Impact of Weak Agostic Interactions in Nickel Electrocatalysts for

Jun 13, 2017 - To understand how H2 binding and oxidation is influenced by [Ni(PR2NR′2)2]2+ catalysts with H2 binding energies close to thermoneutra...
0 downloads 0 Views 1MB Size
Download PDF
Article pubs.acs.org/Organometallics

Impact of Weak Agostic Interactions in Nickel Electrocatalysts for Hydrogen Oxidation Christina M. Klug, Molly O’Hagan, R. Morris Bullock, Aaron M. Appel, and Eric S. Wiedner* Center for Molecular Electrocatalysis, Pacific Northwest National Laboratory, P.O. Box 999, K2-57, Richland, Washington 99352, United States S Supporting Information *

ABSTRACT: To understand how H2 binding and oxidation is influenced by [Ni(PR2NR′2)2]2+ catalysts with H2 binding energies close to thermoneutral, two [Ni(PPh2NR′2)2]2+ (R = Me, C14H29) complexes with phenyl substituents on phosphorus and varying alkyl chain lengths on the pendant amine were studied. In the solid state, [Ni(PPh2NMe2)2]2+ exhibits a weak agostic interaction between the Ni(II) center and the a C−H bond of the pendant N−CH3 group. DFT computations and variable-temperature 31P NMR experiments suggest that the agostic interaction persists in solution. The equilibrium constants for H2 addition to these complexes were measured by 31P NMR spectroscopy, affording free energies of H2 addition (ΔG°H2) of −0.8 kcal mol−1 in benzonitrile and −1.7 to −2.7 kcal mol−1 in THF. The agostic interaction contributes to the low driving force for H2 binding by stabilizing the four-coordinate Ni(II) species prior to binding of H2. The pseudo-first-order rate constants for addition of H2 (1 atm) were measured by variable-scan rate cyclic voltammetry and were found to be similar for both complexes, with rate constants of 3−6 s−1 in THF and less than 0.2 s−1 in benzonitrile. In the presence of exogenous base and H2, turnover frequencies of electrocatalytic H2 oxidation were measured to be less than 0.2 s−1 in benzonitrile and 4−6 s−1 in THF. These complexes are slower electrocatalysts for H2 oxidation in comparison to previously studied [Ni(PR2NR′2)2]2+ complexes because of a competition between H2 binding and formation of the agostic bond. However, the decrease in catalytic rate is accompanied by a beneficial 130 mV decrease in overpotential.

■

INTRODUCTION [FeFe]- and [NiFe]-hydrogenases are well-known enzymes that reversibly catalyze the interconversion of H2 and protons.1−3 Many groups have studied synthetic models that mimic the structure of the enzyme active site,4−8 but only a few of these complexes are catalysts for the oxidation of H2.9−12 Our group has studied electrocatalytic H2 oxidation using complexes of Ni, Fe, and Mn13,14 containing PR2NR′2 ligands (PR2NR′2 = 1,5-R3,7-R′-1,5-diaza-3,7-diphosphacyclooctane) that are functional mimics of the azadithiolate ligand found in the active site of [FeFe]-hydrogenase.2,15,16 In particular, [Ni(PR2NR′2)2]2+ complexes can be tuned to catalyze either H2 oxidation or proton reduction by controlling the free energy for H2 addition (ΔG°H2) through ligand modifications. Catalysts of this type for H2 oxidation typically contain bulky cyclohexyl groups on phosphorus and alkyl groups on the pendant amines. Nickel complexes containing these substituents have a strong driving force for H2 addition (ΔG°H2 < −3 kcal mol−1), which contributes to a large catalytic overpotential (≥300 mV): i.e., wasted electrochemical energy. 17−20 Only a few [Ni(PR2NR′2)2]2+ complexes have been found to reversibly catalyze both the oxidation and production of H2.21−23 Reversible electrocatalysis, as observed in hydrogenase, requires H2 binding to be close to thermoneutral so that both H2 binding and release can occur at an appreciable rate. In this regard, we © XXXX American Chemical Society

are interested in developing a greater understanding of how the reaction profile for H2 binding and release for [Ni(PR2NR′2)2]2+ changes as ΔG°H2 approaches thermoneutrality. The mechanism for H2 binding to [Ni(PR2NR′2)2]2+ has been investigated previously through a combination of experimental and computational studies.20,24 Hydrogen binding likely proceeds through the formation of a high-energy Ni(II) dihydrogen complex that is stabilized through hydrogenbonding interactions between the dihydrogen ligand and the pendant amine groups (Figure 1). The [Ni(PR2NR′2)(H2)]2+ complex then undergoes rapid sequential intramolecular proton transfers from the nickel center to the pendant amines, generating a proton-hydride intermediate before evolving to the endo-endo isomer of a doubly protonated Ni(0) (Figure 1). At high concentrations of base, the rate-limiting step for catalysis is the addition of hydrogen to [Ni(PR2NR′2)2]2+.24−26 Catalysts containing simple alkyl groups on the pendant amine display electrocatalytic turnover frequencies (TOF) of ≤40 s−1 at 25 °C and 1 atm of H2.18,19,26,27 One effort toward improving rates of H2 oxidation incorporated amino acid residues onto the pendant amines.28 Incorporation of an arginine group on the pendant amine affords a catalyst with a TOF of 210 s−1 under Received: February 8, 2017

A

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

Figure 2. Structures of the Ni(0) complexes (left) and Ni(II) complexes (right) used in this study.

0.5 equiv of [Ni(MeCN)6 ](BF 4 ) 2 in acetonitrile, and subsequent reaction with H 2 and base (n-butylamine, nBuNH2) led to precipitation of the Ni(0) complex Ni(PPh2NC142)2 as a yellow crystalline solid (Figure 2).29 [Ni(PPh2NMe2)2](BF4)2 was synthesized by oxidation of Ni(PPh2NMe2)2 with ferrocenium tetrafluoroborate ([Cp2Fe]BF4) in THF. This procedure leads to clean precipitation of the desired Ni(II) product in 1 h, which is a significantly shorter reaction time in comparison to that required for the method previously reported.29 [Ni(PPh2NC142)2](BF4)2 was obtained as a red solid by oxidation of Ni(PPh2NC142)2 with 2 equiv of [Cp2Fe]BF4 in acetonitrile. Due to the limited solubility of [Ni(PPh2NMe2)2](BF4)2 in THF, [B(C6F5)4]− was used as the anion for H2 addition studies in THF. This compound was synthesized in a fashion similar to that for the BF4− salt, using [Cp2Fe][B(C6F5)4] as the chemical oxidant. Cyclic voltammograms of Ni(PPh2NC142)2 and Ni(PPh2NMe2)2 were collected in benzonitrile (PhCN) and THF (Figure S7 in the Supporting Information). Benzonitrile was chosen for these studies over acetonitrile due to the limited solubility of the complexes in acetonitrile. These compounds display two fully reversible one-electron waves assigned to the Ni(I/0) and Ni(II/I) couples in both PhCN and THF (Table 1). Plots of

