Electrocatalytic Oxidation of Ammonia on Transition-Metal Surfaces: A

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Electrocatalytic Oxidation of Ammonia on Transition-Metal Surfaces: A First-Principles Study Jeffrey A. Herron, Peter Ferrin, and Manos Mavrikakis* Department of Chemical and Biological Engineering, University of Wisconsin − Madison, Madison, Wisconsin 53706, United States S Supporting Information *

ABSTRACT: We investigate the catalytic electro-oxidation of ammonia on model close-packed surfaces of Au, Ag, Cu, Pd, Pt, Ni, Ir, Co, Rh, Ru, Os, and Re to derive insights for the reaction mechanism and evaluate the catalysts based on their energy efficiency and activity in the context of their application in fuel cells. Two mechanisms, which are differentiated by their N−N bond formation step, are compared: (1) a mechanism proposed by Gerischer and Mauerer, whereby the N−N bond formation occurs between hydrogenated NHx adsorbed species, and (2) a mechanism in which N−N bond formation occurs between N adatoms. The results of our study show that the mechanism proposed by Gerischer and Mauerer is kinetically preferred and that the formation of N adatoms poisons the surface of the catalyst. On the basis of a simple Sabatier analysis, we predict that Pt is the most active monometallic catalyst followed by Ir and Cu, whereas all other metal surfaces studied here have significantly lower activity. We conclude by outlining some design principles for bimetallic alloy catalysts for NH3 electro-oxidation.

1. INTRODUCTION The catalytic electro-oxidation of ammonia (NH3) is a fundamentally important reaction for a number of applications including electrochemical wastewater cleanup,1,2 chemical sensing,3,4 and onboard hydrogen generation.5,6 In addition, there has been interest in using ammonia as an alternative feed to hydrogen for direct fuel cell applications.7,8 Other alternative fuel cell feeds such as methanol address some of the transportation and storage issues of hydrogen9,10 but suffer from CO poisoning of their anode electrocatalysts11 and methanol crossover to the cathode.12,13 NH3 presents a promising alternative feedstock because it is easily liquefied and its electro-oxidation products (dinitrogen and water) are environmentally benign. Regardless of the catalytic reaction of interest, a fundamental understanding of its reaction mechanism is essential for designing catalysts with improved performance, e.g., activity, selectivity, and reduced cost.14−17 However, in the case of NH3 electro-oxidation, the mechanism is not well-understood and the reaction is slow on typical Pt-based catalysts.7,18 The mostaccepted mechanism was proposed by Gerischer and Mauerer in 19708 with the proposed elementary steps outlined in Scheme 1. NH3 is deprotonated by hydroxyl ions in steps 1, 2, and 3, forming water molecules while simultaneously releasing an electron at each step. N adatoms (formed in step 3) are surface poisons because of a typically large kinetic barrier for N−N bond formation to release N2. Therefore, adsorbed NHx (and NHy) species react with one another to form an N−N bond, thereby producing an HxNNHy species. According to the Gerischer−Mauerer mechanism, this species is then deproto© XXXX American Chemical Society

Scheme 1. Electro-Oxidation Mechanism Proposed by Gerischer and Mauerera

a

The * indicates free surface sites or adsorbed intermediates.

nated to N2, which desorbs from the surface. However, the identity of the NHx and NHy species that react to form the N− N bond remains in dispute. Experimental electrochemical studies have provided important mechanistic insights into the reaction mechanism. Gootzen et al. used differential electrochemical mass spectroscopy (DEMS) and cyclic voltammetry to study the coverage of adsorbates on a platinized platinum electrode.19 Assuming that the most abundant surface intermediate is atomic nitrogen, they estimated a surface coverage of 0.6 ML. De Vooys et al. studied Special Issue: Steven J. Sibener Festschrift Received: December 30, 2014 Revised: February 6, 2015

A

DOI: 10.1021/jp512981f J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C

whereby N−N bond formation is restricted to occur through adsorbed N adatoms. Hereafter, we will refer to the latter mechanism as the “N + N mechanism”. For the metal surfaces studied, we determine the most-favorable reaction pathway(s), estimate activity, and predict ranges of operating potential, as defined by the onset potential of the two mechanisms.

