Understanding Catalytic Activity Trends in the Oxygen Reduction

Feb 6, 2018 - it is clear that developments in the existing material classes can only lead to marginal improvements in the inherent catalytic activity...
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Review Cite This: Chem. Rev. XXXX, XXX, XXX−XXX

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Understanding Catalytic Activity Trends in the Oxygen Reduction Reaction Ambarish Kulkarni,†,§ Samira Siahrostami,†,§ Anjli Patel,† and Jens K. Nørskov*,†,‡ †

SUNCAT Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, 450 Serra Mall, Stanford, California 94305, United States ‡ SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, Menlo Park, California 94025, United States ABSTRACT: Despite the dedicated search for novel catalysts for fuel cell applications, the intrinsic oxygen reduction reaction (ORR) activity of materials has not improved significantly over the past decade. Here, we review the role of theory in understanding the ORR mechanism and highlight the descriptor-based approaches that have been used to identify catalysts with increased activity. Specifically, by showing that the performance of the commonly studied materials (e.g., metals, alloys, carbons, etc.) is limited by unfavorable scaling relationships (for binding energies of reaction intermediates), we present a number of alternative strategies that may lead to the design and discovery of more promising materials for ORR.

CONTENTS

1. INTRODUCTION

1. Introduction 2. Free Energy Diagram and ORR Mechanism 2.1. Four-Electron Reduction to H2O 2.2. Two-Electron Reduction to H2O2 3. Understanding ORR Trends for Transition Metals 3.1. Linear Scaling Relationships 3.2. Limiting Potential and Volcano Plots 3.3. Volcano Plots for Two-Electron Reduction 4. Kinetics of ORR 4.1. Microkinetic Modeling of ORR over Pt(111) 4.2. Rationalizing the Two- vs Four-Electron Selectivity 5. Improving ORR Activity: Alloys and Core−Shell Catalysts 5.1. Thermodynamic Analyses Based on Limiting Potentials 5.2. Related Approaches beyond Transition Metals and Pt Alloys 5.3. Comparing Trends from Kinetic and Thermodynamic Analyses 6. Perspectives and Outlook Author Information Corresponding Author ORCID Author Contributions Notes Biographies Acknowledgments References © XXXX American Chemical Society

Radical innovations in energy generation and utilization technologies are necessary to ensure a safe and sustainable future.1 One aspect of this utopia involves exploiting renewable energy sources such as solar, wind, and hydroelectric power to generate “green” hydrogen and liquid fuels (via electrocatalytic water splitting and reduction reactions) that can be assimilated in the existing energy and transportation infrastructure. In addition, a suite of highly efficient, chemical-to-electrical conversion technologies are required. Fuel cell devices involve electrochemical oxidation of a fuel (typically hydrogen, methanol, etc.) and are a promising avenue for small to intermediate scale operations. The oxygen reduction reaction (ORR) is at the heart of many such fuel cell devices and holds a special place in the field of electrocatalysis. In the ORR, molecular oxygen is electrochemically reduced by four protons and electrons to form water, which is accompanied by generation of an electrical potential. For low-temperature proton exchange membrane (PEM) fuel cells, the lack of a sufficiently good ORR catalyst severely limits the overall efficiency of the device.2−4 Thus, the discovery and optimization of novel fuel cell catalysts has been a consistent driving force for academic and industrial research all over the world. Moreover, ORR is often viewed as a prototype, multielectron process and has been widely used to

A B C C C C C D D D E E E F F F G G G G H H H H

Special Issue: Oxygen Reduction and Activation in Catalysis Received: August 14, 2017

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DOI: 10.1021/acs.chemrev.7b00488 Chem. Rev. XXXX, XXX, XXX−XXX

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comparisons to theoretical work (various computational approaches have been reviewed by Keith et al.14). Depending on the catalyst, both the two- and four-electron oxygen reduction pathways are possible (eqs 1−5). The two-electron, partial reduction pathway results in hydrogen peroxide (H2O2) as the product and only involves OOH* as a reaction intermediate, while the full reduction of oxygen to water (eq 3) involves four reduction steps. Typically, both associative and dissociative mechanisms are considered, depending on whether the O2 molecule dissociates before reduction. The associative mechanism (eq 4) involves three different intermediates, namely, *OOH, *O, and *OH, while the dissociative pathway only involves *O and *OH (eq 5).

