Evaluating the Thermodynamics of Electrocatalytic N2 Reduction in

Sep 9, 2016 - ABSTRACT: The development of a sustainable ammonia synthesis by proton-coupled electroreduction of dinitrogen (N2) requires knowl-...
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Evaluating the Thermodynamics of Electrocatalytic N2 Reduction in Acetonitrile Brian M. Lindley,† Aaron M. Appel,‡ Karsten Krogh-Jespersen,§ James M. Mayer,◊ and Alexander J. M. Miller*,† †

Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-3290, United States Catalysis Science Group, Pacific Northwest National Laboratory, Richland, Washington 99352, United States § Department of Chemistry and Chemical Biology, Rutgers, The State University of New Jersey, New Brunswick, New Jersey 08903, United States ◊ Department of Chemistry, Yale University, New Haven, Connecticut 06520-8107, United States ‡

S Supporting Information *

ABSTRACT: The development of a sustainable ammonia synthesis by proton-coupled electroreduction of dinitrogen (N2) requires knowledge of the thermodynamics described by standard reduction potentials. The first collection of N2 reduction standard potentials in an organic solvent are reported here. The potentials for reduction of N2 to ammonia (NH3), hydrazine (N2H4), and diazene (N2H2) in acetonitrile (MeCN) solution are derived using thermochemical cycles. Ammonia is thermodynamically favored, with a 0.43 V difference between NH3 and N2H4 and a 1.26 V difference between NH3 and N2H2. The thermodynamics for reduction of N2 to the protonated products ammonium (NH4+) and hydrazinium (N2H5+) under acidic conditions are also presented. Comparison with the H+/H2 potential in MeCN reveals a 63 mV thermodynamic preference for N2 reduction to NH3 over H2 production. Combined with knowledge of the kinetics of electrode-catalyzed H2 evolution, a wide working region is identified to guide future electrocatalytic studies. Schrock and Yandulov,9 Nishibayashi and co-workers,10−12 and Peters and co-workers13,14 has produced a series of low-valent Mo and Fe complexes capable of catalytic ammonia production at low temperature. Catalysis is driven by (sometimes incompatible)15 strong chemical reductants and organic acids, which may contribute to relatively low turnover numbers. The energy efficiency of these systems has only been estimated computationally;16 experimental data on N 2 reduction thermodynamics in organic solvents is lacking.

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he Haber−Bosch process utilizes heterogeneous transition-metal catalysts to convert dinitrogen and dihydrogen to ammonia (eq 1).1,2 This century-old process remains the only industrially relevant method for producing nitrogen fertilizers, directly supporting half of all global food production.3 Ammonia has also received increasing attention in the energy sector as a potential nitrogen-based fuel component.4 The energetic and environmental cost of worldscale ammonia synthesis is as staggering as it is indispensable: the combination of H2 synthesis (usually by steam reforming of methane) and N2 hydrogenation (above 300 °C and 125 atm)1 is estimated to consume 1−2% of all global energy resources.5 A low-temperature, sustainably powered ammonia synthesis could reduce fossil energy dependence while meeting the increased demand for fertilizers needed to support a growing population. N2 + 3H 2 ⇌ 2NH3

N2 + 6H+ + 6e− ⇌ 2NH3

Long-term implementation of proton-coupled electron transfer (PCET) reduction of N2 will require replacement of stoichiometric chemical reductants with an applied electrochemical potential. Examples of electrochemical ammonia synthesis are currently limited to a handful of heterogeneous electrode materials capable of N2 reduction.17−19 Homogeneous examples feature only stoichiometric transformations,20,21 including a very recent report wherein Peters and

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The replacement of H2 with proton and electron sources has attracted considerable attention as an alternative route to ammonia, inspired by the tandem protonation and reduction of N2 (eq 2) by bacterial nitrogenase enzymes.6 Building on the pioneering studies of Chatt et al.,7,8 a recent surge of progress in homogeneous transition-metal-catalyzed N2 reduction led by © XXXX American Chemical Society

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Received: August 2, 2016 Accepted: August 31, 2016

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DOI: 10.1021/acsenergylett.6b00319 ACS Energy Lett. 2016, 1, 698−704

