Electrocatalytic Hydrogen Production by a Nickel Complex Containing

Article Cite This: Organometallics XXXX, XXX, XXX−XXX

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Electrocatalytic Hydrogen Production by a Nickel Complex Containing a Tetradentate Phosphine Ligand Christina M. Klug,† William G. Dougherty,‡,§ W. Scott Kassel,‡ and Eric S. Wiedner*,† †

Center for Molecular Electrocatalysis, Pacific Northwest National Laboratory, P.O. Box 999, K2-57, Richland, Washington 99352, United States ‡ Department of Chemistry, Villanova University, 800 East Lancaster Avenue, Villanova, Pennsylvania 19085, United States

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S Supporting Information *

ABSTRACT: A nickel complex has been synthesized using the P4N2 ligand, a tetradentate phosphine ligand with two pendant amines in the ligand backbone. The rigidly square pyramidal [Ni(P4N2)(CH3CN)]2+ complex is an electrocatalyst for the reduction of protons to hydrogen. Using N,N-dimethylformamidium ([DMF(H)+]) as the acid in an acetonitrile/water solution, [Ni(P4N2)(CH3CN)]2+ displays a turnover frequency of 1.6 × 106 s−1, which is among the fastest rates reported for any molecular electrocatalyst. This high catalytic rate comes at the cost of a 1200 mV overpotential at the catalytic half-wave potential. The Ni(II) state of the catalyst was found to be stable under strongly acidic conditions, with only trace decomposition observed over 2 weeks in the presence of 0.1 M trifluoromethanesulfonic acid (HOTf). The catalyst was less stable in the Ni(0) state due to the inability of the rigid P4N2 ligand to adopt a tetrahedral geometry. Using variable scan rate cyclic voltammetry, a first-order rate constant of ∼1 s−1 was measured for dissociation of a phosphine from Ni(0), which is proposed to be a key step for determining the lifetime of the catalyst during electrolysis.

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INTRODUCTION

Hydrogen obtained by electrocatalytic proton reduction has potential application as an energy vector for storage of renewable energy. While metallic Pt is widely recognized for its high activity and low overpotential for electrocatalytic H2 evolution, its scarcity and high cost pose a barrier to its widespread use.1 Efforts to design molecular electrocatalysts for H2 production have focused on complexes of the earthabundant metals,2−4 particularly complexes of cobalt,5−7 nickel,8−10 iron,11−13 and molybdenum.14,15 Catalytic turnover frequency (TOF) and overpotential are frequently the focus of designing improved molecular electrocatalysts. An equally important metric is the catalyst lifetime, particularly under strongly acidic conditions (pH 0−1) that are technologically relevant for use in polymer electrolyte membrane (PEM) electrolyzers. Molecular catalysts are notable for decomposing under highly acidic conditions and, in some cases, decompose into catalytically active heterogeneous films.16−20 Examples of catalysts that can operate under highly acidic conditions include Ni(diphosphine)221−23 and Co(azamacrocycle)6 complexes, as well as Co(dithiolate)2 complexes that are adsorbed to a carbon electrode24,25 or incorporated into coordination polymers.26−28 Tetradentate “P4N2” ligands have previously been used to design catalysts for H2 production,29,30 CO2 hydrogenation,31 and N2 silylation32 (Figure 1). In each of these catalysts, the rigid coordination geometry of the P4N2 ligand was found to © XXXX American Chemical Society

Figure 1. Previously reported molecular catalysts employing a P4N2 ligand.

impart unique characteristics to one or more of the catalytic intermediates. In this study, we describe the synthesis, H2 production activity, and acid stability of [Ni(P 4 N 2 )(CH3CN)]2+ (1). The primary motivation for this work is to Special Issue: Organometallic Electrochemistry: Redox Catalysis Going the Smart Way Received: July 31, 2018

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DOI: 10.1021/acs.organomet.8b00548 Organometallics XXXX, XXX, XXX−XXX

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Organometallics identify how the tetradentate coordination environment and presence of a single P2N2 moiety affects the catalytic response relative to the well-known [Ni(P2N2)2]2+ family of catalysts.8 With TOF > 106 s−1, 1 is one of the fastest molecular electrocatalysts reported for H2 production, though this high rate is accompanied by a 1200 mV overpotential. Additionally, 1 is highly stable under strongly acidic conditions due to the rigid tetradentate ligand framework. Spectroscopic studies provide insight into the mechanism of catalyst decomposition in the presence of acid.

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RESULTS Catalyst Synthesis. Treatment of the P4N2 proligand29 with 1 equiv of [Ni(CH3CN)6]2+ afforded 1 as a yellow powder (eq 1). The 1H NMR spectrum of 1 in CD3CN

Figure 3. X-ray crystal structural depiction of 1. Hydrogen atoms, two BF4− anions, and solvent molecules are omitted for clarity. Thermal ellipsoids are shown with 50% probability.

indicates that both arms of the cyclic P2N2 backbone are magnetically equivalent. In addition, 1 equiv of free CH3CN is observed at 1.96 ppm. These data indicate 1 contains a weakly bound CH3CN ligand that exchanges rapidly with the CD3CN solvent. Complex 1 displays a second-order 31P{1H} NMR spectrum centered at 63.0 ppm, which was simulated as an AA′BB′ coupling pattern with two unique phosphorus resonances at 63.7 and 62.3 ppm (Figure 2).

Figure 4. Cyclic voltammograms of 1 at varying scan rates. The current is normalized to υ−1/2 to facilitate comparison of data collected at different scan rates. Conditions: 1.9 mM 1, 0.2 M [Bu4N][PF6] in benzonitrile, 1 mm diameter glassy carbon working electrode.

chemically irreversible at slow scan rates (υ) and become more reversible at higher scan rates (Figure 4). In acetonitrile solvent, two stripping waves are observed on the return sweep at scan rates less than 1 V s−1 (Figure S14), suggesting that the neutral Ni(0) species are insoluble in acetonitrile and precipitate on the electrode surface.35 No evidence of Ni(0) precipitation was observed in benzonitrile solvent (Figure S14), consistent with previous observations on the solubility of Ni(diphosphine)2 complexes.35,36 At a scan rate of 5 V s−1, the half-wave potentials (E1/2) of the two reduction waves were measured to be −1.52 and −1.68 V versus Cp2Fe+/0 in acetonitrile and −1.50 and −1.70 V in benzonitrile. These reduction waves are assigned to the Ni(II/I) and Ni(I/0) couples of 1, respectively, and are significantly more negative than those typically observed for [Ni(diphosphine) 2]2+ complexes. For example, [Ni(dmpe)2]2+ (dmpe = 1,2-bis(dimethylphosphino)ethane) displays a two-electron Ni(II/0) couple at −1.35 V due to the electron-rich dmpe ligands.37 The negative potentials of 1 are a consequence of the rigidity of the P4N2 ligand, which prevents 1 from adopting a tetrahedral geometry at the Ni(I) and Ni(0) oxidation states.38,39 On the reductive sweep, the cathodic peak current (ipc) for the Ni(II/I) and Ni(I/0) couples is linear versus the square root of the scan rate (Figures S16 and S17), indicating that these reductions are controlled by diffusion of the complex to the electrode surface. Additionally, the ratio of the ipc values of the Ni(II/I) and Ni(I/0) couples does not vary substantially

Figure 2. Experimental (top) and simulated (bottom) 31P{1H} NMR spectra of 1 (200 MHz).