develop a number of concepts that are broadly applicable to other electrocatalytic reactions.5 In this paper, we summarize some of these important concepts and describe our current understanding of the ORR mechanism over the most commonly studied metal catalysts. Over the past decade, density functional theory (DFT) calculations have shown that the energetics of key reaction intermediates are linked by linear scaling relationships that have been indispensable for rationalizing trends in ORR catalytic activity across different materials.5 These analyses often result in a “volcano relation” between catalytic activity and key adsorption energies, which, for instance, explains why Pt and Pd and their alloys are the best catalysts for ORR. More importantly, it also reveals the fundamental factors limiting the efficiency of the known catalysts, explaining why the past decade of catalyst development has resulted in only a modest improvement in the intrinsic activity (Figure 1). Development

O2 + 2(H+ + e−) → H 2O2

E° = 0.70 V

(1)

O2 + 2(H+ + e−) → *OOH + (H+ + e−) → H 2O2 (2)

O2 + 4(H+ + e−) → 2H 2O

E° = 1.23 V

(3)

O2 + 4(H+ + e−) → *OOH + 3(H+ + e−) → *O + 2(H+ + e−) → *OH + 1(H+ + e−) → 2H 2O (4)

1 O + 2(H+ + e−) → *O + 2(H+ + e−) → 2 2 *OH + (H+ + e−) → H 2O

(5)

In principle, we should include the adsorption of molecular O2 on the surface as well, but we view this as a chemical step, and the steady-state adsorption of such a neutral species does not contribute to the potential dependence of the faradic current. The molecular levels of adsorbed O2 couple strongly enough to the metal electrons that electron transfer is unlikely to be rate limiting for this process. We initially only consider the free energies of the intermediates, as a considerable amount of understanding can be achieved on that basis. The electrocatalytic current depends on the “rate” of the reaction, which in turn depends on the potential-dependent activation energy barriers. These barriers are important for describing the ORR kinetics and selectivity and will be discussed in section 4. However, specifically for the ORR on metal surfaces, as the electron−proton transfers to oxygen from hydronium are very facile,15 most potential dependence comes from the energy of the intermediates. DFT calculations work well in describing adsorption energies on metal surfaces;16 the challenge has been to describe the energy of the solvated protons and the electrons in the electrode at a given potential. In the computational hydrogen electrode (CHE) model,5 the free energy of a single proton−electron pair is defined as −eU relative to H2 in the gas phase at standard conditions, where U is the electrode potential with respect to the reversible hydrogen electrode (RHE). If we therefore add the effects of adsorbate solvation (ΔEw),15 electric field effects (Efield),17 zero-point energy (ΔZPE),5 and entropic corrections (−TΔS)5 to a DFT calculated binding energy (ΔEele), we can describe the adsorption free energies of a reaction intermediate (with n proton−electron pairs) as a function of the potential by

Figure 1. Timeline showing that the improvements in ORR performance of materials has plateaued owing to the scaling relationships between ORR intermediates. Nontraditional electrocatalysts and/or process configurations are necessary to make disruptive advances in this field. Adapted with permission from ref 6. Copyright 2017 American Association for the Advancement of Science.

of completely new types of catalysts with active sites that do not obey the known scaling relations is identified as a crucial bottleneck in catalyst development, and a number of alternative strategies are discussed. Although high-temperature fuel cell designs (based on solid oxides) are also an active area of research, their performance is generally not limited by catalysis,7 and they are beyond the scope of this review.8 Thus, we limit our discussion to the lowtemperature PEM devices, as the catalytic turnover of the material is critical for improved performance and decreased cost. While other factors such as material cost and availability, density of active sites, diffusion properties, electrical conductivity, and catalyst stability will be addressed throughout the paper, the focus will be on the intrinsic catalyst performance.