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ACS Energy Letters Scheme 1

N2(g ) + 8H+ + 6e− ⇌ 2NH4 +(aq)

co-workers noted the formation of ammonia in electrochemical reductions of an iron complex.22 The prospects of electrochemical N2 reduction depend upon multiple parameters, including (i) catalytic rate, (ii) selectivity for NH3 over competing products such as H2, and (iii) electrocatalytic overpotential. This last parameter is governed by the difference between the operating potential required to achieve catalysis at a certain rate and the thermodynamic potential of N2 reduction under the same conditions.23,24 The standard potentials at which N2 is reduced to various species will speak to the feasibility of selective NH3 production, as well as provide reference points for defining the overpotential of successful catalysts. Though the standard potentials of N2 reduction are well-known in aqueous solution, no such data is available for organic solvents. This report is the first for the standard N2 reduction potentials in an organic solvent. The thermodynamics of electrochemical N2 reduction are reported here in acetonitrile (MeCN) solvent. MeCN was chosen for this study because a wealth of thermochemical data is available and MeCN is widely used for electroreduction reactions.23,25 The new standard potentials can be used to benchmark N2 reduction electrocatalysts, adding to a growing collection of thermochemical data for the electroreduction of small molecules (H+, O2, and CO2) in organic solvents.26−30 Using thermochemical cycles starting from the known aqueous reduction potentials, the free energy of N2 reduction to N2H2, N2H4, and NH3 are derived. The standard potentials, and their dependence on the pKa of the proton source, show that N2 electroreduction is thermodynamically accessible. The thermodynamic selectivity for NH3 formation is established based on the relative potentials for N2 reduction to form NH3 versus the partially reduced intermediates N2H2 and N2H4. Similarly, a comparison of the N2/NH3 and H+/H2 potentials in MeCN reveals a thermodynamic preference for N2 reduction over proton reduction. The approach to derive standard potentials for the reduction of N2 in MeCN involves starting from the known aqueous N2 reduction potentials and applying appropriate solvation corrections. The values for the aqueous potentials used for the thermochemical cycles are given in eqs 3−5.31,32 Standard reduction potentials are given at standard states of 1 atm gases, 1 M solutes, and 298 K. Because no electrode can rapidly and reversibly interconvert N2 and NH3, the literature potentials are based on data for free energies of formation. Because of several literature inconsistencies (opposite signs or different values of E°), we have independently evaluated the aqueous standard potentials using available thermochemical data (Supporting Information, section S1).

E°(aq) = + 0.275 V vs SHE (3)

+

−

+

N2(g ) + 5H + 4e ⇌ N2H5 (aq)

E°(aq) = − 0.23 V vs SHE (4)

N2(g ) + 2H+ + 2e− ⇌ N2H 2(g )

E°(aq) = − 1.20 V vs SHE (5)

Conversions from aqueous to acetonitrile reduction potentials can be accomplished using the H+/H2 standard reduction potential in both H2O and MeCN. The reduction potential of the H+/H2 couple in water is simply the standard hydrogen electrode (SHE), defined at a potential of 0 V.33,34 The H+/H2 reduction potential in MeCN was recently measured experimentally using open-circuit potentiometry35 and is given relative to the IUPAC-recommended ferrocenium/ ferrocene reference potential,36,37 E°(H+/H2) = −0.028 V vs Fc+/0 in MeCN, in good agreement with a previous report by Kolthoff and Thomas.38 Derived here are standard potentials for reduction of N2 both to neutral products such as NH3(g) and to protonated derivatives such as solvated ammonium, NH4+(MeCN). For the protonated species, two approaches to determining intersolvent transfer free energies were considered. First, the pKa values of NH4+ and N2H5+, which are known in both H2O and MeCN,39 can be used along with the solvation free energies associated with dissolution of the gas-phase conjugate bases in both solvents. In cases where no experimental solvation energy values were available, computational methods were used to estimate the values (Supporting Information, sections S3 and S4). Alternatively, the free energy of cation transfer from water to acetonitrile could be utilized. This method has not been included because it requires the free energy of proton transfer from water to acetonitrile, the value of which is uncertain (Supporting Information, section S5).40 The following discussion derives the N2/NH3, N2/NH4+, N2/N2H4, N2/ N2H5+, and N2/N2H2 standard potentials in MeCN using the available thermochemical data. The uncertainty is estimated to be approximately ±30 mV in most cases, driven in large part by the uncertainty in the free energies of solvation (Supporting Information, section S9). The reduction of N2 to NH3(g) was considered first because ammonia is a gas at ambient temperature and represents the simplest thermochemical case. Scheme 1 shows the stepwise thermochemical cycle used to determine the N2(g)/NH3(g) potential. The aqueous standard potential for N2(g)/NH4+(aq) (eq 3) is modified by the H+/H2 potential in H2O (eq 6) and the H+/H2 potential in MeCN (eq 7), accounting for the change in proton solvation and the change in electrochemical reference. To account for formation of NH3(g), the pKa of 699