ammonia electro-oxidation on a series of transition-metal (Pt, Pd, Rh, Ru, Ir, Cu, Ag, and Au) flag electrodes using these same two techniques.20 Ir was the most active electrode, while Pt was the second most active. Pd, Rh, and Ru were all less active than both Pt and Ir. The coinage metals (Cu, Ag, Au) were inactive for N2 formation, but oxygenated nitrogen species (NO, N2O) were detected at high potential on these metals. Platinum reached a steady-state current up to 0.57 V versus the reversible hydrogen electrode (RHE), above which deactivation occurs. A similar deactivation was seen on Ir above 0.50 V. The structure sensitivity of ammonia electro-oxidation on Pt has been well documented, showing superior activity for Pt(100) than Pt(111).21−26 Rosca and Koper studied the structure sensitivity of ammonia electro-oxidation on Pt(111) and Pt(100) surfaces using cyclic voltammetry, chronoamperometry, and in situ infrared spectroscopy.26 On Pt(111) they found low activity, with NH and N being the dominant surface intermediates. The activity was higher on Pt(100), where adsorbed NH2 was the dominant surface intermediate. On the basis of these results and Tafel slope analysis, they concluded that dimerization of adsorbed NH2 species to hydrazine (H2NNH2) is the rate-determining step for NH3 electrooxidation. Furthermore, they attributed the low activity of Pt(111) to the slow kinetics of N−N bond formation via adsorbed NH or N species which are dominant on that surface. Other research groups have taken advantage of the improved activity of the Pt(100), synthesizing preferentially oriented Pt electrodes.7,21,25,27 Going beyond monometallic catalysts, there have been several experimental studies aimed at developing alloy catalysts for NH3 electro-oxidation.18 A number of studies have shown that PtIr alloys show improved activity when compared to pure Pt.28−35 Endo et al. studied PtIr and PtCu alloy electrodes and showed that Pt0.8Ir0.2 have higher activity and a lower ammonia oxidation potential when compared to Pt.32 In contrast, PtCu electrodes had oxidation potentials similar to those of pure Pt electrodes, yet lower current density. Lomocso and Baranova developed Pt7Pd3 and Pt7Ir3 nanoparticle catalysts and showed that both formulations had improved activity when compared with pure Pt.30 Alloying Ir with Pt also increased the stability of the catalyst, while addition of Pd led to reduced stability. Yao and Cheng alloyed Pt with Ni, producing catalysts with activity similar to that of pure Pt, but with reduced Pt loading.36 In order to design improved alloy catalysts for ammonia electro-oxidation, it is essential to elucidate the detailed atomicscale reaction mechanism and its rate-determining step and to understand how these change with the binding properties of the metal surface. Toward this goal, we present a first-principles density functional theory (DFT) study of ammonia electrooxidation on the close-packed facets of Au, Ag, Cu, Pt, Pd, Ni, Ir, Rh, Co, Os, Ru, and Re. Though experimental studies have demonstrated the structure sensitivity of this reaction,21−26,37 and understanding the structure sensitivity is a longer term goal, we begin our analysis of this reaction by focusing on the most thermodynamically stable, close-packed crystal facet. Importantly, we do not know a priori that a more open (e.g., (100)) facet of all of these metals is more active than the corresponding close-packed facet. Furthermore, as the closepacked facet is most thermodynamically stable, it would be beneficial to design alloy catalysts that are active on the closepacked facet. In studying the electro-oxidation of ammonia, we consider two mechanistic pathways: (1) the mechanism proposed by Gerischer and Mauerer and (2) a mechanism

2. METHODS The free energies of reaction species are calculated using spinpolarized, planewave density functional theory, as implemented in DACAPO,38,39 a total energy code. Calculations are performed on the (111) facet of face-centered cubic (fcc) metals (here, Rh, Ir, Ni, Pd, Pt, Cu, Ag, and Au) and the (0001) facet of hexagonal close-packed (hcp) metals (here, Re, Ru, Os, Co). We drop the facet notation hereafter for brevity. The catalysts are modeled using a periodic 3 × 3 unit cell (hence a surface coverage of 1/9 ML for each individual adsorbed species) with three layers of metal atoms and five equivalent layers of vacuum between successive slabs. The metal atoms are fixed at their optimized bulk positions. The optimized bulk lattice constants are (experimental value40 in parentheses): Ag 4.14 Å (4.09 Å), Au 4.18 Å (4.08 Å), Co 2.50 Å (2.51 Å), Cu 3.67 Å (3.62 Å), Ir 3.86 Å (3.83 Å), Ni 3.52 Å (3.52 Å), Os 2.74 Å (2.73 Å), Pd 3.99 Å (3.89 Å), Pt 4.00 Å (3.92 Å), Re 2.76 Å (2.76 Å), Rh 3.83 Å (3.80 Å), and Ru 2.74 Å (2.71 Å). For Co, Os, Ru, and Re, we used a c/a ratio of 1.63. Adsorption is permitted on only one of the two exposed surfaces, and the dipole moment41,42 is adjusted accordingly. The ionic cores are described using ultrasoft Vanderbilt43 pseudopotentials. The Kohn−Sham one-electron states are expanded in a series of plane waves with an energy cutoff of 25 Ry. The surface Brillouin zone is sampled with 18 special Chadi−Cohen44 kpoints based on convergence of the total energy. The exchange−correlation energy and potential are self-consistently calculated using the PW91 generalized-gradient approximation.45 The electron density is determined by iterative diagonalization of the Kohn−Sham Hamiltonian, Fermi population of the Kohn−Sham states (kBT = 0.1 eV), and Pulay mixing of the resulting electronic density. All total energies are extrapolated to kBT = 0 eV. Convergence with respect to all calculation parameters, including the effect of surface atom relaxation, has been thoroughly checked. The minimum-energy pathways and corresponding activation energy barriers for all N−N bond formation elementary steps are calculated using the climbing image nudged elastic band method (CI-NEB).46 The reaction coordinate for each elementary step considered is discretized with nine images, including end points. All transition states were confirmed by vibrational frequency calculations yielding a single negative curvature mode. These transition states were calculated for only Cu, Pt, Pd, Co, Ir, and Rh. The specific metals were chosen in order to study the activity over a wide-range of binding energies. The zero-point energy (ZPE) is included in the calculation of free energy of adsorbates. The zero-point energies are calculated assuming a quantum harmonic oscillator with the calculated vibrational frequencies. The vibrational frequencies are calculated by numerical differentiation of forces using a second-order finite difference approach with a step-size of 0.010 Å. The Hessian matrix is mass-weighted and diagonalized to yield the frequencies and normal modes of the adsorbed species. The entropies of all surface species are calculated by B

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The Journal of Physical Chemistry C Table 1. Gibbs Free Energy of Adsorbed Intermediates at 0 V Relative to N2(g), H2O(g), and OH−a Ag Au Cu Pd Pt Ni Ir Co Rh Ru Os Re

N

NH

NH2

NNH

HNNH

NNH2

HNNH2

H2NNH2

3.20 3.16 1.94 0.94 0.79 0.23 0.13 −0.11 −0.14 −0.30 −0.82 −1.49

1.75 2.11 0.71 0.40 0.27 −0.33 −0.37 −0.50 −0.39 −0.69 −0.99 −1.62

0.74 1.06 0.27 0.08 0.10 −0.21 −0.28 −0.27 −0.36 −0.59 −0.58 −0.73

2.63 2.47 2.08 1.18 1.16 0.97 0.82 0.82 0.89 0.49 0.33 −0.25

2.12 2.07 1.87 1.27 1.16 1.00 0.68 0.69 0.65 0.25 0.31 −0.28

2.63 2.93 1.78 1.05 1.08 0.58 0.61 0.41 0.48 0.12 0.03 −0.59

1.97 2.13 1.40 1.07 1.03 0.68 0.29 0.61 0.31 −0.09 −0.18 −0.27

1.31 1.61 1.21 0.87 0.75 0.97 0.42 0.99 0.53 0.26 0.24 0.48

All values are in electronvolts. The Gibbs free energy of NH3(g) is calculated to be −0.36 eV. Metals are listed in order of increasing binding strength of N.