Figure 1. Mechanism for H2 binding to Ni(PR2NR′2)22+ complexes. The substituents on P have been omitted for clarity.

ambient conditions, representing a substantial improvement in catalytic activity in comparison to catalysts without amino acid residues. Computational studies of the amino acid modified catalysts suggest that intramolecular interactions between amino acid groups on opposing PR2NR′2 ligands slows the chair−boat isomerization of the pendant amines. This interaction is proposed to lead to improved positioning of the amines relative to the Ni(II) center, stabilizing the interaction between the metal and an incoming H2 molecule.23 This study was designed to gain an understanding of two different design aspects related to H2 binding in [Ni(PR2NR′2)2]2+ complexes and their subsequent impact on electrocatalytic H2 oxidation. First, we desired to understand how the reaction profile for H2 binding changes when the overall driving force is close to thermoneutral. To this end, we performed stoichiometric and electrocatalytic H2 oxidation studies on [Ni(PPh2NMe2)2]2+, which has a phenyl group on phosphorus and a methyl group on nitrogen. Hydrogen binding to this complex has been shown to be slightly downhill in acetonitrile (ΔG°H2 = −1.1 kcal mol−1),29 but the electrocatalytic activity for H2 oxidation has not been previously reported. Second, we sought to learn whether H2 binding can be improved by controlling the structural dynamics affecting the positioning of the pendant amines, and therefore the interaction between the pendant amines and H2. For this objective, we prepared the new complex [Ni(PPh2NC142)2]2+ (C14 = C14H29), where the tetradecyl chain on nitrogen is expected to slow the structural dynamics of the ligand in a manner similar to that for the recently reported [Ni(PR2NR′2)2]2+ electrocatalysts for production of H2.30 Unexpectedly, the complexes investigated in this study display evidence of a weak agostic bond between the Ni(II) center and the α-CH of the pendant amine substituent. We hypothesize that this weak agostic interaction competes with H2 binding, thereby slowing the rate of electrocatalysis relative to other catalysts that do not form a similar agostic interaction.

Table 1. Electrochemical Data for Ni(PPh2NR′2)2 Complexesa E1/2b (ΔEpc) complex

solvent

Ni(PPh2NC142)2

THFd PhCNe THFd PhCNe

Ni(PPh2NMe2)2

Ni(II/I) −0.82 −0.97 −0.84 −0.96

(70) (64) (79) (69)

Ni(I/0) −1.31 −1.27 −1.27 −1.21

(68) (68) (75) (66)

Scan rate 0.1 V s−1. bUnits of V versus Cp2Fe+/0. cUnits of mV. d0.1 M NBu4[B(C6F5)4]. e0.2 M NBu4BF4.

a

the peak current (ip) versus the square root of the scan rate are linear for both the Ni(II/I) and Ni(I/0) waves in both solvents, indicating that the redox events are diffusion-controlled (Figure S8 in the Supporting Information).31 The redox properties are very similar between the two complexes; the potentials of the Ni(I/0) couple for Ni(PPh2NC142)2 are slightly more negative than for Ni(PPh2NMe2)2 due to the electronic effects from the secondary versus primary α-carbon on nitrogen. The Ni(II/I) couple of each complex is 120−150 mV more negative in benzonitrile than in tetrahydrofuran. Nitrile solvents are known to coordinate weakly to [Ni(diphosphine)2]2+ complexes;32−35 therefore, the negative shift of the Ni(II/I) couple in PhCN is due in part to stabilization of the Ni(II) oxidation state through reversible formation of [Ni(PPh2NMe2)2(PhCN)]2+. Additionally, anion pairing between Ni(II) and the electrolyte will be

■

RESULTS Synthesis and Characterization. The synthesis of Ni(PPh2NMe2)2, containing a phenyl group at phosphorus and a methyl group at nitrogen (Figure 2), has been described previously.29 A new ligand containing a tetradecyl group at nitrogen, PPh2NC142, was synthesized by condensing bis(hydroxymethylphenylphosphine) and 1-tetradecylamine in ethanol at 70 °C. Attempts to purify the PPh2NC142 ligand were unsuccessful. The crude ligand mixture was treated with B

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

optimized structure and the solid-state structure was obtained using the B3P86 hybrid functional,40,41 the Stuttgart basis set with effective core potential (ECP) for Ni,42 and the 6-31G** basis set on the H, C, N, and P atoms.43,44 With these parameters, a CH···Ni distance of 2.442 Å and a C···Ni distance of 3.313 Å was calculated for the weak bonding interaction, in good agreement with the distances observed in the solid state. The C−H bond oriented toward Ni is elongated by 0.024 Å relative to the remaining two C−H bonds in the methyl group that points away from the Ni center, corroborating the presence of a weak CH···Ni bonding interaction. The electronic structure of [Ni(PPh2NMe2)2]2+ was investigated in more detail to better understand the nature of the CH···Ni interaction.45−47 Natural bond orbital (NBO) analysis48 is well suited to this task,49,50 since it partitions the computed electron density into Lewis-type valence orbitals, and second-order perturbative theory analysis within the NBO framework provides information on the energetic stabilization resulting from electron delocalization between an occupied donor orbital and a nearby vacant acceptor orbital. Two weakly stabilizing interactions were identified between the C−H bond and Ni using NBO analysis (Figure 4). The primary interaction

greater in THF than in PhCN, which also contributes to the difference in Ni(II/I) couples in these solvents.26,36,37 The solid-state structure of [Ni(PPh2NMe2)2]2+, determined by single-crystal X-ray diffraction, was found to depend on the solvent choice for crystal growth. Crystallization of [Ni(PPh2NMe2)2](BF4)2 from CH3CN/Et2O resulted in a trigonal-bipyramidal structure with each PPh2NMe2 ligand occupying one axial and one equatorial position and a CH3CN ligand occupying the third equatorial position (Figure 3a). Similar

Figure 4. Stabilizing CH···Ni interactions in [Ni(PPh2NMe2)2]2+ that were identified by NBO analysis.