DOI: 10.1021/acsenergylett.6b00319 ACS Energy Lett. 2016, 1, 698−704

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ACS Energy Letters Scheme 2

Scheme 3

NH4+ in H2O (9.25)39 and the aqueous solvation energy of NH3 (ΔG° = −2.4 kcal·mol−1)41 are used (eqs 8 and 9, respectively). In experimental practice, electrocatalytic reductions are carried out with a range of organic acids of varying acidity. As the choice of acid will influence the overall reaction thermodynamics, the acidity is taken into account by eq 10. This affords a N2(g)/NH3(g) standard reduction potential in MeCN of +0.029 V vs Fc+/0, with a cathodic shift of 59 mV per pKa unit of the acid (eq 12), required for a process involving transfer of equal numbers of proton and electron equivalents. With the standard reduction potential for N2(g)/NH3(g) in hand, the corresponding potential for MeCN-solvated ammonia can be derived (Scheme 2). Starting with the free energy for the N2(g)/NH3(g) process (eq 11), only the solvation energy of NH3 in MeCN is needed (eq 13). This value was determined experimentally to be −0.41 kcal·mol−1 based on ammonia solubility measurements (Supporting Information, section S2). The result is a standard N2(g)/ NH3(MeCN) potential of +0.035 V vs Fc+/0, shifted anodically by only 6 mV compared to the N2(g)/NH3(g) potential (eq 15). This highlights the relatively small contribution of solvation to the N2(g)/NH3(MeCN) potential. In contrast to the reductions of O2, CO2, and H+, the products of N2 reduction are moderately basic. This means that the thermodynamic products will depend on the acidity of the medium, which affects the thermochemical analysis. The conjugate acid of ammonia, NH4+, has a pKa value of 16.5 in MeCN.39 If an acid with pKa < 16 is used, the N2(g)/ NH4+(MeCN) potential is more appropriate than the N2(g)/ NH3(MeCN) potential. Scheme 3 illustrates the cycle for determining the N2(g)/NH4+(MeCN) potential, in which the expression for the N2(g)/NH3(MeCN) potential (Scheme 3) is combined with pKa values of NH4+ and the added acid in MeCN (eqs 16 and 17). The N2(g)/NH4+(MeCN) standard reduction potential is +0.361 V vs Fc+/0 (eq 19). The potential shifts negatively by 79 mV per pKa unit of the acid, reflecting the unequal 8H+/6e− stoichiometry. Partial reduction of N2 to N2H4 has been observed in stoichiometric protonation−reduction schemes involving lowvalent transition-metal complexes.7,8,42−44 Hydrazine is also likely present as a cofactor-bound intermediate in biological nitrogen fixation by nitrogenase enzymes.6 Information on the N2/N2H4 potential will elucidate differences in driving force relative to NH3 and indicate applied potentials at which intermediate hydrazine formation will not be competitive.

The thermochemical cycle used for determining the N2(g)/ N 2 H 4 (MeCN) potential is analogous to the N 2 (g)/ NH3(MeCN) potential (Scheme S2). This cycle comprises the H+/H2 potentials in H2O and MeCN, the expression for the acidity of the proton source, the pKa of N2H5+ in H2O (7.96),39 and the solvation energies of gaseous N2H4 in water (−3.9 kcal· mol−1) and MeCN (−1.84 kcal·mol−1). The solvation energies for N2H4 were determined computationally (Supporting Information, sections S3 and S4). The N2(g)/N2H4(MeCN) standard reduction potential is −0.398 V, with a 59 mV cathodic shift per acid pKa unit (eq 21). N2(g ) + 4HBase+(MeCN ) + 4e− ⇌ N2H4(g ) + 4Base ΔG° = 36.73 + (4·1.364·pK a) kcal ·mol−1

(20)

E°(N2(g )/N2H4(MeCN )) = −0.398 V − 0.059· pK a (21)

Hydrazine and ammonia are bases of similar strength, and an analogous approach can be used to derive the N2(g)/ N2H5+(MeCN) standard potential as for NH4+ (Scheme S3). Necessary adjustments to eq 20 include the addition of the N2H5+ pKa in MeCN (16.6).39 The result is a N2(g)/ N2H5+(MeCN) standard potential of −0.153 V vs Fc+/0 (eq 23). The 5H+/4e− stoichiometry gives rise to a cathodic shift of 74 mV per pKa unit. N2(g ) + 5HBase+(MeCN ) + 4e− ⇌ N2H5+(MeCN ) + 5Base