2. FREE ENERGY DIAGRAM AND ORR MECHANISM An obvious starting point for understanding ORR mechanism is the Pt(111) surface. Owing to the large volume of available electrochemical,3,9 structure sensitivity,10 single-crystal,11,12 and spectroscopic13 data, Pt(111) is an ideal benchmark for

ΔG = ΔEele + ΔEw + ΔEfield + ΔZPE − T ΔS − neU (6) B

DOI: 10.1021/acs.chemrev.7b00488 Chem. Rev. XXXX, XXX, XXX−XXX

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Figure 2. (a) Free energy diagram for the four-electron associative ORR on Pt(111)22 and (b) two-electron ORR on PtHg4.21 The free energy diagram is shown at three different potentials: 0 V (blue lines), the corresponding equilibrium potential (green lines), and the limiting potential (black lines). The blue-green arrows indicate the effect of the potential based on the CHE model.5

2.1. Four-Electron Reduction to H2O

electron route, the limiting potential is obtained where all the steps are downhill in free energy, ηtheo ≈ 0.70−0.63 V = 0.07 V.

Figure 2 shows the DFT-calculated free energy diagram for the four-electron ORR over Pt(111) using the CHE model. As the dissociative process is unlikely to be important for Pt(111),5 only the associative pathway is shown. At U = 0.0 V, all the reaction steps are downhill, implying a facile reaction (blue lines, Figure 2a). To model the thermodynamics at the equilibrium potential (i.e., U = 1.23 V), eq 6 suggests that the adsorption energies are shifted by −neU, depending on the number, n, of H+ + e− pairs (blue-green arrows, Figure 2a) corresponding to each intermediate. Figure 2 shows that, at U = 1.23 V (green lines), the reduction of *OH to H2O is uphill in energy, indicating that the surface will be covered by these species and will be inactive for O2 adsorption. The maximum potential where all steps are downhill in free energy is ca. 0.8 V, which agrees well with the experimentally observed maximum potential for large ORR current densities.9,12,18,19 Details of the picture developed on the basis of DFT calculations have been confirmed in operando experimental studies of surface coverages on Pt(111) during the reaction.13 Within the CHE model, the highest potential at which all the reaction steps are downhill in free energy is called the thermodynamic limiting potential (UL) and corresponds to 0.8 V for Pt(111); see Figure 2. The difference between the equilibrium potential of U = 1.23 V and the limiting potential is called the theoretical overpotential, and ηtheo ≈ 1.2−0.8 V = 0.4 V for Pt(111). Note that ηtheo (as well as UL) is just a measure of the activity of a catalyst and should not be compared directly with a measured overpotential, which of course depends on the current density.

3. UNDERSTANDING ORR TRENDS FOR TRANSITION METALS 3.1. Linear Scaling Relationships

In addition to improving the understanding of the ORR mechanism on Pt(111), the CHE-based thermochemical theory has been widely used to explain the trends in catalytic activity for a range of materials, including transition metals, alloys,23 oxides,24−26 perovskites,27 sulfides,28 nitrides,29 and carbonbased materials.30 As the associative ORR pathway involves three intermediates, the theoretical overpotential is a function of three catalystdependent binding free energies. However, as shown in Figure 3a, the binding energies for *OOH, *O, and *OH are strongly correlated and change monotonically for different metals. These linear relationships arise because all the adsorbates bind to the surface through an O atom. In particular, the *OOH vs *OH line has a slope close to unity, indicating a similar metal− oxygen single bond for both adsorbates. In contrast, the *O vs *OH scaling line has a slope close to 2, in line with a picture where the adsorbed O has a double bond to the surface while *OH binds through a single bond.31 More broadly, linear scaling relationships and descriptor-based analysis are important concepts in the field of heterogeneous catalysis and have been crucial for understanding catalytic activity trends for a range of materials and chemical reactions.31−33 3.2. Limiting Potential and Volcano Plots