is an agostic bond between the positioned σ(C−H) orbital of the N-methyl group and the vacant 4pz orbital on Ni (Figure 4a). A donor−acceptor stabilization energy of 6.4 kcal mol−1 was calculated for this agostic interaction, indicating that it is a very weak bond in comparison to conventional agostic bonds that have stabilizing energies of 20−60 kcal mol−1 according to NBO perturbative analysis.50,51 The second interaction revealed by NBO analysis is a hydrogen bond between the filled 3dz2 orbital on Ni and the vacant σ*(C−H) orbital (Figure 4b). With a stabilizing energy of only 1.3 kcal mol−1, the hydrogen bond contributes less to stabilization of the CH···Ni interaction than does the agostic bond. Therefore, the CH···Ni interaction will be referred to as an agostic bond even though it is not the only stabilizing interaction that is present. A pertinent consideration is whether the weak agostic bond persists in solution or is simply an artifact of the solid-state crystal packing. NMR spectroscopic evidence for the agostic bond in [Ni(PPh2NMe2)2]2+ could not be obtained at −90 °C due to the dynamic nature of the PR2NR′2 ligands (Figures S9− S13 in the Supporting Information); therefore, its stability was investigated through DFT analysis. Thermodynamic calculations were performed using the functional and basis set described above in conjunction with the polarizable continuum model of solvation in tetrahydrofuran.52,53 In this manner the isomer with the methyl group positioned for the agostic bond was calculated to be 1.3 kcal mol−1 more stable than the isomer with the methyl group pointed away from nickel, consistent

Figure 3. X-ray crystal structural depictions of (a) [Ni(PPh2NMe2)2(CH3CN)](BF4)2 and (b) [Ni(PPh2NMe2)2][B(C6F5)4]2. The anions, phenyl rings, and most hydrogen atoms are omitted for clarity. Thermal ellipsoids are shown with 40% probability.

five-coordinate structures have been observed for other [Ni(PR2NR′2)2(CH3CN)]2+ complexes.33−35,38 When crystals of [Ni(PPh2NMe2)2][B(C6F5)4]2 were grown from THF/Et2O, the complex was found to adopt a distorted-square-planar coordination geometry (Figure 3b) with no coordinated solvent molecule. Each pendant amine forms a six-membered ring with Ni, with two rings in a chair conformation and two in a boat conformation. The N-methyl group of one of the rings assuming the boat conformation is folded over the Ni(II) center, leading to a short CH···Ni distance of ∼2.4 Å, a C···Ni distance of 3.260(5) Å, and a C−H−Ni angle of 142°, all of which suggest the presence of a weak bonding interaction. This CH···Ni distance is longer than typical for agostic bonds (1.8− 2.3 Å) but is within the range of 2.3−2.9 Å noted for weak bonding interactions in square-planar complexes of d 8 transition-metal centers.39 Density functional theory (DFT) calculations were performed on [Ni(PPh2NMe2)2]2+ to further characterize the weak bonding interaction. Good agreement between the geometryC

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics with the agostic bond being very weak. Breaking the agostic interaction requires inversion of the alkyl substituents on the pendant amine, and the calculated barrier for this process (6.3 kcal mol−1 in THF solvent) is only slightly higher than the barrier for inversion of trimethylamine (6.0 kcal mol−1 in the gas phase).54 A similar agostic interaction is likely present in [Ni(PPh2NC142)2]2+, though calculations were not performed on this complex because of the large number of conformational isomers that result from the tetradecyl chains. Structural Dynamics. The kinetics of boat−chair isomerization of the PR2NR′2 ligand in the Ni(0) and Ni(II) states of each complex were studied to examine the impact on electrocatalytic H 2 oxidation. 30,55 Variable-temperature 31 1 P{ H} NMR spectra of the Ni(0) complexes Ni(PPh2NMe2)2 and Ni(PPh2NC142)2 were collected in THF between 25 and −90 °C (Figures S14 and S15 in the Supporting Information). At 25 °C, Ni(PPh2NMe2)2 and Ni(PPh2NC142)2 each display a single sharp resonance in the 31P{1H} NMR spectrum. Cooling Ni(PPh2NC142)2 to −90 °C leads to partial decoalescence of the 31

Figure 5. Simplified mechanism for chair−boat isomerization in [Ni(PR2NR′2)2(CH3CN)]2+.

PR2NR′2

P resonances, consistent with each ligand having one arm in a chair conformation and one in a boat conformation. A single broad resonance was observed for Ni(PPh2NMe2)2 even at −90 °C, indicating that the 31P nuclei are still undergoing fast exchange due to rapid boat−chair interconversion. These results indicate that the longer alkyl chain of Ni(PPh2NC142)2 slows the rate of boat−chair isomerization relative to Ni(PPh2NMe2)2, consistent with a recent study on [Ni(PR2NR′2)2]2+ electrocatalysts for proton reduction.30 In the Ni(II) complexes, the rates of boat−chair isomerization were measured in 4/1 CH2Cl2/CH3CN, a solvent mixture that was previously used to probe the structural dynamics of other [Ni(PR2NR′2)2(CH3CN)]2+ complexes.30,55 At room temperature, [Ni(PPh2NMe2)2(CH3CN)]2+ and [Ni(PPh2NC142)2(CH3CN)]2+ each display a broad resonance in the 31 1 P{ H} NMR spectrum. When the temperature is lowered to −30 °C, the single resonance decoalesces into two (Figures S16 and S17 in the Supporting Information), which are attributed to the two inequivalent 31P environments for the axial and equatorial positions of the trigonal-bipyramidal geometry of the Ni(II) center with coordinated CH3CN. Previous reports have demonstrated that the exchange between the axial and equatorial positions for the 31P nuclei is accompanied by isomerization of the boat and chair conformers of the sixmembered rings of the ligand (Figure 5).55 The activation parameters for this process were determined by line-shape analysis using a two-site exchange model. Eyring plots of the two compounds show that the rates of 31P exchange at 25 °C are similar (Table 2 and Figure S18 in the Supporting Information), with the methyl substituent appearing to have a slightly slower isomerization rate in comparison to the tetradecyl group.55 The entropy of activation was determined to be slightly negative for both complexes, suggesting that dissociation of the weakly coordinated CH3CN ligand is not the rate-limiting step in the exchange process. This interpretation is consistent with previous DFT calculations on a related complex, where the barriers for sequential ring inversion of the pendant amines were found to be higher than the barrier for dissociation of CH3CN.55 Therefore, the isomerization rates for [Ni(PPh2NMe2)2(CH3CN)]2+ and [Ni(PPh2NC142)2(CH3CN)]2+ are determined by the barriers for ring inversion of the fourcoordinate intermediate. The isomerization barrier for these

Figure 6. Equilibrium between Ni(II) and the doubly protonated Ni(0) isomers under 1 atm of H2 in either PhCN or THF solution.