ΔG° = 14.09 + (5· 1.364· pK a) kcal ·mol−1 (22) +

E°(N2(g )/N2H5 (MeCN )) = −0.153 V − 0.074· pK a (23)

Diazene is a viable substrate for conversion to ammonia by nitrogenase enzymes and is invoked as a metal-bound intermediate in various biological and chemical N2 reduction schemes.6,43 The thermochemical cycle for the reduction of N2 to N2H2 is analogous to the N2(g)/NH3(MeCN) and N2(g)/ N2H4(MeCN) cycles, except that the aqueous N2(g)/N2H2(g) potential is the starting point, so the pKa of N2H3+ is unnecessary (Scheme S4). No solubility data was found for N2H2, likely a consequence of its rapid disproportionation to N2 and N2H4 in solution.45 This problem was circumvented by computing the solvation free energy of N2H2 in MeCN (−0.03 kcal·mol−1). The resulting N2(g)/N2H2(MeCN) potential is −1.22 V vs Fc+/0, with a cathodic shift of 59 mV per pKa unit (eq 25). 700

DOI: 10.1021/acsenergylett.6b00319 ACS Energy Lett. 2016, 1, 698−704

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source pKa. Protonation of NH3 leads to a 325 mV anodic shift in reduction potential between N2(g)/NH3(MeCN) and N2(g)/ NH4+(MeCN). Similarly, a difference of 245 mV is observed for the N2(g)/N2H4(MeCN) and N2(g)/N2H5+(MeCN) potentials. While the aqueous data exhibit the same trend, the anodic shift upon protonation is substantially smaller in water (e.g., 183 mV from NH3 to NH4+), likely a result of the difference in the pKa scales between solvents. Figure 1 highlights the preference for NH3/NH4+ over all other N2 reduction products. The thermodynamic landscape of N2 reduction can be compared to the analogous multielectron, multiproton reductions of CO2. Products of CO2 reduction include the 2e−-reduced products CO and formic acid (HCO2H) along with the 4e−-reduced product formaldehyde (H2CO), the 6e−reduced product methanol, and the 8e−-reduced product methane.46 While highly reduced products such as methanol and methane are thermodynamically favored, CO and HCO2H are much more commonly observed in electrocatalytic reactions.26,25,28 For CO2 reduction in MeCN, the CO2/CO potential has been estimated as 250 mV more negative than the CO2/CH4 potential.26 It is shown here that the N2(g)/ N2H4(MeCN) potential is 430 mV more negative than the N2(g)/NH3(MeCN) potential. The role of acid/base chemistry is also reminiscent of CO2 reduction to HCO2H. The conversion of CO2 and H2 to HCO2H is unfavorable in acid, but the addition of base can generate formate and render the overall reaction exergonic.47 In the case of N2 reduction, NH3 is thermodynamically favored with respect to N2 and H2, but conducting the reaction under acidic conditions significantly increases the thermodynamic driving force via formation of NH4+. In addition to considering selectivity within the N2 reduction manifold, any reductive process involving protons is subject to competing proton reduction to give H2. In current chemical N2 reduction reactions, for example, only moderately selective production of NH3 relative to H2 is achieved.9−13 The relevant acid-dependent electrochemical potential for the H+/H2 couple at 1 M ionic strength in MeCN is −0.028 V vs Fc+/0, with a 59 mV cathodic shift per pKa unit.35,38 N2 reduction to NH3 is favored by 63 mV over proton reduction when the acid source has a pKa value greater than 16.5. When stronger acids are employed, NH4+ is produced with greater thermodynamic selectivity. For example, anilinium salts (pKa = 10.62),48 which are commonly employed in electrocatalytic reductions,49−51 give a thermodynamic potential for N2(g)/NH4+(MeCN) of −0.478 V vs Fc+/0, a full 177 mV positive of the thermodynamic potential for anilinium reduction to generate H2 (−0.655 V vs Fc+/0). This comparison highlights the difference in pKa dependence for the H+/H2 and N2/NH4+ potentials: N2 reduction is favored over H+ reduction by a wider margin when stronger acids are employed. This tunability based on product protonation state is an important departure from the standard 59 mV per pKa unit dependence of the proton reduction process, offering a possible means of influencing the relative product selectivity. Any electrocatalytic N2 reduction will compete with the direct reduction of protons at the electrode. At sufficiently negative potentials, all electrodes will catalyze proton reduction, so each electrode has a “potential window” in which acidpromoted electrochemical reactions can be efficiently catalyzed. Beyond this window, the current is dominated by the background processes of the electrode, thus preventing the