The *OOH vs *OH or *O scaling relations imply that, to a first approximation, there is only one independent variable that describes the binding free energies of all intermediates to the catalyst surface. Although the initial work using this approach was based on the *O binding free energy (ΔGO),5 the *OH binding free energy (ΔGOH) has been more popular recently (owing to a better scaling between *OOH and *OH; see Figure 3a).34 As the theoretical overpotential is a function of the binding energies of *OOH, *O, and *OH, we can use the scaling relations in Figure 3a to define the limiting potentials for the four steps (UL1−L4) as a function of the OH free energy of adsorption as

2.2. Two-Electron Reduction to H2O2

Before moving further, it is useful to briefly examine the twoelectron ORR, which leads to H2O2 as the product. Although the two-electron pathway is undesirable for a fuel cell application, electrochemical H2O2 production is an emerging reaction of interest for onsite applications such as water disinfection.20 Here the sole intermediate is *OOH, and an analogous CHE analysis can be used to develop a thermodynamic picture. The free energy diagram for the twoelectron ORR to H2O2 is shown in Figure 2b for the state-ofthe-art catalyst PtHg4.21 At U = 0.0 V, the reaction is facile (blue lines), but at the equilibrium potential, U = 0.70 V (green lines), formation of OOH* is slightly uphill. Similar to the four-

UL1 = −ΔGOH + 1.72 C

(7) DOI: 10.1021/acs.chemrev.7b00488 Chem. Rev. XXXX, XXX, XXX−XXX

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of the other transition metals, that of Pt lies closest to the top of the limiting potential volcano, in excellent agreement with the experimentally observed trends.9,35−40 In addition to explaining the trends among different metals, this analysis can be extended to explain the structure sensitivity of ORR. As the other fairly close-packed transition-metal facets approximately follow the (111) *OOH vs *OH scaling line,23 they can be described by the 1-D volcano in Figure 3b. Compared to the (111) terraces, the bonding of *OH on the step sites (represented by (211) facets) is considerably stronger (more negative ΔGOH). For instance, although Pt(111) shows excellent ORR activity, the stronger bonding at the (211) step sites leads to poisoning of the surface by adsorbed *OH (red arrow, Figure 3b). This explains the decrease in the activity of smaller Pt particles, as smaller particle sizes lead to higher numbers of inactive step sites.41,42 On the other hand, for a catalyst on the weak binding side (such as Au), steps may enhance the activity and a smaller particle size may be desirable.43,44 We use the *OH adsorption energy as a descriptor of ORR catalytic activity throughout, but it is important to note that this is just a convenient way of monitoring the underlying variations in the surface electronic structure from one catalyst surface to the next. In fact, the *OH and *O adsorption energy varies essentially linearly with the center of the metal d states associated with the bonding; the higher the energy of the d states, the higher the reactivity and, subsequently, the more negative the ΔGOH.45−47 3.3. Volcano Plots for Two-Electron Reduction

For the two-electron process, the limiting steps are either formation or removal of *OOH from the surface. Here, the theoretical overpotential is only a function of the *OOH binding energy. On the basis of the scaling relations, the *OOH binding energy is related to the *OH binding energy, and the limiting potentials can be expressed as Figure 3. (a) Scaling relationships for the chemisorption energies of *OOH and *O for the (111) surface of various metals (black circles) using *OH as a descriptor. The solid black lines correspond to ΔGOOH = ΔGOH + 3.2 and ΔGO = 2ΔGOH. (b) Limiting potentials for individual steps in eqs 7−10, showing the strongly bound *OH region (solid blue line) and weakly bound *OOH region (solid green line) for the four-electron process. (c) Limiting potentials for individual steps in eqs 11 and 12, showing the strongly bound *OH region (solid purple line) and weakly bound *OOH region (solid green line) for the two-electron process. The color gradient indicates the strong *OH and weak *OOH binding regions.

UL2 = −ΔGOH + 3.3

(8)

UL3 = ΔGOH

(9)

UL4 = ΔGOH

(10)

UL1 ′ = −ΔGOH + 1.72

(11)

UL2 ′ = ΔGOH − 0.32

(12)

Figure 3c shows the two-electron volcano (purple line). Since this reaction has only one intermediate (*OOH), the peak of the volcano crosses the equilibrium potential at 0.70 V. This indicates that it is possible, in principle, to find a catalyst with an ideal activity if it binds that single intermediate with optimal strength, not too weak nor too strong. This has been observed for Pt and Pd mercury alloys,21,48 Co−porphyrin,49 and some defective carbon structures.30 Obviously, catalyst surfaces with strong oxygen bonding energy lying on the left side of the two-electron volcano are not efficient for the twoelectron ORR route due to dissociation of the adsorbed O2 or *OOH or the product H2O2.

The above equations result in four limiting potential lines for the four elementary steps as shown in Figure 3b. The lowest limiting potential for the full catalytic reaction defines the overall limiting potential for the reaction, and it is indicated by the blue and green solid lines. For the associative mechanism, we find that the first and last steps in eq 4 are potential limiting for all the transition metals (shown by solid circles, Figure 3b). For metals that bind *OH strongly, *OH→ H2O is potential limiting (solid blue line), whereas for the weakly bonding metals the activation of O2, O2 → *OOH, is potential limiting (solid green line). Encouragingly, compared to the (111) facets

4. KINETICS OF ORR 4.1. Microkinetic Modeling of ORR over Pt(111)

The thermodynamic analysis based on theoretical limiting potentials has played a fundamental role in rationalizing the activity trends for metals and in guiding the design and optimization of various catalysts.6,47 As the ORR current ultimately depends on the rate of the reaction and on electrochemical barriers, we now briefly discuss the role of kinetics in the ORR. We note that calculation of electrochemical barriers and solvation effects is an active research area, D

DOI: 10.1021/acs.chemrev.7b00488 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

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and a variety of computational approaches ranging from implicit solvation50 to explicit solvation models, including few water molecules51,52 or multiple water layers,53 have been used for describing the ORR kinetics. Using an atomistic model of the charged solid−liquid interface,54 the barriers for proton transfer to surface-bound intermediates have been estimated to be quite small, essentially of the same order of magnitude as proton transfer barriers in water.15 Subsequently, a microkinetic model for the ORR based on the DFT calculations was shown to reproduce the experimental kinetic observations (cyclic voltammetry,55 linear sweep voltammetry,53 and Tafel slopes) for Pt(111) (Figure 4).56 Unsurprisingly, as the surface reaction is fast at low

Figure 5. Free energy diagram for the four- and two-electron oxygen reduction in blue and red, respectively, on Au(111). The electrochemical barriers for *OOH to H2O2 or *O are illustrative and indicate the importance of kinetics in determining catalyst selectivity.

As shown by the above equations, the key to avoid the fourelectron pathway is to prevent the O−O bond dissociation in the adsorbed *OOH. This simple analysis excludes all catalysts with strong oxygen binding energies (due to favorable O* formation) and limits the search to materials that bind oxygen weakly. This is evidenced by the numerous experimental observations on weak oxygen binding catalysts such as Au(111)58,59 and carbon-based materials60,61 that selectively reduce O2 to H2O2. Despite several experimental studies,18,59,62 the factors determining the selectivity on Au remain elusive.57 On the basis of a purely thermodynamic analysis, the dissociation of *OOH is more favorable over H2O2 formation (Figure 5) and incorrectly implies a low H2O2 selectivity.57 This clearly shows that thermodynamics is insufficient to capture the experimentally observed trends. Early microkinetic modeling studies56 indicate that kinetic parameters have a high degree of rate control and are essential for describing the selectivity.63 Figure 5 also shows representative electrochemical barriers that would explain the experimentally observed selectivity. Moreover, related phenomena, including the pH-dependent selectivity toward H2O2 or H2O on Au(100)64−67 and carbon-based materials,68−71 are of both fundamental and practical interest and require evaluation of potential-dependent barriers in both alkaline and acidic media.

Figure 4. Comparison of the simulated (black line) and experimentally measured (green and gray lines) polarization curves for Pt(111). The inset shows that the predicted Tafel slope (∼59 mV/decade) agrees well with the experiments. Adapted from ref 56. Copyright 2014 American Chemical Society.

potentials (