Table 2. Boat−Chair Isomerization Rate Constants and Activation Parameters for 31P Exchange of [Ni(PPh2NR′2)2]2+ in 4/1 CH2Cl2/MeCN complex [Ni(PPh2NC142)2]2+ [Ni(PPh2NMe2)2]2+

k298 K (s −1)

ΔH⧧ (kcal mol−1)

ΔS⧧ (cal mol−1 K−1)

ΔG⧧298 K (kcal mol−1)

9.7 × 104 5.2 × 104

9.5 ± 1.5

−3.9 ± 0.4

10.7

9.5 ± 1.8

−5.2 ± 0.4

11.0

complexes is 0.7−1.0 kcal mol−1 higher than for [Ni(PPh2NtBu2)2(CH3CN)]2+ (ΔG⧧298 K = 10.0 kcal mol−1),55 a closely related complex that does not possess an α-CH capable of forming an agostic bond. This suggests that breaking the agostic bond in [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ contributes to their rates of isomerization, which are similar despite the difference in chain length of the substituents on the pendant amines. H2 Addition Studies. The Ni(II) complexes react with H2 (1 atm) at room temperature in benzonitrile or THF to form an equilibrium mixture of unreacted Ni(II) and the doubly protonated Ni(0) complex with a distribution of three structural isomers which vary in the position of the protons relative to the metal center: endo-endo (ee), endo-exo (ex), and exo-exo (xx) (Figure 6). These isomers were identified by their characteristic resonances in the 31P{1H} NMR spectrum (see the Supporting Information for spectral assignments).18,19,27,56,57 Of these three isomers, the ee isomer is D

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics Table 3. Equilibrium Constants and Pseudo-First-Order Rate Constants for H2 Addition THF

a

PhCN

complex

KH2,ee (atm−1)

ΔG° (kcal mol−1)

kH2a (s−1)

KH2,ee (atm−1)

ΔG° (kcal mol−1)

kH2a (s−1)

[Ni(PPh2NC142)2]2+ [Ni(PPh2NMe2)2]2+

96 ± 6 18 ± 3

−2.7 −1.7

6 3

3±2 4±2

−0.7 −0.8

0.2 0.1

PH2 = 1 atm.

Figure 7. (a) Cyclic voltammograms of Ni(PPh2NMe2)2 under 1 atm of H2 in benzonitrile with 0.2 M NBu4BF4 at 0.05 V s−1 (blue) and 0.5 V s−1 (red). [Cp*2Co]PF6 was added as an internal redox reference. (b) Plot of the natural log of the ratio of the cathodic peak current for the Ni(II/I) couple to the anodic peak current for the Ni(I/0) couple as a function of reaction time under 1 atm of H2.

Figure 8. Cyclic voltammograms of Ni(PPh2NC142)2 in absence of base (black trace) and varying concentrations of n-BuNH2 in (a) benzonitrile with 0.2 M [NBu4][BF4] and (b) THF with 0.1 M [NBu4][B(C6F5)4]. Scan rate υ = 0.05 V s−1. [Cp*2Co]PF6 was added as an internal redox reference.

the kinetic product of H2 addition18 and the ex and xx isomers are only formed by intermolecular isomerization with an exogenous base such as water.58−60 The equilibrium constant of H2 addition resulting in the formation of the ee isomer (KH2,ee) was determined by comparing the relative integration of the ee isomer in the 31P{1H} NMR spectrum versus the Ni(II) resonances; the resulting KH2,ee values are presented in Table 3. Hydrogen addition is more favorable in THF than in benzonitrile, presumably due to differences in solvent binding to Ni(II). The kinetics of H2 addition to Ni(II) were measured by recording cyclic voltammograms of Ni(0) at various scan rates under 1 atm of H2.61 Since the nickel complexes do not react with H2 until they are oxidized to the Ni(II) state, the reaction can only take place after Ni(I) is oxidized to Ni(II) on the anodic scan and before Ni(II) is reduced to Ni(I) on the cathodic return scan. As a result, the time scale of the experiment is defined by the scan rate and the magnitude of the voltage sweep between oxidation of Ni(I) and reduction of

Ni(II). At sufficiently slow scan rates (1 V s−1 in THF), H2 binding is kinetically competitive with reduction of Ni(II) to Ni(I), leading to an increase in the cathodic current for the Ni(II) reduction and decreased anodic current for the H2added products (Figure 7a, red trace). The anodic peak current (ipa) of the Ni(I/0) couple is proportional to the initial concentration of Ni before the reaction with H2, and the cathodic peak current (ipc) of the Ni(II/I) couple on the return scan is proportional to the concentration of Ni(II) remaining in the electrochemical diffusion layer that has not reacted with H2. Therefore, the pseudo-first-order rate constant for H2 addition E

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

Table 4. Observed Rate Constants (kobs) and Ecat/2 Values for Electrocatalytic H2 Oxidation under Different Solvent and Base Conditionsa Ni(PPh2NC142)2

a

−1

Ni(PPh2NMe2)2 −1

solvent, base

kobs (s )

Ecat/2 (V)

η (mV)

kobs (s )

Ecat/2 (V)

η (mV)

PhCN, nBuNH2b THF, nBuNH2c THF, NEt3c

0.2 6 4

−0.88 −0.81 −0.80

220 430 370

0.1 4 5

−0.90 −0.82 −0.81

200 420 360

Scan rate at 0.05 V s−1. Potentials are referenced versus Cp2Fe+/0. b0.2 M NBu4BF4 in PhCN. c0.1 M NBu4[B(C6F5)4] in THF.

binding of base to Ni(II) does not compete with H2 binding under electrocatalytic conditions. Overpotentials for H2 oxidation were determined from the difference between the catalytic half-wave potential (Ecat/2) and the thermodynamic potential of the H+/H2 couple for the base used in electrocatalysis and are given in Table 4. A thermodynamic potential of −1.10 V was estimated for nBuNH2 in benzonitrile using the standard state potential of a proton in acetonitrile (E°H+ = −0.028 V)65 and the pKa of nbutylammonium in acetonitrile (18.3).64 Using the open-circuit potential method described by Roberts and Bullock,65 the potential of the H+/H2 couple in THF was measured to be −1.24 V using 1/1 nBuNH2/nBuNH3+ and −1.17 V using 1/1 NEt3/HNEt3+.