a

including translational, vibrational, and freedom. We calculate the Gibbs free energy N2(g) and H2(g). For example, the free calculated from the reaction (1/2)N2 which leads to

of a specific reaction pathway by calculating the activity of the most energetically difficult reaction step within that pathway. The overall activity, A, is given by A = kBT mini(ln(ki/ko)) where kB is the Boltzmann constant, T the absolute temperature, ki the rate constant for elementary step i in a particular mechanism, and ko a normalization factor. The rate constant for elementary step i is given by the Arrhenius expression ki = vi exp(−Eia/kBT) where νi is the prefactor and Eia is the activation energy for elementary step i. For N−N bond formation steps, we use the activation energy calculated using the nudged elastic band method, while for proton-transfer reactions we use the free energy of reaction (when positive, or 0 when negative) as a first approximation for the activation energy. In this formulation of activity, activity is a negative number, and activity increases as the value approaches zero. For comparison purposes, at room temperature a change in activity by 0.06 eV corresponds to an order of magnitude change in the rate constant of the rate-determining step. Finally, we predict the operating potentials of these surfaces as a measure of their energy efficiency (minimization of overpotential) and for comparison with experimental results. First, we calculate the onset potential for the two reaction mechanisms previously described: (1) the Gerischer and Maueuer mechanism and (2) the N + N mechanism. We define the onset potential, for a particular pathway, as the required potential such that all proton−electron transfer reactions (i.e., the elementary steps which are influenced by cell potential) are exothermic (ΔG < 0). Because we have not evaluated N−N bond formation energies for all metals, we cannot predict the most kinetically favored Gerischer−Mauerer pathway on all metals. Therefore, when evaluating the operating potential for this mechanism, we choose the pathway that proceeds via hydrazine (H2NNH2) formation.

rotational degrees of of species relative to energy of ammonia is + (3/2)H2 → NH3,

ΔG NH3 = (TE NH3 − TSNH3 + ZPE NH3) −

1 3 (TE N2 − TSN2 + ZPE N2) − (TE H2 − TSH2 + ZPE H2) 2 2

where TE is the total energy of a species calculated from DFT, T the standard temperature (298 K), S the calculated entropy, and ZPE the zero-point energy for the species. The free energies of other species are calculated similarly. When calculating the free energy of adsorbed species, the total energy, entropy, and zero-point energy are all calculated for the adsorbed state. To correct the calculated Gibbs free energy for a specific applied electrochemical potential, we employ a procedure which has been implemented to study other electrode-mediated reactions, including oxygen evolution,47,48 oxygen reduction,49,50 and methanol oxidation.51,52 We choose the reversible hydrogen electrode as a reference, and we henceforth quote all potentials versus this reference. At standard conditions, hydrogen gas is in equilibrium with protons and electrons, at a defined potential of 0.00 V. When the pH differs from 0, we correct the free energy of H+ ions using the Nernst equation. Simplified, we have G(pH) = −kBT ln[H+], where kB is the Boltzmann constant and T is the absolute temperature. The reactions in this system involve proton−electron abstraction by hydroxyl ions in solution to form water, thereby having the net release of an electron. To account for the freeenergy change of this reaction, we use tabulated half-cell potentials. The reaction OH− + (1/2)H2 → H2O + e− has a potential 0.83 V versus the standard hydrogen electrode. To calculate the overall free energy of a proton−electron transfer reaction, following Hess’s Law, we add the DFT-calculated free energies of reaction for a dehydrogenation analogue, the reversible hydrogen electrode, and our half-cell reaction, corrected for pH 14. If the potential of the cell is increased by U, the free energy of electrons is shifted by −eU, where e is the absolute charge of an electron. To determine the most probable reaction mechanism and quantify the activity of each catalytic surface, we perform a simple Sabatier analysis.50 Specifically, we evaluate the activity

3. RESULTS AND DISCUSSION We studied the adsorption of all species within the reaction network on 12 metal surfaces and present the respective free energies in Table 1. Optimized structures for HxNNHy intermediates (NNH, HNNH, NNH 2 , HNNH 2 , and H2NNH2) in their preferred adsorbed states on Pt(111) are shown in Figure 1. Optimized structures on the other surfaces studied are provided in Supporting Information. 3.1. Electro-Oxidation on Pt(111). We begin our analysis by describing the behavior on Pt and then describing trends across all metals studied. First, we consider the N + N C

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Re

0.92 1.01 1.45 0.66 0.65 0.15 0.01 0.51 −2.63

All energies are given relative to infinitely separated adsorbates as initial states, 1/9 ML coverage. All values are in electronvolts. Metals are listed in order of increasing binding strength of N. a

−0.40 −0.41 0.18 −0.42 0.21 0.49 0.30 0.02 −0.33 −0.41 −0.10 0.39 −0.35 0.21 0.34 0.37 0.24 −0.49

Os Rh

−0.18 −0.03 0.25 −0.21 0.17 0.33 0.41 0.25 −0.89 −0.09 −0.22 0.39 −0.38 −0.20 0.09 0.41 0.12 −0.82

Co Ir

−0.10 −0.09 0.50 −0.13 0.32 0.39 0.21 0.14 −0.82 −0.03 −0.12 0.56 −0.29 −0.09 0.33 0.39 −0.03 −0.97

Ni Pt

0.28 0.17 0.52 0.28 0.05 0.14 0.08 0.00 −1.16 0.26 0.32 0.54 0.20 −0.02 0.20 0.13 −0.09 −1.18