Slow diffusion of diethyl ether into an acetonitrile solution of 1 afforded crystals that were suitable for X-ray diffraction, and the solid-state structure of 1 confirms the proposed geometry (Figure 3). The Ni−P bond lengths of 1 are 2.2031(6) and 2.2310(6) Å for the terminal PPh2 groups, and 2.1533(6) and 2.1683(6) Å for the P2N2 backbone. Notably, 1 displays a square-pyramidal geometry in the solid state with an axial CH3CN ligand, with a τ5 value of 0.02.33 In contrast to 1, the related complex [Ni(tetraphos)(CH3CN)]2+ (tetraphos = (R,R)-Ph2CH2CH2P(Ph)CH2CH2P(Ph)CH2CH2PPh2) displays τ5 values of 0.24 and 0.40 for two crystallographically independent molecules,34 indicating a moderately to strongly distorted square pyramidal geometry. This structural difference is a consequence of the cyclic P2N2 fragment in the P4N2 backbone, which possesses a small P−Ni−P angle of 80.23(2)°. Electrochemistry. Cyclic voltammograms of 1 in acetonitrile or benzonitrile display two reduction waves that are B


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Organometallics with the scan rate. These observations indicate that the irreversible chemical reaction occurs at the Ni(0) state of 1. A rate constant of ∼1 s−1 can be estimated for decomposition of the electrochemically generated Ni0(P4N2) species, on the basis of the reversibility of the Ni(I/0) couple in benzonitrile as a function of the scan rate (see the Supporting Information for details). To characterize the reduction product, a 5 mM solution of 1 was reduced by two electrons using controlledpotential electrolysis. A 31P{1H} NMR spectrum of the resulting solution showed a complex mixture of peaks in the range 70−30 ppm (Figure S20). These resonances are shifted substantially downfield from the resonances for the free P4N2 ligand (−12 and −46 ppm), indicating that the decomposition pathway does not involve complete ligand dissociation. A plausible assignment for the decomposition pathway is dissociation of one or more phosphines from Ni0(P4N2), followed by formation of multimetallic species in which the P4N2 ligand bridges two or more nickel centers. In support of this assignment, Ni0(tetraphos) has been shown to isomerize to [Ni02(tetraphos)2], in which the tetraphos ligands bridge both Ni centers.34 Electrocatalytic Production of H2. Addition of [DMF(H)][OTf] to dilute acetonitrile solutions of 1 (ca. 0.3 mM) results in a significant current enhancement near the Ni(II/I) couple, indicative of electrocatalytic H2 production (Figure 5a). Direct reduction of [DMF(H)]+ by the glassy carbon electrode produces a significant background current at these negative potentials.30,40 However, the background current is less than the current observed in the presence of 0.2−0.3 mM 1; therefore, the catalytic current (icat) resulting from 1 was determined by subtracting the background current from the total observed current. The catalytic turnover frequency (TOF) was determined from eq 2, where F is the Faraday constant, υ is the scan rate in V s−1, n is the number of electrons consumed in the catalytic cycle (two electrons), R is the gas constant, T is the temperature in K, and ip is the peak current for the noncatalytic Ni(II/I) redox couple.41−43 The TOF increases linearly with [DMF(H)]+ concentration at low acid concentrations (≤100 mM) and becomes pseudoindependent of [DMF(H)]+ concentration as the acid concentration is increased further (Figure 5b). This dependence on acid concentration is indicative of saturation kinetics. A TOF of 200000 s−1 was measured for 1 at 0.4 M [DMF(H)]+. TOF =

0.199Fυ ijjj icat yzzz j z n2RT jjk i p zz{

Figure 5. (a) Cyclic voltammograms with varying amounts of 1, [DMF(H)]+, and H2O in 0.2 M [Bu4N][PF6] acetonitrile solution with υ = 15 V s−1: (black trace) [1] = 0.3 mM; (red trace) [DMF(H)]+ = 0.44 M, [H2O] = 6.5 M; (blue trace) [1] = 0.2 mM, [DMF(H)]+ = 0.46 M; (green trace) [1] = 0.2 mM, [DMF(H)]+ = 0.44 M, [H2O] = 6.5 M. (b) Plot of TOF versus concentration of [DMF(H)]+. (c) Plot of TOF versus concentration of H2O with [DMF(H)]+ = 0.46 M.

Acid Stability of 1. To test its stability under acidic conditions, a CD3CN solution of 1 (16 mM) and [DMF(H)]+ (5 M) was monitored by 31P{1H} NMR spectroscopy. The initial spectrum of this reaction solution showed a 1:1 ratio of two resonances, a pseudodoublet at 64 ppm and a broad resonance at 58 ppm (Figure S6). This difference in these chemical shifts (Δδ = 6) is greater than those observed for the two 31P resonances of 1 (Δδ = 1.1) in the absence of acid. Protonation of [Ni(P2N2)2]2+ complexes at a pendant amine shifts the 31P{1H} resonance upfield by 20 ppm.47,48 Since a much smaller upfield shift is observed for 1 in the presence of [DMF(H)]+, an equilibrium between protonated and unprotonated pendant amines is likely present. No further change in the 31P{1H} NMR spectrum was observed over 15 days, indicating that 1 is highly stable under these conditions. A similar study was performed with varying amounts of trifluoromethanesulfonic acid (HOTf) ,which is more acidic (pKaMeCN = 2.6)49 than [DMF(H)]+ (pKaMeCN = 6.1).49 31 1 P{ H} NMR spectra of 1 (10 mM) were recorded 1 h after the addition of HOTf in CD3CN, shown in Figure 6a. The spectrum of 1 with 0.1 M HOTf shows one major species (2,