N2(g ) + 2HBase+(MeCN ) + 2e− ⇌ N2H 2(MeCN ) + 2Base

ΔG° = 56.46 + (2· 1.364· pK a) kcal ·mol−1 (24)

E°(N2(g )/N2H 2(MeCN )) = −1.22 V − 0.059· pK a

(25)

Table 1 summarizes the N2 reduction standard potentials in MeCN for the products described above, along with the Table 1. Summary of N2 Reduction Potentials in MeCN and H2Oa E° (V) b

reduction process

MeCN

N2(g)/NH3(solv) N2(g)/NH4+(solv) N2(g)/N2H4(solv) N2(g)/N2H5+(solv) N2(g)/N2H2(solv)

+0.035 − 0.059·pKa +0.361 − 0.079·pKa −0.398 − 0.059·pKa −0.153 − 0.074·pKa −1.22 − 0.059·pKa

H2O +0.092 − 0.059·pH +0.275 − 0.079·pH −0.33 − 0.059·pH −0.23 − 0.074·pH −1.20 − 0.059·pH

a

The pKa is that of the acid present in the solution that provides the protons. The potentials are for solutions with a 1:1 ratio of [HA]:[A−] (more rigorously, a 1:1 ratio of activities, though these are very rarely available). The standard potentials refer to conditions of 1 atm N2 and 1 M concentration of the reduced product (e.g., NH3). MeCN potentials are referenced to Fc+/0; aqueous potentials are referenced to SHE. bEstimated uncertainty ±30 mV in most cases (Supporting Information, section S9).

corresponding aqueous data for comparison. Comparing the potentials of the neutral species in MeCN is informative. The 4e− reduction of N2 to N2H4 has a standard potential 0.43 V more negative than the 6e− reduction of N2 to NH3. The N2/ N2H2 potential is more negative than the N2/NH3 potential by a full 1.26 V. N2H2 is unlikely to be observed under electrocatalytic conditions, given extremely unfavorable N2 reduction thermodynamics and rapid disproportionation to more easily reduced hydrazine. A viable electrochemical synthesis of ammonia should exhibit high selectivity for ammonia (or ammonium) over other N2 reduction products. The formation of byproducts would reduce Faradaic efficiency and necessitate costly separations. The MeCN data is depicted graphically in Figure 1, with thermodynamic potential (E°) plotted as a function of proton

Figure 1. Comparison of the thermodynamic potential (E°) of each N2 reduction process in MeCN as a function of proton source pKa. The vertical dashed line indicates the pKas of NH4+ and N2H5+. 701

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reaction and the applied potential required to achieve a desired current density.24,55 A larger overpotential corresponds to a less efficient process. Though no molecular electrocatalysts for N2 reduction have been reported, there are a few examples of stoichiometric electroreduction to form NH3.20−22 For example, Pickett and co-workers reported stepwise protonation of a W−N2 complex with tosic acid followed by electroreduction at −2.7 V vs Fc+/0 in THF,20,21 which would correspond to an overpotential of 2.38 V if it were electrocatalytic in MeCN. The standard potentials for N2 reduction provided in this study allow for the calculation of overpotential for future electrocatalysts. The catalytic efficiency of N2 reduction systems driven by chemical reductants can also be interpreted along similar lines. The thermodynamic driving force under catalytic conditions can be estimated by comparing the potentials of the chemical reductant that was employed with the thermodynamic potential for the specific proton source used (although our estimates are for MeCN rather than heptane solvent). Previous studies by Schrock and Yandulov9 and Nishibayashi and co-workers10 utilized decamethylchromocene (E° = −1.44 V vs Fc+/0 in MeCN)56,57 and cobaltocene (E° = −1.305 V vs Fc+/0 in MeCN),57,58 respectively, as reducing agents. 2,6-Lutidinium (pKa = 14.13 in MeCN)48 was the proton source in both cases. Using eq 19 above with 2,6-lutidinium as the acid gives the N2(g)/NH4+(MeCN) potential as E° = −0.76 V vs Fc+/0 in MeCN. The combination of decamethylchromocene and 2,6lutidinium in MeCN provides 0.68 V beyond the thermodynamic potential, leading to 94 kcal·mol−1 of driving force for the six-electron N2 reduction reaction. Similarly, the use of cobaltocene and 2,6-lutidinium to reduce N2 is exergonic by 76 kcal·mol−1, which translates to 12.7 kcal·mol−1 per electron or an overpotential of 0.55 V. DFT calculations for the Schrock system estimate that the reaction is exergonic by more than 120 kcal·mol−1 in heptane.16,59 While it is challenging to make meaningful comparisons between acetonitrile and heptane, and the thermodynamic details of each specific intermediate will play a role in catalysis, the magnitude of the required thermodynamic driving force in each case highlights the energetic inefficiency of current chemical N2 reduction systems that must be addressed in order to compete with the HaberBosch process. A complete thermodynamic evaluation of all products expected from electrochemical N2 reduction has been undertaken for the first time in an organic solvent. With the aid of thermochemical cycles, the standard potentials for the reduction of dinitrogen to ammonia, hydrazine, diazene, ammonium, and hydrazinium in MeCN solution have been derived. In a direct comparison of these values to the known H+/H2 potential in MeCN, the most favorable process is the formation of ammonium, about 400 mV positive of proton reduction. One interesting ramification of N2 reduction in MeCN solution is an increased thermodynamic preference for N2 reduction over H+ reduction: the difference between the N 2 /NH 4 + process and the H + /H 2 process was more pronounced in MeCN (390 mV) than in water (275 mV). The thermodynamic potential for ammonium formation is positive of glassy carbon electrode-mediated H+ reduction by up to 1.4 V, providing a significant potential window in which N2 reduction can proceed. The thermodynamic preference for NH3/NH4+ over all other species is encouraging for future electrochemical studies, especially given the kinetic challenge of N2 reduction under mild conditions.