to Ni(II) (kH2) was determined from plots of ln(ic/ip) versus the reaction time, where the slope is equal to −kH2 (Figure 7b and Figures S23 and S24 in the Supporting Information). A substantial solvent dependence is observed: kH2 values of approximately 0.1 and 0.2 s−1 were measured for [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ in PhCN, while kH2 increases to approximately 3 s−1 for [Ni(PPh2NMe2)2]2+ and 6 s−1 for [Ni(PPh2NC142)2]2+ in THF (Table 3). Electrocatalytic Oxidation of Dihydrogen. The two complexes Ni(PPh2NC142)2 and Ni(PPh2NMe2)2 were tested for electrocatalytic H2 oxidation. Using nBuNH2 as the base in the presence of 1 atm of H2 at a scan rate of 0.05 V s−1, the current for the Ni(II/I) oxidation wave increased slightly and displayed a “plateau” shape indicative of a steady-state catalytic response (Figure 8). The catalytic current (icat) was found to be independent of the nBuNH2 concentration at base concentrations above 10 mM for Ni(PPh2NMe2)2 and between 5 and 12 mM for Ni(PPh2NC142)2, while a slight decrease in icat was observed at base concentrations higher than 12 mM. Pseudofirst-order catalytic rate constants (kobs) were calculated in the base-concentration-independent regime using the ratio of icat to the peak current (ip) of the Ni(I/0) oxidation wave as previously described (Figures S27 and S28 in the Supporting Information)62,63 and are given in Table 4. The kobs values in THF were 40-fold larger in THF than in PhCN, presumably due to the decreased ability of THF to bind to Ni(II) in comparison to PhCN.26 The catalytic rate constants matched the pseudo-first-order rate constant for H2 addition measured by cyclic voltammetry within the experimental error; thus, the rate of H2 oxidation for these compounds appears to be limited by the rate of H2 addition. In the absence of H2, binding of nBuNH2 to Ni(II) was apparent at high concentrations of nBuNH2 (>20 mM), as indicated by the irreversibility of the Ni(II/I) couple (Figure S29 in the Supporting Information). To test the influence of base binding on kobs, H2 oxidation studies were also performed using triethylamine (NEt3) as the base, since it has a basicity similar to that of nBuNH2 but is more sterically bulky.64 Cyclic voltammograms recorded on the Ni(0) complexes and NEt3 in the absence of H2 show two fully reversible couples for the Ni(II/I) and Ni(I/0) redox events. In the presence of NEt3 and H2, a catalytic current enhancement is observed positive of the Ni(II/I) oxidation (Figures S25 and S26 in the Supporting Information), indicative of an alternative H2 oxidation mechanism in which the doubly protonated Ni(0) complex must be oxidized before deprotonation.18 On comparison of the two catalysts, the current enhancement at more negative potentials is greater for Ni(PPh2NMe2)2 than for Ni(PPh2NC142)2; thus, deprotonation of the doubly protonated Ni(0) with NEt3 is easier for Ni(PPh2NMe2)2, as the pendant amine is less sterically encumbered. Similar kobs values were measured for Ni(PPh2NMe2)2 using either nBuNH2 or NEt3, suggesting that

■

DISCUSSION

The positioning of the pendant amines in [Ni(PR2NR′2)2]2+ complexes plays a critical role in stabilizing a nickel dihydrogen complex, a high-energy intermediate of electrocatalytic H2 oxidation.20,24 One goal for this work was to determine whether the positioning of the pendant amines in [Ni(PR2NR′2)2]2+ complexes could be influenced indirectly by controlling the ligand structural dynamics. To our surprise, the structural dynamics of [Ni(P Ph 2 N Me 2 ) 2 ] 2+ and [Ni(PPh2NC142)2]2+ were the same within experimental error, despite the difference in chain lengths between the methyl and tetradecyl chains on the pendant amines. Structural analysis of [Ni(PPh2NMe2)2]2+ suggests the presence of a weak agostic interaction between the Ni(II) center and a methyl group from one of the pendant amines, with DFT analysis providing additional support for this interaction. A similar agostic interaction is presumed for [Ni(PPh2NC142)2]2+ between the α-CH of the alkyl chain and the Ni(II) center, since these compounds exhibit similar structural dynamics and catalytic activities. DFT calculations indicate the agostic interaction between the α-CH and the Ni(II) center stabilizes the fourcoordinate [Ni(PPh2NMe2)2]2+ species by 1.3 kcal mol−1, thereby inhibiting H2 binding by blocking the open site at Ni(II). Agostic interactions in [Ni(PR2NR′2)2]2+ complexes have not been previously studied; therefore, it is of interest to consider how prevalent these interactions may be in this class of compounds. In the solid-state structure of [Ni(PCy2N(CH2)2OMe2)2]2+, one of the pendant amine alkyl chains is in close proximity to the Ni(II) center with a Ni(II)−CH2 distance of 3.258(4) Å,19 suggesting the presence of an agostic interaction. In contrast, the solid-state structure of [Ni(PCy2NBn2)2]2+ shows a four-coordinate Ni(II) center with the α-CH2 groups not pointed toward the Ni,33 possibly due to prohibitive steric interactions between the benzyl group on nitrogen and the opposing bulky PCy2NBn2 ligand. When only F

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

Figure 9. Mechanism of H2 binding to [Ni(PPh2NMe2)2]2+ showing the influence of the agostic interaction. The phenyl groups on P have been omitted for clarity.