2

(2) 2+

44

As is typical for [Ni(P2N2)2] electrocatalysts, addition of H2O subsequent to the [DMF(H)]+ led to a further increase in the catalytic current. A maximum current enhancement is observed at approximately 6 M H2O; further increases in H2O concentration led to a slight decrease in TOF. Under these conditions the catalytic TOF is 1600000 s −1, which corresponds to the maximum turnover frequency (TOFmax) attainable by 1 under these conditions (Figure 5c). The catalytic overpotential of 1 is 1200 mV, on the basis of the difference between the half-wave potential for the catalytic wave (−1.63 V) and the equilibrium potential of 1/1 DMF/ [DMF(H)]+ in wet acetonitrile under 1 atm of H2 (−0.43 V).45,46 Hydrogen was confirmed as the product with a Faradaic efficiency of 90 ± 5% by gas chromatographic analysis of the headspace gas obtained from a bulk electrolysis experiment. C


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be an exo-pinch complex in which one of the terminal PPh2 groups has dissociated from Ni and become protonated (Figure 6c, right). Dissociation of the PPh2 ligand is likely accompanied by coordination of either an acetonitrile or triflate ligand to occupy the fourth coordination site. Full details of the spectroscopic assignment of 3 are provided in the Supporting Information. Over the course of 2 weeks, 2 persists as the primary species in the presence of 0.1 M HOTf. At a higher concentration of 0.5 M HOTf, 2 and 3 gradually decompose into a complex mixture of products over 2 weeks (Figure S8). Several of the 31 1 P{ H} NMR spectroscopic resonances are similar to those observed when the free P4N2 proligand is treated with HOTf (Figure S10). When 1 is treated with 1.0 M HOTf, complexes 2 and 3 are observed as the primary species after 1 h of reaction, but after 3 days the complexes have decomposed into unbound protonated P4N2 (Figure S9). As a surrogate for the protonated species, a complex containing a methylated pendant amine (4) could be obtained by treating 1 with trimethyloxonium tetrafluoroborate (eq 3).

The 1H NMR spectrum of 4 indicates the presence of a single regioisomer containing two inequivalent pendant amine arms and one N−CH3 group (Figure S3). The 31P{1H} NMR spectrum of 4 showed a second-order AA′BB′ pattern centered at 58.5 ppm (Figure S4). This spectrum is similar in appearance and is shifted downfield from the spectrum of 1. The solid state structure of 4, obtained by X-ray diffraction analysis, reveals that the N−CH3 group is located exo with respect to the Ni center (Figure 7). In contrast to 1, complex 4

Figure 6. (a) 31P{1H} NMR spectra of 1 in CD3CN with varying amounts of HOTf after 1 h of reaction. (b) Expanded 31P{1H} NMR spectra of 1 in CD3CN with 0.5 M HOTf showing the four resonances for species 3. (c) Proposed products formed in the reaction of 1 with HOTf.

green circles) with a pseudodoublet at 63.6 ppm and a broad resonance at ∼44 ppm. This spectrum resembles that observed using 5 M [DMF(H)]+, except the difference in chemical shifts is much greater when HOTf is used (Δδ = 19.6). When the concentration of HOTf is increased to 0.5 M, the broad resonance resolves into a doublet with a P−P coupling constant that matches the resonance at 63.6 ppm. This species is assigned to an “exo-pinch” complex (2), where both pendant amines are bound to a single proton (Figure 6c, left). In the case of [Ni(P2N2)2]2+ complexes, protonation of Ni(II) results exclusively in exo-pinch species,47,48 whereas the pendant amines of Ni(0) are protonated both exo and endo relative to Ni.50 A characteristic resonance for the exo proton of 2 was not observed at 10−12 ppm in the 1H NMR spectrum,30,47 presumably due to rapid exchange with the excess HOTf present in solution. After 1 h of reaction with acid, a second species (3) is observed in trace quantities with 0.1 M HOTf and is more pronounced with 0.5 M HOTf (Figure 6a, purple diamonds). The 31P{1H} NMR spectrum of this minor species displays four resonances in a 1:1:1:1 ratio, indicating that the complex has lost its 2-fold symmetry. Two of these resonances appear as a doublet of doublets (dd), one is a doublet of doublets of doublets (ddd), and the fourth resonance is a simple doublet (Figure 6b). In the 1H-coupled 31P NMR spectrum, the doublet at 9.6 ppm splits into a broad doublet with JHP ≈ 540 Hz, which is characteristic for a protonated phosphine (Figure S12). On the basis of these spectroscopic data, 3 is proposed to

Figure 7. X-ray crystal structural depiction of 4. Hydrogen atoms, three BF4− anions, and solvent molecules are omitted for clarity. Thermal ellipsoids are shown with 50% probability.

shows a square-planar geometry without a coordinated acetonitrile ligand. However, the phenyl ring located on the quaternary pendant amine is folded over the nickel center, resulting in a close approach of the ipso (3.216 Å) and ortho (3.147 Å) carbons to nickel. The terminal PPh2 groups of 4 show Ni−P bond lengths of 2.2199(5) and 2.2367(5) Å for the terminal PPh2 groups, while the P2N2 backbone shows Ni− D


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Organometallics P bond lengths of 2.1606(5) and 2.1707(5) Å. Despite the higher molecular charge of 4, its average Ni−P bond lengths are statistically indistinguishable from those of 1. To determine the stability of 1 under electrocatalytic turnover, an extended controlled-potential electrolysis was conducted at −1.84 V versus Cp2Fe+/0 on an acetonitrile solution containing 1.0 mM 1, 0.46 M [DMF(H)]+, 2.84 M H2O, and 4.6 mM [PPh4][BF4] as an integration standard. A total of 194 C was passed over a period of 3 h, corresponding to ∼50% consumption of the [DMF(H)]+. A 31P{1H} NMR spectrum recorded after electrolysis showed only 1 and the PPh4+ integration standard, with no change in relative intensity in comparison to a spectrum recorded prior to electrolysis (Figure S27). Therefore, the decomposition of 1 during electrolysis was less than the 31P{1H} NMR spectroscopic detection limit (∼10%). A rinse test was not performed, since background reduction of [DMF(H)]+ by the carbon electrode would obscure the results. As a result, we were unable to definitively rule out deposition of a small amount of catalytically active nanoparticles on the electrode surface.