Pd Cu Ag elementary step

Au

0.45 0.44 1.22 0.19 0.38 0.47 0.30 0.21 −2.08

Ru D

NH3 + OH− + * → NH2* + H2O + e− NH2* + OH− → NH* + H2O + e− NH* + OH− → N* + H2O + e− H2NNH2* + OH− → HNNH2* + H2O + e− HNNH2* + OH− → NNH2* + H2O + e− HNNH2* + OH− → HNNH* + H2O + e− NNH2* + OH− → NNH* + H2O + e− HNNH* + OH− → NNH* + H2O + e− NNH* + OH− → N2* + H2O + e−

Table 2. Free Energies of Deprotonation Steps for NHx and HxNNHy Intermediates at 0 V on Close-Packed Monometallic Surfacesa

mechanism. On Pt at 0 V, the free energy of reaction for proton transfer to hydroxyl (while releasing an electron) from NH3 to adsorbed NH2* is 0.28 eV, from NH2* to NH* is 0.17 eV, and from NH* to N* is 0.52 eV (see Table 2). The N* adatoms recombine exothermically (−1.58 eV), with a high activation energy barrier of 2.24 eV. A free-energy diagram illustrating the reaction coordinate is shown in Figure 2. Using the methods described previously, we calculate an onset potential of 0.52 V, which is the potential required to facilitate the most difficult proton−electron transfer in this pathway. In this case, this reaction step is NH oxidation to N, which is an important result to be revisited later. We call this elementary step the potential determining step for this mechanism. This onset potential and those of the other metals studied are plotted in Figure 3 for comparison. In the mechanism proposed by Gerischer and Mauerer, we must consider additional proton−electron transfer and N−N bond formation steps (see Tables 2 and 3). First, we analyze the dehydrogenation of the most hydrogenated HxNNHy species, hydrazine (H2NNH2), which is formed by dimerization of NH2. Hydrazine reacts with hydroxyl ions in solution to produce water and an HNNH2 species. The reaction energy on Pt is 0.28 eV, as shown in Table 2. HNNH2 can be oxidized to HNNH, NNH, and finally N2 (with production of water with each reaction step). The free energies of reaction are 0.14, 0.00, and −1.16 eV, repsectively. Alternatively, HNNH2 can be oxidized to NNH2 (instead of HNNH), which is further oxidized to NNH as before. The respective free energies of reaction are 0.05 and 0.08 eV. The activation energy for formation of these hydrogenated N2 species are shown in Table 3. Of these, the reaction NH* + NH* → HNNH* (Ea = 2.69 eV) is more activated than N* + N* → N2* (Ea = 2.24 eV). N* + NH* → NNH* (Ea = 2.26 eV) formation has a barrier similar to that of N2 formation. The formation of the other HxNNHy species have significantly reduced activation energy barriers as compared with that of N2 formation: N* + NH2* → NNH2* (Ea = 1.21 eV), NH* + NH2* → HNNH2* (Ea = 1.52 eV), and NH2* + NH2* → H2NNH2* (Ea = 1.07 eV). A freeenergy diagram illustrating these results (at 0 V) is shown in Figure 2. From the Sabatier analysis, the minimum-energy pathway for the Gerischer−Mauerer mechanism goes through deprotonation of NH3 to NH2, dimerization of NH2 species to hydrazine (H2NNH2), and then successive deprotonation to N2. Because the deprotonation steps are not rate-limiting, pathways which proceed through HNNH or NNH2 have the same activity via this measure. However, we note that because NNH2 is lower in energy than HNNH, we would expect the reaction to proceed through the NNH2 intermediate. The rate-determining step is N−N bond formation to make hydrazine (NH2* + NH2* →

1.24 1.04 1.06 0.52 0.80 −0.06 −0.47 0.40 −2.47

Figure 1. Optimized geometry of HxNNHy species in their minimumenergy structures on Pt(111): (a) NNH, (b) HNNH, (c) NNH2, (d) HNNH2, and (e) H2NNH2. Platinum atoms are shown in silver, nitrogen atoms in blue, and hydrogen atoms in white.

−0.55 −0.89 0.13 −0.75 −0.32 −0.01 0.34 0.04 0.25

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The Journal of Physical Chemistry C