one bulky PR2NR′2 ligand is present, such as the case of [Ni(PtBu2NBn2)2(CH3CN)2]2+, the decreased steric constraint allows the α-CH of the N-benzyl group to form an agostic interaction with Ni.66 Additionally, a similar agostic interaction was observed in [Mn(PPh2NBn2)(CO)(bppm)]+ (bppm = (PArF2)2CH2, ArF = 3,5-(CF3)2C6H3), where the low electron density at five-coordinate Mn(I) overcomes the steric constraints imparted by bppm for bringing the pendant amine near the metal center.67 A second goal of this study was to flatten the free energy profile for H2 oxidation by using [Ni(PR2NR′2)2]2+ complexes with a free energy of H2 addition close to thermoneutral. This was accomplished by using phenyl substituents on the coordinating phosphorus atoms and basic pendant amines with alkyl substituents. 29 Hydrogen addition to [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ to give the doubly protonated endo-endo isomers is only slightly downhill in benzonitrile solution (ΔG°H2 = −0.8 kcal mol−1), while H2 addition to previously studied [Ni(PCy2NR′2)2]2+ catalysts for H2 oxidation is much more favorable (ΔG°H2 ≤ −3.2 kcal mol−1) due to the bulky cyclohexyl groups on phosphorus.20 The free energy of H2 addition to [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ is more favorable in THF (ΔG°H2 = −1.7 to −2.7 kcal mol−1) because of the absence of solvent binding to the Ni(II) center.26 The agostic interactions in [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ contribute to the free energy of H2 addition by stabilizing the four-coordinate Ni(II) complex. As a result, the Ni(II) complex must undergo two pre-equilibrium steps, breaking the agostic interaction and binding of H2, before the rate-limiting conversion of the high-energy Ni(H2) complex into the proton-hydride intermediate can occur (Figure 9).24 Consistent with this interpretation, the experimental barrier for H2 addition to [Ni(PPh2NMe2)2]2+ (ΔG⧧H2 = 16.6 kcal mol−1) is 1.3 kcal mol−1 higher than for [Ni(PCy2NBn2)2]2+ (ΔG⧧H2 = 15.3 kcal mol−1),26 a difference that is equal to the computed strength of the agostic interaction in [Ni(PPh2NMe2)2]2+ (1.3 kcal mol−1). In comparison to the previously studied H2 oxidation catalysts [Ni(PCy2NtBu2)2]2+ and [Ni(PCy2NBn2)2]2+, the Ni(II/ I) couples for [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ are 150−160 mV more negative in benzonitrile68 and 130−220 mV more negative in THF.26 This shift in potentials results from the smaller size of the phenyl substituents on phosphorus, which allows the Ni(II) complex to better adopt a squareplanar coordination geometry.69 As a consequence, the overpotentials of [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ (420 mV in THF) are much lower than the overpotentials of [Ni(PCy2NtBu2)2]2+ and [Ni(PCy2NBn2)2]2+ (540−560 mV in THF). Thus, while the agostic interaction negatively affects the turnover frequency for catalysis, the decreased steric profile of [Ni(PPh2NMe2)2]2+ and [Ni(PPh2NC142)2]2+ affords a favorable decrease in the overpotential.

■

SUMMARY

■

EXPERIMENTAL SECTION

The structural dynamics and H2 oxidation properties of two [Ni(PR2NR′2)2]2+ complexes incorporating phenyl substituents on phosphorus were examined. This study involved an expanded investigation of [Ni(PPh2NMe2)2]2+ and the synthesis and characterization of the related complex [Ni(PPh2NC142)2]2+, which differs in the length of the alkyl chain on the pendant amine. Despite the significant difference in length of the alkyl groups on the pendant amines, similar chair−boat isomerization rates were measured for the complexes in the Ni(II) oxidation state. The similarities between the two Ni(II) complexes are attributed to the formation of a weak agostic interaction between nickel and the α-CH of the alkyl chain, which was observed crystallographically for [Ni(PPh2NMe2)2]2+ and supported by DFT analysis. Addition of H2 to these complexes is favorable and close to thermoneutral in THF. These complexes were shown to be competent catalysts for electrochemical H2 oxidation with turnover frequencies ranging from 4 to 6 s−1 in THF using either nBuNH2 or NEt3 as the exogenous base. In comparison to previously studied [Ni(PR2NR′2)2]2+ complexes for H2 oxidation, the overpotential has been decreased by 130 mV in THF. Importantly, these studies highlight the level of control that is needed to create efficient mimics of natural enzymes. Hydrogenases have evolved to possess highly tailored active sites that effectively balance the energetic requirements for rapid and reversible cleavage of H2 into protons and electrons. It is difficult to design a synthetic catalyst that operates as efficiently as the enzyme, as many structural modifications can have unintended consequences on catalytic performance. In this work, creating a [Ni(PR2NR′2)2]2+ with a low free energy for H2 addition resulted in catalysts with lower overpotentials, though H2 binding is inhibited by formation of an agostic bond between the pendant amine and the less sterically protected Ni site. Successful creation of a molecular catalyst that matches the activity of the hydrogenases will require precise control over multiple aspects of the catalyst’s structure and function.

General Experimental Procedures. All manipulations were performed under a dinitrogen atmosphere using standard Schlenk techniques or in a Vacuum Atmospheres glovebox. All protio solvents were deoxygenated by sparging with nitrogen and dried by passage through neutral alumina columns in an Innovative Technology, Inc., PureSolv solvent purification system. Acetonitrile-d3 (MeCN-d3) was dried over P2O5 and purified by vacuum transfer. Tetrahydrofuran-d8 (THF-d8) was dried over NaK and vacuum-transferred. 1H, 31P{1H}, and 13C{1H} NMR spectra were recorded on a 500 MHz 1H frequency Agilent spectrometer. 31P{1H} NMR spectra were referenced to an external phosphoric acid standard. All 1H chemical shifts were referenced to the residual protio solvent, and all J values are given in Hz. MS analysis was performed using a 12 T Fourier transform ion cyclotron resonance mass spectrometer (FTICR-MS) (Bruker SolariX, Billerica, MA) outfitted with a standard electrospray G