Deactivation of 1 is presumed to initiate by dissociation of a terminal PPh2 ligand, followed by other chemical steps that drive dechelation of the P4N2 ligand. Several different pathways for deactivation of 1 during electrocatalysis were considered in which the phosphine dissociates from different intermediate states of the catalyst. One possible deactivation route is dissociation of PPh2 from the Ni(II) state of 1 (Figure 9, Path

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DISCUSSION Mechanism and Stability. In this work, we examined the effect of a tetradentate P4N2 ligand on the stability and catalytic activity for H2 production by 1. By analogy to [Ni(P2N2)2]2+ complexes, H2 production by 1 is proposed to operate by competing ECEC and EECC pathways,51−53 where E represents a 1e− reduction of the catalyst by the electrode and C represents a multistep sequence of protonation and deprotonation of the pendant amines. For example, the pendant amines of 1 can be protonated in either an endo or exo position with respect to the nickel center, as illustrated in Figure 8 for a proposed nickel hydride intermediate of 1. Endo

Figure 9. Proposed paths for deactivation of 1.

A). This deactivation step is thermodynamically unfavorable due to the chelate effect but could be driven toward dechelation by protonation of the dangling PPh2 group by an exogeneous acid. On the basis of the acidity of [HPMePh2]+ (pKaMeCN = 9.96),56 a PPh2 group that dissociates from 1 should be protonated by [DMF(H)]+ (pKaMeCN = 6.1).49 However, no evidence of phosphine dissociation and protonation was observed over 15 days in the presence of 5 M [DMF(H)]+, suggesting that Path A is not a major deactivation pathway of 1. A second proposed deactivation pathway is protonation of 1 to generate an exo-pinch species (2), followed by dissociation of the PPh2 group (Figure 9, Path B). Complex 2 is the major product observed after treating 1 with 0.1 M HOTf (pKaMeCN = 2.6,49 pKaH2O = −14.757), and only minimal decomposition of 2 was observed over 2 weeks in the presence of 0.1 M HOTf. In contrast, 2 displays a half-life of ∼2 days in the presence of 0.5 M HOTf. Complex 2 is estimated to have pKaMeCN ≈ 3 on the basis of the variance of its 31P{1H} NMR chemical shifts in the presence of different concentrations of HOTf (see the Supporting Information for more details). These data suggest that 2 is in equilibrium with ∼4% of 1 in the presence of 0.1 M HOTf and that phosphine dissociation is promoted by shifting the protonation equilibrium toward 2 at higher acid concentration. The underlying cause for the increased lability of the phosphine ligands in 2 relative to 1 is not clear. One possibility is that protonation of the P4N2 ligand decreases the strength of the Ni−P bonds. However, the Ni−P bond lengths of 1 and the N-alkylated complex 4 are statistically indistinguishable. A second possibility is that the tetradentate coordination mode of the P4N2 ligand is geometrically destabilized by constraining both pendant

Figure 8. Protonation at an endo site leads to fast catalysis, while protonation at an exo site leads to inactive species.

protonation places the proton near the nickel center and leads to fast catalysis, while exo protonation slows catalysis by trapping the proton away from the nickel center. The pendant amines can be protonated either endo or exo at the Ni(I), Ni(0), or Ni(II)-H states, leading to multiple branching points in the catalytic cycle. At high concentrations of acid, the TOF becomes independent of the acid concentration and is controlled by rate-limiting isomerization of these exo-pinch intermediates to catalytically active endo-protonated species. This exo to endo isomerization is facilitated by H2O and by DMF base that is generated in the diffusion layer during electrocatalytic turnover.54,55 E


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Organometallics amines to a boat conformation via the exo-pinch, though no structural metrics could be obtained for 2 to provide evidence in support of this hypothesis. A third deactivation pathway is dissociation of a phosphine from the Ni(0) state of 1 (Figure 9, Path C). In contrast to the high stability of the Ni(II) state of 1, the Ni(0) state is highly unstable, since the rigid tetradentate P4N2 ligand is unable to adopt a tetrahedral geometry. Electrochemical measurements afforded a first-order rate constant of ∼1 s−1 for decomposition of Ni(0), providing a relative ordering of C > B > A for the rate of phosphine dissociation via these three mechanisms. As a result, the Ni(0) state represents a key branching point, since this species can either react with acid to generate H2 or dissociate a phosphine, leading to deactivation of the catalyst. This hypothesis indicates a theoretical maximum of ∼1.6 × 106 catalytic turnovers prior to complete catalyst deactivation, obtained from the ratio of the catalytic TOFmax (1.6 × 106 s−1) and the first-order rate constant for decomposition of Ni(0) (∼1 s−1). We were unable to measure a maximum turnover number for 1, since a high catalyst concentration (1 mM) was necessary to ensure the catalytic current was substantially larger than background reduction of the acid by the electrode. However, no evidence of catalyst decomposition was observed over 120 turnovers during electrolysis, as expected on the basis of the theoretical maximum turnover number. Catalyst Activity. With a maximum TOF of 1.6 × 106 s−1, 1 is among the fastest reported molecular electrocatalysts for production of H2. To the best of our knowledge, [Ni(P2N2)2]2+ complexes22,44,54 and a binuclear Cu fused porphyrin complex58 are the only other H2 production electrocatalysts reported to have TOF > 106 s−1. Examples of other fast H2 production electrocatalysts (Figure 10a) include Fe(TPP)Cl (TOF = 4 × 105 s−1),59 a homobimetallic FeFe complex encapsulated in a polymer matrix (TOF = 2.5 × 105 s−1),60 and a heterobimetallic NiFe complex (TOF = 2.5 × 104 s−1).61 The TOF values reported for Fe(TPP)Cl and the NiFe complex are theoretical TOF values that would be observed at 1 M acid concentration (TOF1 M), on the basis of extrapolation of kinetic data obtained at low acid concentration in an acid concentration dependent regime. The actual TOFs measured for these two catalysts at 1 M acid could be substantially lower than those predicted due to either a change in the rate-limiting step of catalysis at higher acid concentration or catalyst instability in the presence of 1 M acid. In addition to the TOF, the catalytic overpotential is also an important metric for comparison of different electrocatalysts. Molecular Tafel plots are commonly used to plot the variation in TOF of a single catalyst over a wide potential range, where the Nernst equation is applied to account for the equilibrium between the oxidized and reduced forms of the catalyst at different potentials.59 Identification of the most efficient catalysts is facilitated by Tafel plot analysis through extrapolation of TOF to zero overpotential (TOFo). Tafel analysis, however, is not conducive to identification of structure−activity relationships that lead to rational improvement of catalyst structure. In this regard, examining how the maximum TOF varies as a function of the overpotential at the catalytic half-wave potential provides greater insight. This analysis has been used to identify scaling relationships between rate and overpotential within different catalyst families for H2 production,46,62 O2 reduction,63−65 and CO2 reduction.8,66 A more detailed comparison of these two methods is presented in the Supporting Information.