N−N bond formation, the rate-determining step, is ∼20 orders of magnitude higher for the Gerischer−Mauerer mechanism) and a lower onset potential. Because the lower onset potential also corresponds to the mechanism with higher activity, we would expect N2 product formation to begin at this potential, 0.28 V. As the potential is raised to 0.52 V (the onset potential for the N + N mechanism), N adatoms begin to form. Because the removal of N adatoms is very sluggish, the catalyst essentially becomes poisoned. Therefore, there exists a narrow range of operating potentials for Pt(111): activation at 0.28 V and then deactivation above 0.52 V. Experimental results on Pt(111) have shown similar behavior, with onset of N2 production at 0.5 V and deactivation above 0.65 V26 (deactivation above 0.57 V20). The fact that we underpredict the onset for N2 formation by ∼0.2 V reflects the approximate nature of our analysis. In particular, all of our calculations were performed at 1/9 ML coverage, while the surface coverage found by Gootzen et al. was much higher (θ = 0.6). This may have a profound effect on the observed chemistry and the onset potential19 as high surface coverage has been shown to weaken adsorption of species.53,54 Additionally, our analysis neglects any solvation of the intermediates and does not account for the presence of electrolyte. The ammonia electro-oxidation mechanism on Pt(100), Pt(111), and Pt(110) was investigated using ab initio molecular dynamics simulations by Skachkov et al.55 In contrast to what we found on Pt(111), they find that, at moderate potentials (>0.5 V), the reaction proceeds through the N + N mechanism on Pt(100). Importantly, the binding of N adatoms to Pt(100) bridge sites and the geometry of the Pt(100) facet both facilitate N−N bond formation. However, at lower potentials, they suggest that the Gerischer−Mauerer and N + N mechanisms are competitive, though the activity will be much lower than at the higher potential range. Though their analysis of the mechanism on Pt(111) was not as extensive as that for Pt(100), they suggest that strong N binding to Pt(111) hollow sites leads to sluggish N−N bond formation and thereby N poisoning, which is in agreement with our results. 3.2. NH3 Electro-Oxidation on Other Close-Packed Transition-Metal Surfaces. 3.2.1. Metals Binding N More Weakly than Pt. Having established the most likely NH3 electro-oxidation mechanism for Pt(111), we now look at trends across metals. First, we consider metals which bind N more weakly than Pt. This includes the coinage metals Cu, Au, and Ag, which bind much weaker than Pt, and Pd, which is similar to Pt with respect to binding N. First, we will comment on the coinage metals. The coinage metals bind species much more weakly than Pt and can easily form N−N bonds between surface species. On Cu, we rank the difficulty of the N−N bond formation barriers leading to the respective products listed in ascending order as N* + NH2* → NNH2* (1.07 eV) < N* + N* → N2* (1.16 eV) < N* + NH* → NNH (1.25 eV) < NH* + NH2* → HNNH2* (1.49 eV) < NH* + NH* → HNNH* (1.66 eV) < NH2* + NH2* → H2NNH2 (1.67 eV). From this, we see that the N−N bond formation barrier is only slightly reduced (i.e., the lowest barrier is for N* + NH2* → NNH2, 1.07 eV) in the Gerischer−Mauerer mechanism when compared to the N + N mechanism, which has a barrier of 1.16 eV. On the other hand, dehydrogenation of NHx or HxNNHy species is quite difficult. A potential energy surface illustrating these differences on Cu, at 0 V, is shown in Figure 4.

Figure 2. Free-energy diagram for ammonia electro-oxidation on Pt(111) at 0 VRHE. In purple are N−N bond formation reaction steps with the respective transition-state energies. Stoichiometry is balanced with OH−, H2O, H+, and e−, which are not shown explicitly. All energies are given with adsorbates at infinite separation, 1/9 ML coverage. Zero energy corresponds to N2(g).

Figure 3. Estimated onset potential for close-packed facets of transition metals. For each metal, the left bar corresponds to the onset potential for the N + N mechanism, while the right bar corresponds to the onset potential for the Gerischer−Mauerer mechanism (proceeding through hydrazine formation). The potential determining steps for these mechanisms are color-coded. Absolute N* binding energy increases to the right. In the legend, and for the A + OH− → B + H2O + e− generic reaction step, the species on the left of the arrow is A, right of the arrow is B.

H2NNH2*), with a calculated activity of −1.07 eV. We note that the previously described N + N mechanism is rate-limited by N−N bond formation to make N2 (N* + N* → N2*) and has a calculated activity of −2.24 eV (recall 0.06 eV corresponds to 1 order of magnitude change in the rate constant at room temperature). The onset potential for the Gerischer−Mauerer mechanism is calculated to be 0.28 V. This is due to two equivalently difficult potential determining steps: NH3 + OH− + * → NH2* + H2O + e− and H2NNH2* + OH− → HNNH2* + H2O + e−. For comparison, the N + N mechanism has an onset potential of 0.52 V, with NH* + OH− → N* + H2O + e− as the potential determining step. Comparing these two pathways, we see that the pathway proposed by Gerischer and Mauerer has a much higher predicted activity (at room temperature, the rate constant for E

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The Journal of Physical Chemistry C Table 3. Free Energies of Reaction and Activation Energy Barrier for N−N Bond Formation Reaction Stepsa N* + N* → N2 Ag Au Cu Pd Pt Ni Ir Co Rh Ru Os Re

N* + NH* → N2H*

N* + NH2* → N2H2*

NH* + NH* → HNNH*

NH* + NH2* → N2H3*

NH2* + NH2* → N2H4*

ΔGrxn

Ea

ΔGrxn

Ea

ΔGrxn

Ea

ΔGrxn

Ea

ΔGrxn

Ea

ΔGrxn

Ea

−6.41 −6.33 −3.88 −1.89 −1.58 −0.46 −0.26 0.23 0.28 0.60 1.64 2.98

− − 1.16 2.16 2.24 − 2.10 2.00 2.55 − − −

−2.32 −2.8 −0.57 −0.17 0.10 1.08 1.07 1.43 1.42 1.48 2.14 2.86

− − 1.25 2.07 2.26 − 2.31 2.05 2.32 − − −

−1.32 −1.29 −0.43 0.02 0.19 0.56 0.77 0.80 0.98 1.01 1.43 1.64

− − 1.07 1.59 1.21 − 1.29 1.67 1.73 − − −

−1.38 −2.14 0.44 0.47 0.62 1.67 1.43 1.69 1.42 1.63 2.30 2.96

− − 1.66 2.26 2.69 − 2.68 2.27 2.38 − − −

−0.52 −1.04 0.41 0.59 0.66 1.22 0.95 1.38 1.06 1.19 1.40 2.08

− − 1.49 1.61 1.52 − 1.51 2.03 1.71 − − −

−0.16 −0.52 0.66 0.71 0.55 1.39 0.99 1.54 1.25 1.44 1.40 1.95

− − 1.67 1.57 1.07 − 1.21 2.00 1.62 − − −

a Activation energies are calculated only for Cu, Pd, Pt, Ir, Co, and Rh. All energies are given relative to infinitely separated adsorbates as initial states, 1/9 ML coverage. All values are in electronvolts. Metals are listed in order of increasing binding strength of N.