observation, interrogation, or productive use of the electrochemical processes of interest. In preliminary screening of soluble electrocatalysts, glassy carbon is the most commonly used electrode for electroreductions in organic solvents. Dempsey and co-workers have recently examined hydrogen production at glassy carbon electrodes and reported the potentials at which a wide range of Brønsted acids begin to be reduced at the electrode in MeCN.52 These values provide an estimate of the cathodic limit at which electrochemical processes can be studied for each particular acid. This discussion is for glassy carbon; other electrode materials will have different solvent- and acid-specific H2 evolution behavior.53,54 Figure 2 shows how the thermodynamic potentials for H+ reduction and N2 reduction to NH3/NH4+ fit within the

Figure 2. Thermodynamic potentials of N2 reduction to NH4+ (green line) or NH3 (red line) and H+ reduction to H2 (purple line) as a function of proton source pKa. The vertical dashed line indicates the pKa of NH4+. The shaded region represents the applied potentials where the kinetics of H+ reduction to H2 at a glassy carbon electrode would lead to large currents. The dotted line represents the linear fit of E° vs pKa for several well-behaved organic aids taken from Dempsey and co-workers (E° = −1.08(10) V − 0.056(6)·pKa).52

effective potential range on glassy carbon for various acids in MeCN. (Figure 1 can be used to make similar comparisons for other N2 reduction processes.) The shaded region in Figure 2 represents potentials where electrode-catalyzed proton reduction is anticipated to dominate the observed currents. This onset potential is based on a linear fit of E vs pKa for a subset of well-behaved organic acids described by Dempsey and coworkers.52 There is a potential window of up to 1.4 V wherein N2 reduction to NH4+ or NH3 is thermodynamically favorable before the kinetic onset of glassy carbon electrode-mediated H+ reduction. Because of the difference in pKa dependence between the two processes, this gap is largest when strong acids are employed. This result is encouraging because it indicates that even N2 reduction processes with reasonably high catalytic overpotentials may be observed and studied with limited interference by background electrode processes. Equipped with the standard potentials for N2 reduction, we can now address the issue of catalytic efficiency. One metric for electrocatalytic efficiency is the overpotential, defined as the difference between the thermodynamic potential of a particular 702

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsenergylett.6b00319. Derivation of aqueous standard reduction potentials, experimental details, computational details, and additional detailed thermochemical cycles (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS B.M.L., K.K.-J., J.M.M., and A.J.M.M. gratefully acknowledge support from the NSF Center for Enabling New Technologies through Catalysis (CENTC), CHE-1205189. For A.M.A., this research was supported as part of the Center for Molecular Electrocatalysis, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences. Pacific Northwest National Laboratory is operated by Battelle for the U.S. Department of Energy.

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REFERENCES

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DOI: 10.1021/acsenergylett.6b00319 ACS Energy Lett. 2016, 1, 698−704