Figure 10. (a) Selected examples of molecular electrocatalysts for production of H2. (b) Plot of log TOFmax versus overpotential at Ecat/2 for selected molecular electrocatalysts. The solid line illustrates the electronic scaling relationship of [Ni(PR2NPh2)2]2+ catalysts. Data used for construction of this plot are given in Table S2 in the Supporting Information.

The electronic scaling relationship previously reported for [Ni(PR2NPh2)2]2+ catalysts46 serves as a useful reference point for comparison of the tradeoff between rate and overpotential for different catalysts (Figure 9). Changing the identity of the phosphorus substituent in [Ni(PR2NPh2)2]2+ catalysts affects the electron density at nickel through a combination of inductive and steric effects.62,67 Furthermore, the pendant amines of [Ni(PR2NPh2)2]2+ complexes are inductively coupled to the electron density of the nickel center.68 As a result, the pendant amines are more basic for [Ni(P R 2 N Ph 2 ) 2 ] 2+ complexes that possess a more negative reduction potential (and higher overpotential), leading to a lower barrier for F


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a long alkyl chain in the outer coordination sphere and by using adiponitrile as the solvent to increase the solution viscosity.46 Molecular Tafel plot analysis of this catalyst system results in TOF0 = 9 s−1, which is the largest value reported to date for a molecular H2 production electrocatalyst.

protonation of the pendant amines and a higher catalytic TOF.46,52 The shallow slope of this scaling relationship (1/200 mV−1) has been attributed to the competing formation of catalytically inactive exo-pinch isomers that become more difficult to deprotonate as the pendant amines become more basic in catalysts with a larger overpotential.46,52,53 As seen in Figure 10, the high TOF of 1 results primarily from its large overpotential of 1200 mV. Notably, 1 lies slightly below the electronic scaling relationship of the [Ni(PR2NPh2)2]2+ catalysts. This result is unexpected, since 1 possesses only one P2N2 moiety than can form a catalytically inactive exo-pinch, while [Ni(PR2NPh2)2]2+ complexes possess two P2N2 ligands that can each form an exo-pinch. For comparison, the catalyst [Ni(7PPh2NPh)2]2+ (light blue bar) is incapable of forming an exo-pinch and lies substantially above the scaling relationship.69 One possible explanation is that the rigid P4N2 ligand causes the reduced forms of 1 to adopt a square-planar geometry, which could lead to a large barrier for endo protonation at the pendant amine due to steric interactions between the exogeneous acid and the opposing phenyl groups of the terminal phosphines. In contrast to 1, [Ni(PR2NPh2)2]2+ complexes have greater structural flexibility and can more easily adopt a tetrahedral geometry upon reduction, which could provide easier access to the pendant amines by the acid substrate. Selected examples of other molecular H2 production electrocatalysts are also shown in Figure 10. Comparison of the previously mentioned catalysts displaying high TOFs reveals that they lie very close to the [Ni(PR2NPh2)2]2+ scaling relationship. One exception is the binuclear Cu fused porphyrin catalyst,58 which lies 700 mV to the right of the scaling relationship. As an additional example, the Co(dmgBF2)2 catalyst is well-known for having a low overpotential for H2 production.70,71 In a detailed kinetic analysis, Dempsey demonstrated that the TOF of the Co(dmgBF2)2 catalyst is limited by a chemical step that is independent of the acid concentration.72 These kinetic measurements allow placement of Co(dmgBF2)2 on the plot of TOF versus overpotential,46 where it is found to coincide with the [Ni(PR2NPh2)2]2+ scaling relationship. None of these catalysts possess a pendant amine to serve as a proton relay; therefore, they cannot follow a mechanism similar to that for the [Ni(PR2NPh2)2]2+ catalysts. As such, there is no obvious reason these catalysts should lie so close to the [Ni(PR2NPh2)2]2+ scaling relationship. The observation of a general scaling trend between TOF and overpotential across several different catalyst families could be indicative of an upper threshold to the performance of traditional molecular catalysts that solely rely on modification of the ligand environment immediately surrounding the active metal site. Recent advances in [Ni(P2N2)2]2+ catalyst design have shown that the parent electronic scaling relationship can be circumvented by slowing the rate of boat to chair isomerization of the pendant amines to kinetically inhibit the formation of noncatalytic exo-pinch species. This has previously been accomplished by incorporating long alkyl chains on the outer coordination sphere of the pendant amines22 or by performing electrocatalysis in either a protic ionic liquid73 or a viscous dinitrile solvent.46 One example from our group is shown in Figure 10 (green diamond), where the overpotential was lowered using an electron-withdrawing phosphorus substituent, and the boat to chair isomerization rate of the pendant amine was slowed through incorporation of

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CONCLUSIONS A nickel complex containing a tetradentate P4N2 ligand, [Ni(P4N2)]2+, was synthesized and tested for electrocatalytic proton reduction. The rigid square-planar coordination geometry of the P4N2 ligand affected the electrocatalytic activity of the catalyst in several ways. First, the ligand was found to stabilize the Ni(II) state, leading to a very high catalyst stability under acidic conditions and a large overpotential of 1200 mV for production of H2. Second, the Ni(0) state of the catalyst was destabilized by the P4N2 ligand, which cannot adopt a tetrahedral coordination geometry. As a result, the Ni(0) state is a highly reactive species that can either enter a catalytic cycle through reaction with exogenous acid, or dissociate one or more arms of the phosphine ligand, leading to catalyst deactivation. The catalytic turnover frequency (1.6 × 106 s−1) was found to be much faster than the rate of catalyst deactivation (∼1 s−1).