with a maximum activity of −1.16 eV. Overall, the Gerischer− Mauerer mechanism is predicted to have activity higher than that of the N + N mechanism, and the maximum activity of −1.07 eV is equivalent to that of Pt. However, if we now consider the onset potentials for these two pathways, we see an important result. The onset potential for the N + N mechanism is 1.22 V, far outside the range of feasibility. For the Gerischer−Mauerer pathway, the onset potential is 0.46 V. However, this latter value corresponds to a pathway proceeding through hydrazine formation, which is a kinetically difficult reaction. The minimum-energy pathway proceeds through NNH2 formation, requiring the presence of N adatoms. Therefore, the onset potential for the minimumenergy pathway is 1.22 V. This conclusion is in agreement with experimental findings suggesting that Cu is not active toward ammonia electro-oxidation unless oxidized, and the catalyst undergoes electrodissolution by formation of Cu(NH3)2+ complexes.20 We caution that our calculations were performed on the close-packed facet of Cu (and other metals), whereas these experiments were conducted on flag electrodes, which may have multiple crystal facets exposed. As previous work has demonstrated the structure sensitivity of ammonia electrooxidation on Pt,26 similar caution is advised when directly comparing against experiments not performed on single crystals of other metals. Although we have not evaluated the detailed reaction kinetics of N−N bond-forming steps on Ag and Au, we can estimate properties based on reaction thermodynamics (see Supporting Information for free-energy surfaces). By calculating onset potentials for these metals, we see that the reaction proceeds at unfeasible potentials along both mechanisms (Figure 3). Experiments on flag electrodes have shown that both Ag and Au are inactive for NH3 electro-oxidation, instead dissolving similarly to Cu.20 Palladium is an interesting metal in that its binding characteristics are similar to those of Pt, yet experimentally Pd is inactive for NH3 electro-oxidation.20 It has been postulated that the low activity for Pd is due to formation of N* at lower potentials than Pt, thereby poisoning the surface.20 However, the results of our calculations do not support this argument. The free energy of adsorbed N is actually 0.15 eV more unstable on Pd when compared to Pt. Furthermore, when we calculate the onset potential for the formation of N adatoms

Figure 4. Free-energy diagram for ammonia electro-oxidation on Cu(111) at 0 VRHE. In purple are N−N bond formation steps with respective transition-state energies. Stoichiometry is balanced with OH−, H2O, H+, and e−, which are not shown explicitly. All energies are given with adsorbates at infinite separation, 1/9 ML coverage.

First, considering the N + N mechanism, we see that N−N bond formation (Ea = 1.16 eV) is not limiting the reaction rate at 0 V. In fact, the most difficult reaction step at 0 V is NH* + OH− → N* + H2O + e− (ΔG = 1.22 eV, see Table 2). In contrast, the minimum-energy pathway for the Gerischer− Mauerer mechanism features N−N bond formation between N and NH2 with a barrier of 1.07 eV. Still, the most energetically difficult step is dehydrogenation of NH to N (ΔG = 1.22 eV at 0 V). Because both pathways are limited by NH* + OH− → N* + H2O + e−, their calculated activities at 0 V are both −1.22 eV. This corresponds to a few orders of magnitude lower in the rate constant versus Pt(111), which had a calculated activity of −1.07 eV. For the Gerischer−Mauerer mechanism, if the potential were increased to 0.15 V, the reaction energy for NH* + OH− → N* + H2O + e− would decrease by 0.15 eV from 1.22 to 1.07 eV, making it of equal difficulty as N−N bond formation (N* + NH2* → NNH2*). If the potential were increased beyond 0.15 V, then the rate-determining step would change to N−N bond formation (N* + NH2* → NNH2*) and the activity would plateau at −1.07 eV. In contrast, the rate-determining step of the N + N mechanism would change at potentials greater than 0.06 V. At potentials more positive than 0.06 V, the activity would be limited by N−N bond formation (N* + N* → N2*), F

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Gerischer−Mauerer mechanism may be more active via formation of NNH2, but the activity is significantly lower that that of Pt (−1.67 eV for Co versus −1.07 eV for Pt). Furthermore, in order to form NNH2 directly from N and NH2 on Co, NH must first be dehydrogenated to N, which requires a high potential of 0.39 V. For Ni, we can make some estimates of the activity based on the thermochemistry of the elementary steps. In particular, we see that H2NNH2 binds very weakly on Ni when compared to its rather strong binding of N (this is similar to the behavior of Co, see Table 1). This leads to a very endothermic hydrazine formation energy (ΔG = 1.39 eV) from adsorbed NH2 species. The corresponding energy barrier (small contribution from zero-point energy and entropy) must be at least this value, which is already greater than the barrier for hydrazine formation on Pt (Ea = 1.07 eV), suggesting that the activity of Ni would be lower than that of Pt. Therefore, on the basis of thermochemistry, we can estimate that Ni will be less active than Pt. Yet, this argument assumes that the minimum-energy pathway on Ni would proceed through the hydrazine intermediate, which may not be true if all of the activation energy barriers were calculated. Experimental studies on Ni electrodes have shown that Ni can be active for ammonia electro-oxidation;56 however, under alkaline conditions, it can become oxidized and deactivate.36,57 Rh, Ru, Os, and Re bind N significantly stronger than Pt. Furthermore, the thermochemistry shows that for Rh, Os and Re, the onset potential for the Gerischer−Mauerer mechanism is higher than that of the N + N mechanism. The onset potential for the Gerischer−Mauerer mechanism (N + N mechanism onset potential in parentheses) is 0.33 V (0.25 V) for Rh, 0.30 V (0.18 V) for Os, and 0.25 V (0.13 V) for Re. For Ru, the onset potential for the Gerischer−Mauerer (0.33 V) is lower than the N + N mechanism (0.39 V), but only by 0.06 V. As a consequence of this, the Gerischer−Mauerer mechanism will be active only at potentials above the point at which N* forms (for Ru they occur at a similar potential). Therefore, these catalysts will be poisoned by N* at all operational potentials. Experimental results agree with these findings, showing no steady-state production of N2 on Rh and Ru, but there is a small amount of N2 formed between 0.3 and 0.8 V on Rh and between 0.5 and 0.8 V on Ru during cyclic voltammetry.20 Iridium is another metal with significantly stronger N binding than Pt (0.66 eV stronger). Yet, experimentally, it is active for N2 formation.20,31,32 From the data presented in Table 1, we see that the free energies of adsorbed species on Ir are similar to those of species on Rh. One important difference, however, is that Ir binds N 0.27 eV weaker than Rh. Therefore, the onset potential for electro-oxidation of NH to N (0.50 V) is higher than the onset potential for the Gerischer−Mauerer mechanism (0.32 V), whereas the case is the opposite for Rh. From calculated onset potentials, Ir should be active between 0.32 and 0.50 V. Above 0.5 V, deactivation will occur, in close agreement with experimental findings.20 Previously, we saw that Pd (also Ni and Co) have favorable onset potentials, yet the formation of hydrazine is calculated to be kinetically difficult. On Ir, the hydrazine formation barrier is modest, only 1.21 eV as compared to 1.07 eV on Pt. Therefore, we predict that Ir should be active, through the Gerischer−Mauerer mechanism, proceeding through a hydrazine intermediate. 3.3. Energy Efficiency and Activity Trends. For a catalyst to be active for NH3 electro-oxidation, it is essential