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

−CH2CH3). 13C{1H} NMR (THF): 130.4, 130.1, 129.5, 127.4, 61.0, 31.8, 29.6, 29.5, 29.2, 27.2, 26.2,, 22.5, 13.4. ESI-MS (positive mode, MeCN): calcd for {[Ni(PPh2NC142)2Cl]}+, 1482.00 m/z; found, 1482.01 m/z. Formation of a chloride adduct is attributed to residual Cl− in the instrument. [Ni(PPh2NC142)2][B(C6F5)4]2. A solution of [Cp2Fe][B(C6F5)4] (60 mg, 0.069 mmol) in 4 mL of THF was added to a solution of Ni(PPh2NC142)2 (46 mg, 0.032 mmol) in 3 mL of THF. The reaction mixture was stirred for 1 h at room temperature, resulting in a color change of the solution from light yellow to orange. The resulting solution was filtered through a Celite plug, and the solvent was removed under reduced pressure. The resulting solid was triturated with hexanes (2 × 15 mL), the solvent was decanted, and the resulting tacky solid was dried under vacuum for 2 h (80 mg, 0.028 mmol, 90%). The spectral properties are consistent with those of the previously described complex with a BF4− anion. [Ni(PPh2NMe2)2](BF4)2. A solution of [Cp2Fe]BF4 (42 mg, 0.15 mmol) in 4 mL of THF was added to a solution of Ni(PPh2NMe2)2 (30 mg, 0.042 mmol) in 5 mL of THF. The resulting solution was stirred for 4 h at room temperature, resulting in the formation of a brick red precipitate. The solid was filtered, washed with THF (3 × 3 mL), and dried under vacuum for 4 h (33 mg, 0.037 mmol, 89%). The spectral properties are consistent with those previously reported.29 [Ni(PPh2NMe2)2][B(C6F5)4]2. A solution of [Cp2Fe][B(C6F5)4] (92 mg, 0.106 mmol) in 4 mL of THF was added to a solution of Ni(PPh2NMe2)2 (33 mg, 0.045 mmol) in 3 mL of THF. The reaction mixture was stirred for 1 h at room temperature, resulting in a color change of the solution from light yellow to red. The resulting solution was filtered through a Celite plug, and the solvent was removed under reduced pressure. The resulting solid was triturated with hexanes (15 mL), filtered, and washed with hexanes (3 × 15 mL) and diethyl ether (2 × 15 mL), yielding a free-flowing microcrystalline solid (84 mg, 0.040 mmol, 88%). The spectral properties are consistent with those previously reported.29 X-ray Structural Determination. Crystals of [Ni(PPh2NMe2)2][B(C6F5)4]2 were grown by removing THF under reduced pressure from [Ni(PPh2NMe2)2][B(C6F5)4]2, resulting in a tacky solid, to which diethyl ether was added. After 3 days at room temperature, crystals suitable for X-ray diffraction were obtained. Crystals of [Ni(PPh2NMe2)2(CH3CN)](BF4)2 were grown from diffusion of diethyl ether into an acetonitrile solution of the complex at room temperature. Single crystals were coated in Paratone-N oil under a dinitrogen atmosphere. The crystals were mounted on a Bruker Kappa Apex 2 CCD diffractometer on a Cryoloop under a stream of N2. All data were collected using a Mo Kα radiation source and a graphite monochromator. Data were integrated and corrected for absorption effects with the APEX3 software package.74 The structure was solved using direct methods and refined by least-squares methods using the SHELXTL software package.75 Unless otherwise noted, thermal parameters for all non-hydrogen atoms were refined anisotropically. Hydrogen atoms were added at the ideal positions and refined using a riding model where the thermal parameters were set at 1.2 times those of the attached carbon (1.5 times for methyl carbons). The disordered components of the cocrystallized diethyl ether molecules in the structure of [Ni(PPh2NMe2)2][B(C6F5)4]2 were modeled into two parts using the PART command and refined isotropically. The disordered components of the fluorine atoms of one tetrafluoroborate anion for [Ni(PPh2NMe2)2(CH3CN)](BF4)2 were modeled into two parts using the PART command and refined anisotropically. NMR Line Shape and Kinetic Analysis. The gNMR program as used to analyze the exchange process observed in the variabletemperature 31P{1H} spectra.76 The spectra were processed in gNMR using a Lorentzian function with line broadening up to 6 Hz. A twosite exchange model was used to determine the rates of nuclei exchange by iteration of the simulated line widths, using both manual and gNMR algorithm methods. Eyring parameters were determined using the equation

ionization (ESI) interface. Bis(hydroxymethylphenylphosphine),70 [Ni(CH3CN)6](BF4)2,71 Ni(PPh2NMe2)2,29 [Cp2Fe][B(C6F5)4],72 and [NBu4][B(C6F5)4]73 were synthesized as previously described. Electrochemical measurements were recorded using a CH Instruments 620D potentiostat equipped with a standard three-electrode cell. The working electrode was a 1 mm PEEK-encased glassy-carbon disk, and the counter electrode was a glassy-carbon rod. For experiments performed in THF, the pseudoreference electrode was a platinum wire. For experiments performed in benzonitrile, the pseudoreference electrode was a silver wire suspended in a 0.2 M NBu4BF4 solution in benzonitrile and separated from the analyte solution by a Vycor frit. Prior to the collection of each voltammogram, the working electrode was polished using 0.25 μm diamond paste on a polishing pad lubricated with ethylene glycol, and then the electrode was rinsed with solvent. Decamethylcobaltocenium hexafluorophosphate (E1/2 = −1.91 V in benzonitrile and −1.94 V in THF versus ferrocenium/ferrocene) was used as an internal reference for H2 oxidation experiments, and all potentials are referenced to the ferrocenium/ferrocene couple at 0.0 V. PPh2NC142. In a 100 mL Schlenk flask, bis(hydroxymethyl)phenylphosphine (1.10 g, 6.47 mmol) was dissolved in 25 mL of absolute ethanol. Tetradecylamine (1.35 g, 6.47 mmol) was dissolved in 25 mL of absolute ethanol and placed in the Schlenk flask. The reaction mixture was stirred at 75 °C for 18 h, during which time a colorless oil formed. The reaction mixture was cooled to room temperature, and the mother liquor was removed from the flask. Additional ethanol (25 mL) was added, and the contents were stirred for 1 h to obtain a white powder. The solution was filtered, and the resulting solid was washed with 10 mL of absolute ethanol and dried under vacuum. The resulting solid was used without further purification (1.14 g). 31P{1H} NMR (toluene): δ −50.1 (br peak). Ni(PPh2NC142)2. A solution of [Ni(MeCN)6](BF4)2 (82 mg, 0.16 mmol) in acetonitrile (3 mL) was added to a suspension of PPh2NC142 (227 mg, 0.33 mmol) in acetonitrile (3 mL). The solution was stirred at room temperature for 16 h, during which time the solution became brown-red. The solution was then sparged with H2 for 10 min, and then n-butylamine (0.065 mL, 1.2 mmol) was added dropwise while the solution was continuously sparged. The headspace of the vial was sparged with H2 for an additional 5 min after the complete addition of the amine, and then the solution was allowed to sit for 1 h. The resulting solid was isolated by vacuum filtration and washed with 5 mL of acetonitrile to isolate a light yellow solid. Crystals were grown from a diethyl ether/acetonitrile solution (120 mg, 0.083 mmol, 50%). 31 1 P{ H} NMR (THF-d8): δ −2.6. 1H NMR (THF-d8): δ 7.77 (br t, J = 6 Hz, 8H, o-C6H5); 7.26 (t, J = 7.4 Hz, 8H, m-C6H5); 7.19 (t, J = 7.3 Hz, 4H, p-C6H5); 3.13 (d, J = 9.2 Hz, 8H, PCH2N); 2.73 (d, J = 11.4 Hz, 8H, PCH2N); 2.57 (t, J = 7.4 Hz, 8H, NCH2CH2−); 1.46 (br t, J = 6.1 Hz, 8H, NCH2CH2CH2−); 1.30 (s, 88 H, alkyl chain); 0.89 (t, J = 6.9 Hz, 12H, −CH2CH3). 13C{1H} NMR (THF-d8): δ 143.4, 131.1, 127.7, 127.3, 62.2, 57.4, 31.9, 29.7, 29.4, 27.1, 26.5, 22.6, 13.5. ESI-MS (positive mode, MeCN/THF): calcd for {[Ni(PPh2NC142)2Cl]}+, 1482.00 m/z; found, 1482.01 m/z. Formation of a chloride adduct is attributed to residual Cl− in the instrument. [Ni(PPh2NC142)2](BF4)2. A solution of [Cp2Fe]BF4 (23 mg, 0.084 mmol) in 4 mL of acetonitrile was added to a suspension of Ni(PPh2NC142)2 (40 mg, 0.028 mmol) in 3 mL of acetonitrile. The reaction mixture was stirred for 1 h at room temperature, resulting in a color change of the solution from light yellow to red. The reaction mixture was filtered through a Celite plug and concentrated to ∼25% of the original reaction volume, and diethyl ether (15 mL) was added to precipitate excess [Cp2Fe]BF4. The solution was filtered, and the solvent was removed under reduced pressure. The resulting solid was triturated with hexanes (15 mL), filtered, and washed with hexanes (15 mL) until the filtrate was colorless, resulting in a red free-flowing solid (40 mg, 0.024 mmol, 87%). 31P{1H} NMR (CD3CN): δ 0.6. 1H NMR (THF-d8): δ 7.35 (t, J = 7.1 Hz, 4H, p-C6H5); 7.29 (br s, 8H, o-C6H5); 7.22 (t, J = 6.6 Hz, 8H, m-C6H5); 3.87 (d, J = 12.7 Hz, 8H, PCH2N); 3.33 (d, J = 12.6 Hz, 8H, PCH2N); 2.98 (t, J = 6.3 Hz, 8H, NCH2CH2); 1.66 (t, J = 6.3 Hz, 8H, NCH2CH2CH2); 1.36 (s, 16 H, alkyl chain); 1.27 (s, 72 H, alkyl chain); 0.88 (t, J = 6.2 Hz, 12H,