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EXPERIMENTAL SECTION

General Procedures. All manipulations were conducted under a N2 atmosphere using standard Schlenk techniques or in a Vacuum Atmospheres glovebox. Bulk solvents were deoxygenated by sparging with N2 and dried by passage through neutral alumina column in an Innovative Technology, Inc., PureSolv solvent purification system. Benzonitrile was sparged with N2 and dried over 5 Å molecular sieves. Water was dispensed from a Millipore Milli-Q purifier and sparged with N2. Acetonitride-d3 (CD3CN) was dried over P2O5 and purified by vacuum transfer. Ferrocene was purified by sublimation, and [Bu4N][PF6] and [Bu4N][BF4] were recrystallized two to three times and dried under vacuum. All other commercial reagents were used as received. [Ni(CH3CN)6][BF4]2,74 P4N2 ligand,29 and [DMF(H)][OTf]75 were prepared using literature procedures. NMR spectra were recorded on a 500 MHz (1H frequency) Agilent spectrometer. 1H NMR chemical shifts were referenced to the residual protio solvent, and 13C{1H} NMR spectra were referenced to naturally abundant 12C solvent.76 31P{1H} NMR spectra were referenced to an external phosphoric acid standard (0 ppm). Second-order 31P{1H} NMR spectra were simulated using the NUMMRIT algorithm77 as implemented in SpinWorks.78 Cyclic voltammetry experiments were performed using a CH Instruments 620D potentiostat. A standard three-electrode cell was employed, using a 1 mm diameter PEEK-encased glassy carbon disk as the working electrode, a glassy carbon rod as the counter electrode, and a silver wire suspended in 0.2 M [Bu4N][PF6] acetonitrile and separated from the analyte solution by a Vycor frit as the pseudoreference electrode. Ferrocene was used as an internal standard, and all potentials are referenced to the Cp2Fe+/0 couple (0 V). Controlled-potential electrolysis experiments were performed using either a CH Instruments 620D potentiostat or a BASi Epsilon potentiostat interfaced with a PWR-3 power module. The working electrode was either a Pt mesh (noncatalytic experiments) or a cylinder of reticulated vitreous carbon that was attached to a copper wire using conductive silver epoxy (catalytic experiments). The counter electrode was a Nichrome wire suspended in 0.2 M [Bu4N][BF4] acetonitrile solution and was contained inside a glass tube capped at one end with a 12 mm diameter ultrafine frit. The pseudoreference electrode was identical with that used in cyclic voltammetry experiments. The electrolysis cell was a cylindrical 30 mL glass vessel which was capped with a rubber septum that was pierced to accept the electrodes. G


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Organometallics

was recorded at υ = 15 V s−1. Failure to polish the electrode resulted in a diminished catalytic current due to electrode fouling. After the acid addition was complete, the described method was repeated for the addition of 160 μL of H2O in 5−20 μL increments. At each acid and water concentration, the ip value for 1 was corrected for dilution by normalizing against the anodic peak current for the Cp2Fe+/0 couple. A potential of −1.85 V was selected for measuring icat, which was determined by finding the most positive potential for which the second derivative of the catalytic voltammogram is constant. A calibration curve for direct reduction of [DMF(H)][OTf] by the working electrode was constructed, and all icat values for 1 were corrected for the background current due to direct acid reduction. Determination of Faradaic Efficiency. Controlled-potential electrolysis was performed on 9.2 mL of an acetonitrile solution containing 1 mM 1, 0.5 M [DMF(H)][OTf], 4 M H2O, 0.2 M [Bu4N][BF4], and a small amount of ferrocene as an internal potential reference. Electrolysis was conducted at −1.84 V until 10.01 C of charge had passed. The headspace was sampled using a gastight syringe and analyzed using an Agilent 6850 gas chromatograph that was equipped with a 10 ft length Supelco 1/8 in. Carbosieve 100/120 column and a thermal conductivity detector. The percentage of H2 in the headspace was determined to be 6.9% by calibration against two different gas standards of known composition. The initial volume of N2 in the headspace (16.6 mL) was determined by subtracting the volume of analyte solution from the total cell volume with the electrodes attached. On the basis of the volume of N2 and the %H2 in the postelectrolysis headspace, 4.7 × 10−6 mol of H2 was produced during electrolysis, corresponding to a current efficiency of 90 ± 5% for production of H2. Stability of 1 During Electrocatalysis. Controlled-potential electrolysis was performed on 8.8 mL of an acetonitrile solution containing 1.0 mM 1, 0.46 M [DMF(H)][OTf], 2.8 M H2O, 0.2 M [Bu4N][BF4], and 4.6 mM [PPh4][BF4] as an internal integration standard for 31P{1H} NMR spectroscopic analysis. Several needles were placed in the septum cap of the electrolysis cell to vent the H2 that formed. Electrolysis was conducted at −1.84 V, and the solution in the anode compartment was periodically replaced with fresh electrolyte solution. After 194 C of charge had passed, the electrolysis solution was analyzed by 31P{1H} NMR spectroscopy (60 s relaxation delay, 700 transients). The relative integration between 1 and [PPh4]+ was consistent between spectra collected before and after electrolysis. Acid Stability of 1. A 50 mM stock solution of 1 and 2 M stock solution of HOTf were made in CD3CN. In a J. Young NMR tube, aliquots of the stock solutions of 1 (120 μL) and HOTf (30 μL for 0.1 M, 150 μL for 0.5 M, and 300 μL for 1 M) were combined with additional CD3CN to bring the total solution volume to 600 μL. The solutions were thoroughly mixed and then analyzed by 31P{1H} NMR spectroscopy periodically over a 15 day period. The 31P{1H} NMR spectra were collected using a 1 s relaxation delay, 5 μs pulse width, and 110 transients. Over a period of 1 week, the color of the solution containing higher acid concentrations changed from light yellow to light red. Ligand Protonation. In a representative experiment, 5.2 mg of P4N2 was suspended in 712 μL of CD3CN, and then 38 μL of a 2 M stock solution of HOTf in CD3CN was added to give final concentrations of 10 mM ligand and 0.1 M HOTf. A similar procedure was followed to make solutions containing 0.5 and 1 M HOTf. The solutions were transferred to a NMR tube and analyzed by 31P{1H} NMR after 1 h and 3 days. Over a period of 3 days, the solutions containing 0.5 and 1 M HOTf changed color from light yellow to light red. X-ray Diffraction Analysis. Crystals suitable for X-ray diffraction studies were grown by vapor diffusion of diethyl ether into an acetonitrile solution of 1 or 4. Crystals were mounted onto a nylon fiber loop with NVH immersion oil and then cooled to the data collection temperature of 100(2) K. Data collection was performed using a Bruker-AXS Kappa APEX II CCD diffractometer with 0.71073 Å Mo Kα radiation. Unit-cell parameters were obtained from 90 data frames at 0.3° Φ from three different sections of the Ewald sphere. Data were integrated and corrected for absorption effects with the

[Ni(P4N2)(CH3CN)](BF4)2 (1). A solution of [Ni(CH3CN)6][BF4]2 (258 mg, 0.54 mmol, 1 equiv) in acetonitrile was added to a stirred suspension of P4N2 (392 mg, 0.54 mmol, 1 equiv) in acetonitrile (20 mL) over the course of 5 min. The mixture was stirred for 1 h 30 min, and then the solvent was removed under vacuum. The resulting residue was triturated with hexane (5 mL) until a free-flowing powder formed, and then the solvent was removed and the solid dried under vacuum. Yield: 494 mg (0.49 mmol, 91%). 1H NMR analysis of 1 in CD3CN revealed the presence of 0.6 equiv of free CH3CN, indicating that the coordinated CH3CN ligand is weakly bound and labile under vacuum. Complex 1 was observed to be pure by NMR spectroscopy but consistently failed to pass elemental analysis within 0.4%. Similar results were observed for cobalt analogues of the same P4N2 ligand.29 Although these results are outside the range viewed as establishing analytical purity, they are provided to illustrate the best values obtained to date. Anal. Calcd for C44H46B2NiF8N2P4·0.6C2H3N: C, 55.18; H, 4.90; N, 3.69. Found: C, 54.49; H, 4.92; N, 3.16. 1H NMR (500 MHz, CD3CN): δ 7.48 (t, 4H, 3JHH = 7.4 Hz, ArH), 7.43 (t, 4H, 3 JHH = 7.9 Hz, ArH), 7.19 (m2, 12H, ArH), 7.13 (t, 2H, 3JHH = 7.3 Hz, ArH), 6.99 (m, 4H, ArH), 4.18 (dd, 4H, 2JHH = 14.2 Hz, 2JHP = 1.0 Hz, PCH2N), 4.02 (d, 4H, 2JHH = 14.2 Hz, PCH2N), 2.87 (m, 4H, P(CH2)2P), 2.59 (m, 8H, P(CH2)2P), 1.96 (s, 1.8H, CH3CN). 31 1 P{ H} NMR (CD3CN): δA 63.67, δB 62.26; 2JAA′ = 85.3 Hz, 2JAB = 28.9 Hz, 2JAB′ = −162.1 Hz, 2JBB′ = 34.2 Hz. 13C{1H} NMR (CD3CN): δ 151.5 (m), 133.7 (m), 133.1 (s), 130.8 (s), 130.2 (m), 128.5 (m), 123.8 (s), 119.8 (s), 49.7 (m), 31.3 (m), 24.0 (m). [Ni(P4N2-Me)](BF4)3 (4). A solution of 1 (100 mg, 0.10 mmol, 1.0 equiv) in CH2Cl2 (4 mL) was added to a suspension of [(CH3)3O]BF4 (85 mg, 0.57 mmol, 5.7 equiv) in CH2Cl2 (5 mL). After the mixture was stirred for 3 days, the yellow precipitate was collected on a bed of Celite via filtration and was rinsed with CH2Cl2 (5 mL). The product was then rinsed through the Celite with acetonitrile until the filtrate ran colorless, and then the filtrate was concentrated to dryness. The residue was crystallized by vapor diffusion of Et2O into CH3CN to afford yellow needles, which were dried under vacuum. 1H NMR analysis indicates the presence of 2 equiv of free CH3CN, presumably resulting from noncoordinated CH3CN in the crystal lattice. Yield 4·2CH3CN: 74 mg (0.065 mmol, 65%). Trace amounts of unidentified impurities were observed in the 1 H NMR spectrum, indicating that the isolated product is not analytically pure. 1H NMR (500 MHz, CD3CN): δ 7.69−7.59 (m, 5H, ArH), 7.52 (t, 4H, 3JHH = 7.8 Hz, ArH), 7.46 (d, 2H, 3JHH = 8.0 Hz, ArH), 7.41−7.34 (m, 10H, ArH), 7.24 (t, 1H, 3JHH = 7.3 Hz, ArH), 7.3 (dd, 4H, 3JHP = 11.0 Hz, 3JHH = 8.0 Hz, ArH), 6.49 (dd, 4H, 3 JHP = 10.6 Hz, 3JHH = 8.0 Hz, ArH), 5.53 (d, 2H, 2JHH = 15.4 Hz, PCH2N), 4.88 (m, 2H, PCH2N), 4.65 (dd, 2H, 2JHH = 2JHP = 13.9 Hz, PCH2N), 4.06 (m, 2H, PCH2N), 3.80 (t, 3H, 4JHP = 2.5 Hz, NCH3), 3.15 (m, 4H, P(CH2)2P), 2.87 (m, 2H, P(CH2)2P), 2.71 (m, 2H, P(CH2)2P). 31P{1H} NMR (CD3CN): δA 59.51, δB 57.46; 2JAA′ = 85.0 Hz, 2JAB = 27.3 Hz, 2JAB′ = −170.8 Hz, 2JBB′ = 32.4 Hz. 13C{1H} NMR (CD3CN): δ 150.8 (t), 145.0 (s), 134.7 (d), 134.6 (s), 134.4 (d), 134.1 (s), 131.1 (m), 130.8 (m), 130.7 (s), 126.2 (m), 125.6 (s), 124.4 (m), 66.6 (s), 61.4 (m), 49.8 (m), 31.1 (m), 23.3 (m). Stoichiometric Electroreduction. Controlled-potential electrolysis was performed on 8 mL of a benzonitrile solution containing 5 mM 1, 0.2 M [Bu4N][BF4], and a small amount of ferrocene as an internal potential reference. Electrolysis was conducted at −1.81 V until 7.84 C of charge had passed, corresponding to 2.1 equiv of electrons. An aliquot of the analyte solution was removed and analyzed by 31P{1H} NMR spectroscopy. Catalytic Hydrogen Production. A 2.0 mL acetonitrile solution of 0.3 mM 1, 0.2 M [Bu4N][PF6], and a small amount of ferrocene was prepared, and a baseline voltammogram was recorded at υ = 15 V s−1. A separate solution was prepared containing 1.5 M [DMF(H)][OTf] and 0.2 M [Bu4N][PF6]. A total of 0.98 mL of the acid solution was added to the catalyst solution in 30−250 μL increments. After each addition of acid, the working electrode was polished with diamond paste (Buehler, 0.25 μm) on a polishing pad wetted with H2O and rinsed with acetonitrile, and then a cyclic voltammogram H


Article

Organometallics APEX2 software package.79 All non-hydrogen atoms were refined with anisotropic displacement parameters, and all hydrogen atoms were treated as idealized contributions. Complex 1. The systematic absences in the data were consistent with the monoclinic space group P21/n. The asymmetric unit contains one [Ni(P4N2)(CH3CN)] 2+ cation, two [BF4]− anions, and acetonitrile solvent and diethyl ether solvent. One of the [BF4]− anions is disordered over two positions, which were located from the difference map and refined using SIMU, DELU, and SAME commands. There is one fully occupied acetonitrile solvent site and one site that has mixed occupancy with acetonitrile and diethyl ether. The molecules in the mixed site were located from the difference map and refined using SIMU, DELU, SAME, and DFIX commands to keep the refinement stable. Complex 4. The systematic absences in the data were consistent with the monoclinic space group P21/c. The asymmetric unit contains one [Ni(P4N2-Me)]2+ cation, three [BF4]− anions, and three molecules of acetonitrile solvent. One of the [BF4]− anions is disordered over two positions, which were located from the difference map and refined using SIMU, DELU, and SAME commands.