(the onset potential for the N + N mechanism), we obtain 0.54 V, which is close to the value on Pt (0.52 V). Most importantly, this potential is significantly higher than the calculated onset potential for the Gerischer−Mauerer mechanism, which is 0.26 V. This allows for oxidation to occur at a potential lower than N* formation, preventing poisoning. With these results, why might Pd be inactive while Pt is active? The key difference between Pt and Pd appears to be the difficulty in forming hydrazine, the rate-determining step of the minimum-energy pathway (see Tables 2 and 3; the free-energy surface is provided in Supporting Information). The hydrazine formation barrier is 1.57 eV for Pd, but only 1.07 eV for Pt. This difference is due to the reaction coordinate whereby NH2* species, which are adsorbed on bridge sites, must first diffuse to top sites prior to forming the N−N bond in H2NNH2. Binding on top sites of Pd are 0.58 eV less stable than on bridge sites, whereas on Pt top sites are only 0.40 eV less stable than on bridge sites. This additional diffusion energy cost that is included in the activation barrier (barriers are referenced to infinite-separation best adsorbed states) is quite significant. In reality, we would expect a significantly lower coverage of topsite bound NH2 species, which are required for N−N bond formation in hydrazine. Overall, Pd features onset potentials similar to those of Pt, but the activity at 0 V is −1.57 eV, as compared to −1.07 eV for Pt, which would translate to ∼8 orders of magnitude difference in reaction rate constant between the two surfaces at room temperature. To summarize these results, we conclude that metals which bind N more weakly than Pt are inactive. For Cu, Ag, and Au, the onset potential is too high for realistic operation, and experiments show enhanced electrodissolution of the metal in the presence of ammonia.20 For Pd, which has similar (yet slightly weaker) binding, the hydrazine formation barrier is high because of unfavorable binding of NH2, while N poisoning is likely not a problem. 3.2.2. Metals Binding N More Strongly than Pt. Next, we examine the energetics of NH3 electro-oxidation on metal surfaces binding N more strongly than Pt. Those studied include (in increasing absolute N binding strength) Ni, Ir, Co, Rh, Ru, Os, and Re. First we consider the 3d metals, Ni and Co; then discuss Rh, Ru, Os, and Re; and finally conclude with describing our results on Ir. From calculated thermochemistry (shown in Table 1), we can evaluate the onset potential on Ni and Co. For Ni we calculate 0.56 and 0.33 V for the N + N mechanism and the Gerischer−Mauerer mechanism, respectively. For Co, we calculate 0.39 and 0.12 V, respectively. Importantly, we see that the onset potential for the Gerischer−Mauerer mechanism is lower than that of the N + N mechanism, meaning that the surface should not be poisoned by N* at lower operational potentials. On the basis of these measures, it seems as though Ni and Co may be good NH3 electro-oxidation catalysts. Co, in particular, has the lowest calculated onset potential of all the metal surfaces that we investigated. Therefore, we evaluated the activation energy barriers for N−N bond formation on Co. Evaluation of the N−N bond formation barriers on Co shows that, though the onset potential may be low, sluggish N− N formation kinetics will likely limit the reaction rate. The N− N bond formation barriers are calculated as N* + NH2* → NNH2* (1.67 eV) < N* + N* → N2* (2.00 eV) = NH2* + NH2* → H2NNH2* (2.00 eV) < NH* + NH2* → HNNH2* (2.03 eV) < N* + NH* → NNH* (2.05 eV) < NH* + NH* → HNNH* (2.27 eV). From these barriers we see that the G

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The Journal of Physical Chemistry C that N2 is formed at a lower potential than N*. If N* adatoms form, it is likely that the catalyst will become poisoned. Therefore, if the onset potential for the Gerischer−Mauerer mechanism is lower than the onset potential for the N + N mechanism, N2 formation can occur without N* poisoning. From the data shown in Figure 3, we see that Re, Os, and Rh will be poisoned at potentials required for N2 formation. For Ru, there is only a small difference between the two mechanisms’ onset potentials, and for Au, the two onset potentials are the same. Therefore, both Ru and Au will likely be poisoned at operating potentials. From the metals studied here, this leaves Ag, Cu, Pd, Pt, Ni, Ir, and Co as viable catalysts, based on this measure. However, the required onset potential for Ag (0.92 V) is unrealistically high for stable operation. Of these remaining catalysts, we can now estimate their energy efficiency. Because the reaction occurs at the anode, the lowest onset potential (and therefore lowest overpotential) will preserve the maximum amount of chemical energy. Because the Gerischer−Mauerer mechanism appears to be the operable mechanism, we rank based on that mechanism’s onset potential (see Figure 3). We obtain the following ranking (from best to worst): Co < Pd < Pt < Ir < Ni < Cu. However, as discussed earlier, although Co has a low onset potential, the calculated N−N bond formation barriers will likely limit Co’s activity. From the Sabatier analysis, we have evaluated a number of surfaces based on calculated activation energies (Table 3) and thermochemistry (Tables 2 and 3). From this, we generate a volcano plot (Figure 5) of each surface’s activity at 0 V against the binding energy of N. We only include surfaces for which we have calculated N−N bond formation barriers. Then, we rank the activity of the catalysts as Pt > Ir > Cu > Pd > Rh > Co. Experiments show Pt is the most active catalyst, followed by Ir, with Co, Pd, and Rh being inactive.20,32