k= H

kBT ΔH ⧧/ R ΔS⧧/ R e e h DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics Errors in the activation parameters were determined by two standard deviations from the fit of the line. Determination of Keq for H2 Addition. Solutions of the [Ni]2+ compounds in the desired solvent (PhCN or THF) were added to J. Young NMR tubes under a N2 atmosphere. The tube was subjected to three freeze−pump−thaw cycles to remove the dinitrogen, and then H2 was added to the tube. The solutions were vortexed to ensure thorough mixing, and the J. Young NMR tube was back-filled with H2 and again vortexed. The sample was allowed to equilibrate at room temperature for at least 1 h before analysis, and the equilibrium was confirmed by recording a second spectrum after 24 h. Quantitative 31 1 P{ H} NMR spectra were collected using a 5 s relaxation delay. Computational Details. Calculations were performed using Gaussian 09.77 Structures were optimized in the gas phase without symmetry constraints using the B3P86 functional,40,41 the Stuttgart basis set with effective core potential (ECP) for Ni,42 and the 6-31G** basis set for H, C, N, and P.43,44 Frequency calculations were performed on each structure to confirm the absence of imaginary frequencies for ground-state structures and the presence of a single imaginary frequency for transition-state structures. Solvation free energies were calculated in tetrahydrofuran using the polarizable continuum model52,53 with Bondi78 atomic radii.

■

thank Dr. Rosalie Chu for mass spectroscopy analysis and Dr. Allan Cardenas, Dr. Monte Helm, and Bojana Ginovska for helpful discussions. PNNL is operated by Battelle for DOE.

■

(1) Vincent, K. A.; Parkin, A.; Armstrong, F. A. Chem. Rev. 2007, 107, 4366−4413. (2) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y. Chem. Rev. 2007, 107, 4273−4303. (3) Lubitz, W.; Ogata, H.; Rüdiger, O.; Reijerse, E. Chem. Rev. 2014, 114, 4081−4148. (4) Simmons, T. R.; Berggren, G.; Bacchi, M.; Fontecave, M.; Artero, V. Coord. Chem. Rev. 2014, 270−271, 127−150. (5) Schilter, D.; Camara, J. M.; Huynh, M. T.; Hammes-Schiffer, S.; Rauchfuss, T. B. Chem. Rev. 2016, 116, 8693−8749. (6) Behnke, S. L.; Shafaat, H. S. Comments Inorg. Chem. 2016, 36, 123−140. (7) Rauchfuss, T. B. Acc. Chem. Res. 2015, 48, 2107−2116. (8) Tard, C.; Pickett, C. J. Chem. Rev. 2009, 109, 2245−2274. (9) Camara, J. M.; Rauchfuss, T. B. Nat. Chem. 2011, 4, 26−30. (10) Wang, N.; Wang, M.; Wang, Y.; Zheng, D.; Han, H.; Ahlquist, M. S. G.; Sun, L. J. Am. Chem. Soc. 2013, 135, 13688−13691. (11) Yuki, M.; Sakata, K.; Hirao, Y.; Nonoyama, N.; Nakajima, K.; Nishibayashi, Y. J. Am. Chem. Soc. 2015, 137, 4173−4182. (12) Yuki, M.; Sakata, K.; Kikuchi, S.; Kawai, H.; Takahashi, T.; Ando, M.; Nakajima, K.; Nishibayashi, Y. Chem. - Eur. J. 2017, 23, 1007−1012. (13) Shaw, W. J.; Helm, M. L.; DuBois, D. L. Biochim. Biophys. Acta, Bioenerg. 2013, 1827, 1123−1139. (14) Bullock, R. M.; Helm, M. L. Acc. Chem. Res. 2015, 48, 2017− 2026. (15) Berggren, G.; Adamska, A.; Lambertz, C.; Simmons, T. R.; Esselborn, J.; Atta, M.; Gambarelli, S.; Mouesca, J. M.; Reijerse, E.; Lubitz, W.; Happe, T.; Artero, V.; Fontecave, M. Nature 2013, 499, 66−69. (16) Schilter, D.; Rauchfuss, T. B. Angew. Chem., Int. Ed. 2013, 52, 13518−13520. (17) Fraze, K.; Wilson, A. D.; Appel, A. M.; Rakowski DuBois, M.; DuBois, D. L. Organometallics 2007, 26, 3918−3924. (18) Yang, J. Y.; Smith, S. E.; Liu, T.; Dougherty, W. G.; Hoffert, W. A.; Kassel, W. S.; Rakowski DuBois, M.; DuBois, D. L.; Bullock, R. M. J. Am. Chem. Soc. 2013, 135, 9700−9712. (19) Das, P.; Ho, M.-H.; O’Hagan, M.; Shaw, W. J.; