝑘5

)

𝑘 ―5

𝑘7

𝑂𝑂𝐻 ∗ + (𝐻 + + 𝑒 ― )

[𝑂2― (𝑎𝑞)]𝑃50 +

∗ + (𝐻2𝑂 + 𝑒 ― )

𝑂𝑂𝐻 ∗ + 𝑂𝐻 ― + 𝑒 ―

∗ + (2𝐻 + + 𝑒 ― )

𝑘11

𝑘7

𝑘10

𝐻2𝑂2(𝑎𝑞) + ∗

∗ + (𝐻2𝑂 + 𝑒 ― ) ∗

[𝑂2― ] + (𝐻2𝑂 + 𝑒 ― )

𝐻𝑂2― (𝑎𝑞) +∗ + 𝑂𝐻 ―

Scheme 1. 2e- ORR mechanism on the mrGO/P50 catalyst at pH=1 (top) and pH=13 (bottom). Note that, even at U ~ 0 VSHE, OOH* has states at the Fermi level available to accept electron transfer since its adsorption reintroduces a half-filled state at the Fermi energy (see Fig. S25). However, the last step described by k10 is likely not a CPET because of the slow e- mobility from the conducting P50 support to OOH*. Similarly, at pH = 13 the mechanism is also summarized in Scheme 1. Figure 8 shows the free energy diagram for the above mechanisms at pH=1 (a) and pH=13 (b). All values are referenced to the free energy of O2(g) at one bar pressure and

U = 0 VSHE. Thus, the energy levels are aligned such that they represent the experimental ORR onset at 𝑈𝑒𝑞 5 ~ 0.0 VSHE with an O2-(aq) concentration of ~10-6M so that the O2/O2- states are at equilibrium (G5 = 0). The free energy of the protons is included by appropriate pH-dependent shifts. Note that the energy levels do not distinguish [𝑂2― (𝑎𝑞)]𝑃50 from [𝑂2― (𝑎𝑞)]𝑚𝑟𝐺𝑂 nor [𝐻𝑂2(𝑎𝑞)]𝑃50 from [𝐻𝑂2(𝑎𝑞)]𝑚𝑟𝐺𝑂 since they represent the same aqueous species at only slightly different concentrations.


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Figure 8. Free-energy diagram for the proposed ORR pathway at a) pH=1, and b) pH=13. All energy levels are aligned such that they represent the experimental ORR onset at ~ 0.0 VSHE with an O2-(aq) concentration of ~10-6M and where the O2/ O2- states are at equilibrium. The free energy of the protons is included by the appropriate pH-dependent shifts. The shaded grey region in b) represents the uncertainty in that intermediate state due to the unknown change of water density at the charged surface and the GGA-DFT error in describing the superoxide state.

The simple statement that eq. 5 is the potential-limiting step at all pH explains the dependence of ORR onset with VSHE instead of VRHE and the absence of a H/D isotope effect for the ORR onset potential. This is clearly illustrated in Fig. S26 of the SI which shows the corresponding free-energy diagrams at pH=1 (a) and pH=13 (b) as in Fig. 8, but at the experimental 2e- ORR equilibrium potential of 0.7 VRHE. It also rationalizes why the onset potential is nominally the same for P50 and mrGO/P50. The diabatic barrier schematically indicated in Fig. 8 for this step has been estimated as ~ 0.9 eV by Hartnig et al. using force-field based molecular dynamics.42 Based on the thermodynamic analysis of pH=1 in Fig. 8(a), this is likely the largest barrier encountered along the acidic ORR pathway. This makes the first electrochemical step both potential-limiting and kinetically rate-limiting (cf. also Fig. S27 in the SI). Neglecting any potential dependent surface coverage contributions to the kinetics, the Tafel slope b is given as 𝑏~59 𝑎 mV/decade, with the transfer coefficient 𝑎 = 𝑛𝑝 + 𝑛𝑞𝛽𝑓; where np is the number of electrons transferred prior to the rate-limiting step, nq is the number of electrons transferred in the rate-limiting step and βf is the symmetry factor in the rate-limiting step.43 This predicts b ~ 120 mV/decade for pH = 1, in reasonable agreement with a measured Tafel slope of ~179 mV/decade, and implying a transfer coefficient α = 0.3-0.4. The difference is probably due to large potential dependent surface concentration effects.44 At pH=13, however, Fig. 8(b) suggests that the kinetics of overcoming the chemical barrier of ~0.4 eV for the

protonation of [O2-]* to form OOH* is likely rate-limiting. Assigning this chemical step as kinetically rate-limiting instead of the initial e- transfer step leads to excellent agreement with the measured alkaline Tafel slope of ~59 mV/decade and the equation above.43 Transition between these two limiting mechanisms at intermediate pH is likely a gradual process that depends on many kinetic parameters. This transition and its dependence on current and proton transport is the likely origin of the kinks observed at higher ORR currents at pH =5 and pH = 9 in Fig. 1. This assignment of rate-limiting steps is also consistent with the observed H/D isotope effects on the electrochemical current in Fig. 2 and S8. At pH=1 the absence of any H/D effect (𝑖𝐻/𝑖𝐷 = 1.0) at all U confirms that the acidic ratelimiting step does not involve a proton transfer. At pH=13 we find a small inverse isotope shift that remains constant at 𝑖𝐻/𝑖𝐷 = 0.9, independent of U. Small inverse isotope shifts have previously been observed in OER45 and in a wide range of chemical studies.46 All that is required for an inverse isotopic shift is for the ZPE at the transition state to be greater than that of the reactants for the rate-limiting step. At pH = 13 and pD = 13, our DFT calculations predict an inverse isotope effect 0.2 < 𝑖𝐻/𝑖𝐷 < 1, depending upon the location of the transition state along the reaction path (cf. details regarding this estimate in the SI). Therefore, we believe this inverse isotope effect is in qualitative agreement with the assignment of eq. (9) as ratelimiting in base. We note that if an electrochemical step was


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rate-limiting, there would be slight changes in the Tafel slope so that 𝑖𝐻/𝑖𝐷 would depend on U, in disagreement with Fig. S8.47 Recognizing that different steps are rate-determining at pH=1 and pH=13 also qualitatively rationalizes the different O2(g) reaction orders shown in Fig. 4(a). We model 𝜌(𝜂) by deriving analytic expressions for the overall current with a simple kinetic model based on a single rate-determining step (RDS) with all other preceding kinetic steps in equilibrium. The SI includes detailed derivations for expressing 𝑖 as a function of 𝜂 and 𝑃𝑂2for both the acid and base cases. The modeling results are shown in Fig. 4(b) and yield reasonable agreement with experiment. We believe most of the disagreement for pH = 1 with experiment is related to the aforementioned ambiguities in defining the experimental 𝑈𝑒𝑞 5 and therefore the magnitude of 𝜂. In base, 𝜌(𝜂) follows a more complicated relationship. The abrupt drop of 𝜌 at low 𝜂 is related to the reversibility of the 𝑂2/𝑂2― redox couple and consequently small fraction of [𝑂2― (𝑎𝑞)]𝑃50 that can proceed on to ORR electrochemistry, as discussed in detail in the SI. It is important to note that we have not included in the mechanisms the adsorption of [𝑂2― (𝑎𝑞)]𝑃50 on P50 since it is an electrically conducting support and electron reservoir. The cycle of O2- formation from P50 and its ensuing adsorption on P50 would thus not lead to significant electrochemical current, consistent with Fig. S5. Instead, adsorption of [𝑂2― (𝑎𝑞)]𝑚𝑟𝐺𝑂 on the semiconducting mrGO or irreversible chemical reaction of [𝑂2― (𝑎𝑞)]𝑃50 in solution will produce electrochemical current and ultimately production of H2O2 or HO2The mechanisms discussed above for mrGO/P50 have some similarities to older mechanisms suggested for the 2e- ORR on various forms of much less catalytically active carbon systems (glassy carbon, graphite).48 Yeager, McCreery and others have suggested that the 2e- ORR on carbon is initiated by a pure electron transfer, either in a single step to directly form [O2-]* or in a two-step process after O2 has been adsorbed. It has then been argued that there are at least two types of carbon-bound [O2-]* where one is a quite localized species to ultimately produce electrochemistry. However, a range of different mechanisms has been suggested to ultimately form H2O2 from some localized defect containing [O2-]*.7, 8, 49 An important obstacle in this respect is that for these various carbon surfaces, the identity of the catalytically active site (or sites) is unknown and in fact has been the subject of much discussion.48 In contrast, the catalytically active site for the mrGO/P50 catalyst is uniquely identified here both from experiment and theory. As mentioned in the introduction, thermodynamics always favors the 4e- over the 2e- ORR. The origin of 2e- selectivity in all ORR catalysts must thus reside in the reaction kinetics. Since both paths involve a OOH* intermediate (after the first electron, proton transfers), it has been suggested that the barrier to form H2O along the 4e- pathway (e.g. via OOH* + H+ + e-  O* + H2O) must be larger than that of forming H2O2 (OOH* + H+ + e-  H2O2).15 However, since the calculation of accurate electrochemical barriers is at present quite difficult, this argument has remained largely untested. The origin of the very high 2e- selectivity at the mrGO/P50 catalyst suffers from this same uncertainty, as the formation of H2O from OOH* is 1.3 eV more exothermic than the formation of

Page 10 of 21

H2O2. Therefore, we simply assume that its selectivity is also based on the kinetics of this step, i.e. we contend that it is easier to break the weaker C-O bond rather than the stronger O-O bond. This is generally consistent with the empirical observation that weakly bound OOH* (as at the at the OH/Hsaturated mrGO) tend to selectively follow the 2e- rather than the 4e- ORR.6, 15, 21

CONCLUSIONS In this paper we present a comprehensive combination of experimental and first-principles theoretical evidence to unravel the mechanism behind the 2e- ORR on a catalyst of mildly reduced graphene oxide supported on P50 carbon paper (mrGO/P50). This catalyst is unique in that it shows > 99% selectivity towards H2O2, exceptionally high gravimetric activity and essentially zero overpotential in base. A comparison of the mass activity of the mrGO electrocatalyst at pH = 13 with other H2O2 catalysts in the literature is given in Fig. 4 of Ref. 13. Similar data for further comparison are presented in Ref. 50. However, since the mrGO/P50 acts as a co-catalyst, its mass activity must be somewhat smaller when considering the full electrode mass. While mrGO/P50 is an excellent electrocatalyst in base, it is not a good electrocatalyst in acidic environments as there is a significant thermodynamic overpotential. This arises from the potential-limiting step of outer-sphere electron transfer to O2 that is independent of pH. Furthermore, the catalytically active site is well identified both by experiment and theory. A wide range of experiments at varying pH designed to probe the mechanism are reported; ORR onset potential, Tafel slopes, H/D kinetic isotope effects and O2 reaction order as a function of overpotential. Firstprinciples theory shows that the mrGO is semiconducting, with all dangling bonds quenched by adsorption of OH- or H+ in solution (depending upon pH) so that it cannot support CPET ORR as is typical of metals. Therefore, decoupled electron and proton transfers must drive the ORR. With DFT reaction energies and known thermodynamic parameters, we calculate the potential and pH dependent free energies of all possible intermediates in the decoupled protonelectron ORR and thereby propose simple kinetic models that give semi-quantitative agreement with all experiments. These show that both the reaction mechanism and identity of the rate-limiting step depend on pH. The wide range of agreement between experiments and first-principles theory firmly establish this as a uniquely confirmed example of ORR electrocatalysis driven solely by decoupled proton-electron transfers. This evidence thus embraces the notion of interplaying reaction mechanisms across carbon chemistries and calls for developing new descriptors to understand activity trends and establish design principles in electrocatalysis.51 Both experiment and theory have played an essential role in determining the mechanisms for ORR at different pH. The combination provides evidence that the only viable mechanisms to explain all results involves the mrGO/P50 acting as a co-catalyst. Identifying mrGO/P50 as a co-catalyst is founded on the following findings: i) the considerable mrGO band-gap as calculated from DFT that makes it highly unlikely to initiate the outer-sphere O2 reduction, ii) the identical onset potentials measured for mrGO/P50 and standalone P50 (but the significantly lower activity of stand-alone P50), and iii) the low mrGO electrical conductivity that is