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Generation Under Aqueous Conditions. Chem. Soc. Rev. 2013, 42, 2388−2400. (3) McKone, J. R.; Marinescu, S. C.; Brunschwig, B. S.; Winkler, J. R.; Gray, H. B. Earth-abundant Hydrogen Evolution Electrocatalysts. Chem. Sci. 2014, 5, 865−878. (4) Fukuzumi, S.; Lee, Y.-M.; Nam, W. Thermal and Photocatalytic Production of Hydrogen with Earth-abundant Metal Complexes. Coord. Chem. Rev. 2018, 355, 54−73. (5) Kaeffer, N.; Chavarot-Kerlidou, M.; Artero, V. Hydrogen Evolution Catalyzed by Cobalt Diimine−Dioxime Complexes. Acc. Chem. Res. 2015, 48, 1286−1295. (6) McCrory, C. C. L.; Uyeda, C.; Peters, J. C. Electrocatalytic Hydrogen Evolution in Acidic Water with Molecular Cobalt Tetraazamacrocycles. J. Am. Chem. Soc. 2012, 134, 3164−3170. (7) Schnidrig, S.; Bachmann, C.; Muller, P.; Weder, N.; Spingler, B.; Joliat-Wick, E.; Mosberger, M.; Windisch, J.; Alberto, R.; Probst, B. Structure-Activity and Stability Relationships for Cobalt PolypyridylBased Hydrogen-Evolving Catalysts in Water. ChemSusChem 2017, 10, 4570−4580. (8) DuBois, D. L. Development of Molecular Electrocatalysts for Energy Storage. Inorg. Chem. 2014, 53, 3935−3960. (9) Luo, S.; Siegler, M. A.; Bouwman, E. Dinuclear Nickel Complexes of Thiolate-Functionalized Carbene Ligands and Their Electrochemical Properties. Organometallics 2018, 37, 740−747. (10) Virca, C. N.; Lohmolder, J. R.; Tsang, J. B.; Davis, M. M.; McCormick, T. M. Effect of Ligand Modification on the Mechanism of Electrocatalytic Hydrogen Production by Ni(pyridinethiolate)3− Derivatives. J. Phys. Chem. A 2018, 122, 3057−3065. (11) Schilter, D.; Camara, J. M.; Huynh, M. T.; Hammes-Schiffer, S.; Rauchfuss, T. B. Hydrogenase Enzymes and Their Synthetic Models: The Role of Metal Hydrides. Chem. Rev. 2016, 116, 8693−8749. (12) Seo, J.; Manes, T. A.; Rose, M. J. Structural and Functional Synthetic Model of Mono-iron Hydrogenase Featuring an Anthracene Scaffold. Nat. Chem. 2017, 9, 552−557. (13) Connor, G. P.; Mayer, K. J.; Tribble, C. S.; McNamara, W. R. Hydrogen Evolution Catalyzed by an Iron Polypyridyl Complex in Aqueous Solutions. Inorg. Chem. 2014, 53, 5408−5410. (14) Zee, D. Z.; Chantarojsiri, T.; Long, J. R.; Chang, C. J. Metal− Polypyridyl Catalysts for Electro- and Photochemical Reduction of Water to Hydrogen. Acc. Chem. Res. 2015, 48, 2027−2036. (15) Appel, A. M.; DuBois, D. L.; Rakowski DuBois, M. Molybdenum-Sulfur Dimers as Electrocatalysts for the Production of Hydrogen at Low Overpotentials. J. Am. Chem. Soc. 2005, 127, 12717−12726. (16) Fang, M.; Engelhard, M. H.; Zhu, Z.; Helm, M. L.; Roberts, J. A. S. Electrodeposition from Acidic Solutions of Nickel Bis(benzenedithiolate) Produces a Hydrogen-Evolving Ni−S Film on Glassy Carbon. ACS Catal. 2014, 4, 90−98. (17) Kaeffer, N.; Morozan, A.; Fize, J.; Martinez, E.; Guétaz, L.; Artero, V. The Dark Side of Molecular Catalysis: Diimine-Dioxime Cobalt Complexes are not the Actual Hydrogen Evolution Electrocatalyst in Acidic Aqueous Solutions. ACS Catal. 2016, 6, 3727−3737. (18) Waldie, K. M.; Kim, S.-K.; Ingram, A. J.; Waymouth, R. M. Cyclopentadienyl Cobalt Complexes as Precatalysts for Electrocatalytic Hydrogen Evolution. Eur. J. Inorg. Chem. 2017, 2017, 2755− 2761. (19) Anxolabéhère-Mallart, E.; Costentin, C.; Fournier, M.; Robert, M. Cobalt-Bisglyoximato Diphenyl Complex as a Precatalyst for Electrocatalytic H2 Evolution. J. Phys. Chem. C 2014, 118, 13377− 13381. (20) Cherdo, S.; Ghachtouli, S. E.; Sircoglou, M.; Brisset, F.; Orio, M.; Aukauloo, A. A Nickel Dimethyl Glyoximato Complex to Form Nickel Based Nanoparticles for Electrocatalytic H2 Production. Chem. Commun. 2014, 50, 13514−13516. (21) Tsay, C.; Yang, J. Y. Electrocatalytic Hydrogen Evolution under Acidic Aqueous Conditions and Mechanistic Studies of a Highly Stable Molecular Catalyst. J. Am. Chem. Soc. 2016, 138, 14174− 14177.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.organomet.8b00548. NMR spectra and electrochemical data (PDF) Accession Codes

CCDC 1859499−1859500 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing [email protected], or by contacting The Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

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AUTHOR INFORMATION

Corresponding Author

*E-mail for E.S.W.: [email protected]. ORCID

William G. Dougherty: 0000-0001-7016-384X W. Scott Kassel: 0000-0002-6764-9045 Eric S. Wiedner: 0000-0002-7202-9676 Present Address

§ W.G.D.: Department of Chemistry, Susquehanna University, 514 University Ave., Selinsgrove, PA 17870, United States.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS 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 (DOE), Office of Science, Office of Basic Energy Sciences. Pacific Northwest National Laboratory is operated by Battelle for DOE.

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REFERENCES

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Electrocatalytic Hydrogen Production by a Nickel Complex Containing

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