From Figure 5, it is clear that the activity of the Gerischer− Mauerer mechanism is many orders of magnitude higher than that of the N + N mechanism for metals with stronger binding of N than Cu (to the left of Cu on the plot). The activity in this regime is limited by N−N bond formation barriers, and the barriers are, generally, reduced in the Gerischer−Mauerer mechanism. For Cu, we see that both mechanisms have the same activity because the activity is limited by deprotonation rather than N−N bond formation. Similar behavior would be expected of metals which bind more weakly than Cu. It is important to note that the activity predicted for Cu (−1.22 eV) is similar to that of Ir (−1.21 eV), yet experimentally Cu has been shown to be inactive. It is likely that this discrepancy is due to directly comparing activation energies of N−N bond formation steps with Gibbs free energies of reaction of proton− electron transfer steps in our analysis. It is likely that there is some additional kinetic barrier beyond the thermochemical “barrier.” However, still the analysis illustrates the importance for balancing the binding properties of a surface in order to maximize catalytic activity. 3.4. Catalyst Design Principles. Though experimental studies have demonstrated the structure sensitivity of ammonia electro-oxidation on Pt, with Pt(100) more active than Pt(111), it would be beneficial to identify catalysts which are active on their close-packed facets because the close-packed facet is more thermodynamically stable. According to the Sabatier principle, the development of improved catalysts requires a balance between the ability to activate bonds and to make bonds. From the predicted activity data shown in Figure 5, we suggest that Pt is the best monometallic close-packed surface with Cu and Ir being the next best. The activity for Pt is limited by N−N bond formation rates (Pt binds N too strongly), whereas on Cu, deprotonation is rate-limiting (Cu binds N too weakly). Therefore, the peak of a hypothetical volcano should lie between Pt and Cu, where the rate-limiting step switches between these two kinds of steps. Note, however, that this intersection point will vary with the potential of the electrode because the activity of the deprotonation steps is a function of the applied potential. As the potential increases, the peak of the volcano will shift to the right (weaker binding of N). With this in mind, in searching for improved alloy catalysts, one should try to find surfaces with binding properties between those of Pt and Cu.

4. CONCLUSION We have investigated the mechanism of NH3 electro-oxidation on model close-packed surfaces of Au, Ag, Cu, Pd, Pt, Ni, Ir, Co, Rh, Ru, Os, and Re. Our results support the dominance of the reaction mechanism proposed by Gerischer and Mauerer, that is, a mechanism whereby the N−N bond formation occurs via hydrogenated NHx species rather than via N adatoms. In particular for Pt and Ir, the minimum-energy pathway involves dimerization of adsorbed NH2 species into hydrazine, which is subsequently deprotonated to form N2. We note that the dominant mechanism may be different on other crystal facets, which may account for the demonstrated structure sensitivity on Pt. We have shown that there exists a specific range of operating potentials for each metal in which N2 is formed without poisoning the surface with N adatoms, and the predicted ranges are in reasonable agreement with experimental evidence. Surfaces where the difference in the onset potential of the two mechanisms is very small, here Ag, Rh, Os, and Re (within error also Ru), are poisoned by strongly adsorbed N

Figure 5. Activity as predicted by Sabatier analysis for both mechanisms at 0 VRHE. Less negative activity is better. The activity is calculated from the thermochemistry of electrochemical steps and the activation energy of N−N bond-forming steps and is a measure of the maximum resistance of the minimum-energy pathway for each mechanism. For Rh, Co, Ir, Pt, and Pd, the activity is limited by (ratedetermining step) N−N bond formation, while for Cu the activity is limited by proton−electron transfer. H

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adatoms, prior to forming N2, and therefore are inactive. On the basis of a simple Sabatier analysis, we predict that Pt is the most active metal followed by Ir and Cu, while all the other metals studied here have significantly lower activity. We suggest that improved alloy catalysts should bind N stronger than Cu but weaker than Pt.

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ASSOCIATED CONTENT

S Supporting Information *

Optimized geometries for adsorbed HxNNHy species on Re(0001), Ru(0001), Os(0001), Co(0001), Rh(111), Ir(111), Ni(111), Pd(111), Cu(111), Ag(111), and Au(111); calculated free-energy diagrams for ammonia electro-oxidation on Re(0001), Ru(0001), Os(0001) Co(0001), Rh(111), Ir(111), Ni(111), Pd(111), Ag(111), and Au(111). This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Phone: 608-262-9053. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This paper is dedicated to Prof. Steven Sibener. Work at UWMadison was supported by DOE-BES, Office of Chemical Sciences. J.A.H. thanks Air Products & Chemicals, Inc. for partial support through a graduate fellowship. We thank Luke T. Roling for his comments on the manuscript. Computational work was performed in part using supercomputing resources at the following institutions: EMSL, a national scientific user facility at Pacific Northwest National Laboratory (PNNL); the Center for Nanoscale Materials at Argonne National Laboratory (ANL); and the National Energy Research Scientific Computing Center (NERSC). EMSL is sponsored by the Department of Energy’s Office of Biological and Environmental Research located at PNNL. CNM and NERSC are supported by the U.S. Department of Energy, Office of Science, under contracts DE-AC02-06CH11357 and DE-AC0205CH11231, respectively.

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DOI: 10.1021/jp512981f J. Phys. Chem. C XXXX, XXX, XXX−XXX