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ACS Catalysis

predicted based on our XPS data combined with experimental studies of rGO conductivity (Fig. S9). As to selectivity in ORR, neither we nor anyone else in the literature has given a theoretical quantitative explanation for selectivity of 2e- vs 4eORR. Since this likely depends upon quantitative electrochemical kinetics and not thermodynamics, this is beyond the current state of the art for ab-initio calculations. There are, however, several generalizations proposed that seem to correlate with H2O2 selectivity that agree with our calculations for ORR on mrGO.50 The predicted weak OOH* binding is generally consistent with H2O2 formation where it appears easier to break the weaker C-O bond rather than the stronger O-O bond. In addition, once OOH* is formed, it is unlikely to dissociate to O* and OH* without significant rearrangements on the surface. A unique aspect of the mrGO/P50 catalyst is that it combines a highly electrical conducting electrode with a semiconducting catalyst. The conducting P50 provides electrons for initiating the ORR by driving the outer-sphere electron transfer to O2(aq), while the semiconducting mrGO provides the active catalytic sites for adsorption of O2-(aq) or HO2(aq), depending upon the pH. Therefore, we have described the mrGO/P50 as a co-catalyst. This concept suggests a general way to combine semiconducting materials with outer-sphere electron transfers to design and optimize new co-catalysts. We suspect that catalytic activity could be even increased if the two materials could be combined on the nm scale instead of the μm scale as done here. In a related context, Surendranath and collaborators have attached specifically designed molecular catalysts to graphite in order to exploit the controllable chemistry of homogenous electrocatalysts, while maintaining the activity and ease of heterogeneous electrocatalysts.52, 53 Again, the graphite provides the electrons for the outer-sphere electrochemistry but the molecular catalyst determines the selectivity and activity for the catalysis.

AUTHOR INFORMATION Corresponding Author * [email protected], [email protected]

Present Addresses # Robert Sengpiel, Chemical Process Engineering, RWTH Aachen University, Forckenbeckstrasse 51, 52074 Aachen, Germany

Author Contributions The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript.

Funding Sources VJB and ACL acknowledge support from a grant of the U.S. Department of Energy, Office of Basic Energy Sciences, to the SUNCAT Center for Interface Science and Catalysis. B.D.M. and H.W.K. gratefully acknowledge support from the National Science Foundation under Grant No. CBET-1604927. This research used resources of the National Energy Research Scientific Computing Center, a DOE Office of Science User Facility supported by the Office of Science of the U.S. Department of Energy under Contract No. DE-AC0205CH11231. R.S. was supported by the Cluster of Excellence "Tailor-made Fuels from Biomass" funded by the Excellence Initiative of the German federal and state governments.

ASSOCIATED CONTENT Supporting Information The Supporting Information is available free of charge on the ACS Publications website. pH-dependent measurements for H2O2 selectivity, Raman spectroelectrochemistry, mrGO/P50 vs P50 activity, O2 reaction orders and H/D isotope effects; DFT set-up and computational details; discussion on calculating the adsorption of the superoxide ion from DFT; methodology for calculating free energies; table of all applied zero-point energy and entropic corrections; supporting data on the mrGO electronic structure; methodology behind the thermodynamic and Pourbaix analyses; force field- based molecular dynamics simulations for simulating the H2O solvent; discussion on the room-temperature structure of interfacial H2O and its effect on adsorbate stability; supporting data on the mrGOORR mechanism; analytic derivations of the ORR reaction in O2 as a function of overpotential in acid and base; atomic coordinates for model structures (PDF).

ACKNOWLEDGMENT We thank H. H. Heenen for help with the force field water simulations. H. Oberhofer and K. Reuter are also thanked for many stimulating discussions.

ABBREVIATIONS ORR oxygen reduction reaction; GO graphene oxide; rGO reduced graphene oxide; mrGO mildy reduced graphene oxide; SHE standard hydrogen electrode; RHE reversible hydrogen electrode; LSV linear sweep voltamettry; (p)DOS (projected) density of states.

REFERENCES 1. Bard, A. J.; Parsons, R.; Jordan, J., Standard Potentials in Aqueous Solution, 1st ed., Marcel. Dekker: New York, 1985, p. 787-802. 2. Campos-Martin, J. M.; Blanco-Brieva, G.; Fierro, J. L. G., Hydrogen Peroxide Synthesis: An Outlook beyond the Anthraquinone Process. Angewandte Chemie International Edition 2006, 45, 6962-6984. 3. Samanta, C., Direct synthesis of hydrogen peroxide from hydrogen and oxygen: An overview of recent developments in the process. Applied Catalysis A: General 2008, 350, 133-149. 4. Rodriguez, P.; Koper, M. T. M., Electrocatalysis on gold. Physical Chemistry Chemical Physics 2014, 16, 13583-13594. 5. Blizanac, B. B.; Ross, P. N.; Markovic, N. M., Oxygen electroreduction on Ag(111): The pH effect. Electrochimica Acta 2007, 52, 2264-2271. 6. Siahrostami, S.; Verdaguer-Casadevall, A.; Karamad, M.; Deiana, D.; Malacrida, P.; Wickman, B.; Escudero-Escribano, M.; Paoli, E. A.; Frydendal, R.; Hansen, T. W.; Chorkendorff, I.; Stephens, I. E. L.; Rossmeisl, J., Enabling direct H2O2 production through rational electrocatalyst design. Nature Materials 2013, 12, 1137-1143. 7. Yang, H. H.; McCreery, R. L., Elucidation of the Mechanism of Dioxygen Reduction on Metal-Free Carbon Electrodes. Journal of The Electrochemical Society 2000, 147, 3420-3428. 8. Yeager, E., Electrocatalysts for O2 reduction. Electrochimica Acta 1984, 29, 1527-1537. 9. Chen, S.; Chen, Z.; Siahrostami, S.; Kim, T. R.; Nordlund, D.; Sokaras, D.; Nowak, S.; To, J. W. F.; Higgins, D.; Sinclair, R.; Nørskov, J. K.; Jaramillo, T. F.; Bao, Z., Defective Carbon-Based Materials for the Electrochemical Synthesis of Hydrogen Peroxide. ACS Sustainable Chemistry & Engineering 2018, 6, 311-317. 10. Wang, H.; Maiyalagan, T.; Wang, X., Review on Recent Progress in Nitrogen-Doped Graphene: Synthesis, Characterization, and Its Potential Applications. ACS Catalysis 2012, 2, 781-794. 11. Dai, L.; Xue, Y.; Qu, L.; Choi, H.-J.; Baek, J.-B., Metal-Free Catalysts for Oxygen Reduction Reaction. Chemical Reviews 2015, 115, 4823-4892. 12. Kong, X.-K.; Chen, C.-L.; Chen, Q.-W., Doped graphene for metal-free catalysis. Chemical Society Reviews 2014, 43, 2841-2857.


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13. Kim, H. W.; Ross, M. B.; Kornienko, N.; Zhang, L.; Guo, J.; Yang, P.; McCloskey, B. D., Efficient hydrogen peroxide generation using reduced graphene oxide-based oxygen reduction electrocatalysts. Nature Catalysis 2018, 1, 282-290. 14. 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. The Journal of Physical Chemistry B 2004, 108, 17886-17892. 15. Kulkarni, A.; Siahrostami, S.; Patel, A.; Nørskov, J. K., Understanding Catalytic Activity Trends in the Oxygen Reduction Reaction. Chemical Reviews 2018, 118, 2302-2312. 16. Suntivich, J.; Gasteiger, H. A.; Yabuuchi, N.; Nakanishi, H.; Goodenough, J. B.; Shao-Horn, Y., Design principles for oxygenreduction activity on perovskite oxide catalysts for fuel cells and metal–air batteries. Nature Chemistry 2011, 3, 546-550. 17. Viswanathan, V.; Hansen, H. A.; Rossmeisl, J.; Nørskov, J. K., Unifying the 2e– and 4e– Reduction of Oxygen on Metal Surfaces. The Journal of Physical Chemistry Letters 2012, 3, 2948-2951. 18. To, J. W. F.; Ng, J. W. D.; Siahrostami, S.; Koh, A. L.; Lee, Y.; Chen, Z.; Fong, K. D.; Chen, S.; He, J.; Bae, W.-G.; Wilcox, J.; Jeong, H. Y.; Kim, K.; Studt, F.; Nørskov, J. K.; Jaramillo, T. F.; Bao, Z., High-performance oxygen reduction and evolution carbon catalysis: From mechanistic studies to device integration. Nano Research 2017, 10, 1163-1177. 19. Lu, Z.; Chen, G.; Siahrostami, S.; Chen, Z.; Liu, K.; Xie, J.; Liao, L.; Wu, T.; Lin, D.; Liu, Y.; Jaramillo, T. F.; Nørskov, J. K.; Cui, Y., High-efficiency oxygen reduction to hydrogen peroxide catalysed by oxidized carbon materials. Nature Catalysis 2018, 1, 156-162. 20. Pegis, M. L.; Wise, C. F.; Martin, D. J.; Mayer, J. M., Oxygen Reduction by Homogeneous Molecular Catalysts and Electrocatalysts. Chemical Reviews 2018, 118, 2340-2391. 21. Koper, M. T. M., Theory of multiple proton-electron transfer reactions and its implications for electrocatalysis. Chemical Science 2013, 4, 2710-2723. 22. Wellendorff, J.; Lundgaard, K. T.; Møgelhøj, A.; Petzold, V.; Landis, D. D.; Nørskov, J. K.; Bligaard, T.; Jacobsen, K. W., Density functionals for surface science: Exchange-correlation model development with Bayesian error estimation. Physical Review B 2012, 85, 235149 1-23. 23. Giannozzi, P.; Baroni, S.; Bonini, N.; Calandra, M.; Car, R.; Cavazzoni, C.; Ceresoli, D.; Chiarotti, G. L.; Cococcioni, M.; Dabo, I.; Dal Corso, A.; de Gironcoli, S.; Fabris, S.; Fratesi, G.; Gebauer, R.; Gerstmann, U.; Gougoussis, C.; Kokalj, A.; Lazzeri, M.; Martin-Samos, L.; Marzari, N.; Mauri, F.; Mazzarello, R.; Paolini, S.; Pasquarello, A.; Paulatto, L.; Sbraccia, C.; Scandolo, S.; Sclauzero, G.; Seitsonen, A. P.; Smogunov, A.; Umari, P.; Wentzcovitch, R.M., QUANTUM ESPRESSO: a modular and open-source software project for quantum simulations of materials. Journal of Physics: Condensed Matter 2009, 21, 395502 1-19. 24. Garrity, K. F.; Bennett, J. W.; Rabe, K. M.; Vanderbilt, D., Pseudopotentials for high-throughput DFT calculations. Computational Materials Science 2014, 446-452. 25. Blum, V.; Gehrke, R.; Hanke, F.; Havu, P.; Havu, V.; Ren, X.; Reuter, K.; Scheffler, M., Ab initio molecular simulations with numeric atom-centered orbitals. Computer Physics Communications 2009, 180, 2175-2196. 26. Perdew, J. P.; Burke, K.; Ernzerhof, M., Generalized Gradient Approximation Made Simple. Physical Review Letters 1996, 77, 38653868. 27. Tkatchenko, A.; Scheffler, M., Accurate Molecular Van Der Waals Interactions from Ground-State Electron Density and Free-Atom Reference Data. Physical Review Letters 2009, 102, 073005 1-4. 28. Berendsen, H. J. C.; Grigera, J. R.; Straatsma, T. P., The missing term in effective pair potentials. The Journal of Physical Chemistry 1987, 91, 6269-6271. 29. Jorgensen, W. L.; Maxwell, D. S.; Tirado-Rives, J., Development and Testing of the OPLS All-Atom Force Field on Conformational Energetics and Properties of Organic Liquids. Journal of the American Chemical Society 1996, 118, 11225-11236. 30. Plimpton, S., Fast Parallel Algorithms for Short-Range Molecular Dynamics. Journal of Computational Physics 1995, 117, 1-19. 31. Lerf, A.; He, H.; Forster, M.; Klinowski, J., Structure of Graphite Oxide Revisited. The Journal of Physical Chemistry B 1998, 102, 4477-4482. 32. Bagri, A.; Mattevi, C.; Acik, M.; Chabal, Y. J.; Chhowalla, M.; Shenoy, V. B., Structural evolution during the reduction of chemically derived graphene oxide. Nature Chemistry 2010, 2, 581-587.

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33. Gómez-Navarro, C.; Weitz, R. T.; Bittner, A. M.; Scolari, M.; Mews, A.; Burghard, M.; Kern, K., Electronic Transport Properties of Individual Chemically Reduced Graphene Oxide Sheets. Nano Letters 2007, 7, 3499-3503. 34. Li, J.-L.; Kudin, K. N.; McAllister, M. J.; Prud’homme, R. K.; Aksay, I. A.; Car, R., Oxygen-Driven Unzipping of Graphitic Materials. Physical Review Letters 2006, 96, 176101-176105. 35. Eda, G.; Chhowalla, M., Chemically Derived Graphene Oxide: Towards Large-Area Thin-Film Electronics and Optoelectronics. Advanced Materials 2010, 22, 2392-2415. 36. Lundie, M.; Sljivancanin, Z.; Tomic, S., Electronic and optical properties of reduced graphene oxide. Journal of Materials Chemistry C 2015, 3, 7632-7641. 37. Bielski, B. H. J.; Allen, A. O., Mechanism of the disproportionation of superoxide radicals. The Journal of Physical Chemistry 1977, 81, 1048-1050. 38. Kuo, I. F. W.; Tobias, D. J., Thermal Fluctuations of the Unusually Symmetric and Stable Superoxide Tetrahydrate Complex: An ab Initio Molecular Dynamics Study. The Journal of Physical Chemistry A 2002, 106, 10969-10976. 39. Kelly, C. P.; Cramer, C. J.; Truhlar, D. G., Aqueous Solvation Free Energies of Ions and Ion−Water Clusters Based on an Accurate Value for the Absolute Aqueous Solvation Free Energy of the Proton. The Journal of Physical Chemistry B 2006, 110, 16066-16081. 40. Viswanathan, V.; Pickrahn, K. L.; Luntz, A. C.; Bent, S. F.; Nørskov, J. K., Nanoscale Limitations in Metal Oxide Electrocatalysts for Oxygen Evolution. Nano Letters 2014, 14, 5853-5857. 41. 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. The Journal of Physical Chemistry C 2014, 118, 6706-6718. 42. Hartnig, C.; Koper, M. T. M., Molecular dynamics simulation of the first electron transfer step in the oxygen reduction reaction. Journal of Electroanalytical Chemistry 2002, 532, 165-170. 43. Fletcher, S., Tafel slopes from first principles. Journal of Solid State Electrochemistry 2009, 13, 537-549. 44. Holewinski, A.; Linic, S., Elementary Mechanisms in Electrocatalysis: Revisiting the ORR Tafel Slope. Journal of The Electrochemical Society 2012, 159, H864-H870. 45. Tse, E. C. M.; Hoang, T. T. H.; Varnell, J. A.; Gewirth, A. A., Observation of an Inverse Kinetic Isotope Effect in Oxygen Evolution Electrochemistry. ACS Catalysis 2016, 6, 5706-5714. 46. Winkler, F. J., Reaction Rates of Isotopic Molecules. Von L. Melander und W. H. Saunders, Jr. Wiley, New York 1980. XIV, 391 S., geb. £ 16.30. Angewandte Chemie 1981, 93, 220-220. 47. Malko, D.; Kucernak, A., Kinetic isotope effect in the oxygen reduction reaction (ORR) over Fe-N/C catalysts under acidic and alkaline conditions. Electrochemistry Communications 2017, 83, 67-71. 48. McCreery, R. L., Advanced Carbon Electrode Materials for Molecular Electrochemistry. Chemical Reviews 2008, 108, 2646-2687. 49. Xu, J.; Huang, W.; McCreery, R. L., Isotope and surface preparation effects on alkaline dioxygen reduction at carbon electrodes. Journal of Electroanalytical Chemistry 1996, 410, 235-242. 50. Yang, S.; Verdaguer-Casadevall, A.; Arnarson, L.; Silvioli, L.; Čolić, V.; Frydendal, R.; Rossmeisl, J.; Chorkendorff, I.; Stephens, I. E. L., Toward the Decentralized Electrochemical Production of H2O2: A Focus on the Catalysis. ACS Catalysis 2018, 8, 4064-4081. 51. Hong, W. T.; Stoerzinger, K. A.; Lee, Y.-L.; Giordano, L.; Grimaud, A.; Johnson, A. M.; Hwang, J.; Crumlin, E. J.; Yang, W.; Shao-Horn, Y., Charge-transfer-energy-dependent oxygen evolution reaction mechanisms for perovskite oxides. Energy & Environmental Science 2017, 10, 2190-2200. 52. Fukushima, T.; Drisdell, W.; Yano, J.; Surendranath, Y., Graphite-Conjugated Pyrazines as Molecularly Tunable Heterogeneous Electrocatalysts. Journal of the American Chemical Society 2015, 137, 10926-10929. 53. Ricke, N. D.; Murray, A. T.; Shepherd, J. J.; Welborn, M. G.; Fukushima, T.; Van Voorhis, T.; Surendranath, Y., Molecular-Level Insights into Oxygen Reduction Catalysis by Graphite-Conjugated Active Sites. ACS Catalysis 2017, 7, 7680-7687.


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a)

i(mA)

i(mA)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25

−0.5 −1.0 −1.5 −2.0

−0.4 −0.2

0.0 0.2 0.4 U (VRHE )

0.6

0.8

0.0

b) −0.5 −1.0 −1.5 −2.0

pH pH pH pH

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0.6

i (mA)

Page 0.0 15 of 21 ACS Catalysis pH(D)=1 pH(D)=13 −0.5 1 2 −1.0 3 4 5 −1.5 6 7 −2.0 ACS Paragon Plus Environment −0.4 −0.2 0.0 0.2 0.4 0.6 8 9 U (VRHE )

0.8

−2

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−3

log[i(A)]

1 2 −4 3 4 −5 5 6 −6 7 8 9

pH 1 pH 13

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0.1

Reaction order, ρ

0.917 of 21 Page ACS Catalysis a) experiment

Reaction order, ρ

0.6 1 2 0.3 3 4 0.0 5 6 7 0.9 b) model 8 9 0.6 10 11 0.3 12 13 0.0 ACS Paragon Plus Environment −0.3 −0.2 −0.1 14 −0.4 η (V) 15

pH 1 pH 13

0.0

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pDOS (states/eV)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17

C(pACS x)

C(py ) Catalysis

C(pz )

a) clean mrGO

5 0 10

b) OH@ mrGO

5 0 10

c) 2OH@ mrGO

5 0 ACS Paragon Plus Environment −4 −2 0 2 ε − εF (eV)

4

∆GOH∗ (eV)

1.5 1.0

0.5 1 0.0 2 3 −0.5 4 −1.0 5 6 7

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∆G (eV)

∆G (eV)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23

1.5 1.0 0.5 0.0 −0.5 −1.0 −1.5 −2.0 0.5 0.0 −0.5 −1.0 −1.5 −2.0 −2.5 −3.0 −3.5

ACS Catalysis

a) pH=1

HO2 + ∗ + (H+ + e− ) O2 + ∗ + 2(H+ + e− )

+ − O− 2 + ∗ + (2H + e )

OOH∗ + (H+ + e− )

H 2 O2 + ∗

b) pH=13 [O2 ∗ ]− + (H2 O + e− ) − − O2 + ∗ + (H2 O + 2e− ) O2 + ∗ + (H2 O + e )


Reaction −→

OOH∗ + (OH− + e− )

HO2 − + ∗ + (OH− )

Combining Experiment and Theory to Unravel the Mechanism of Two

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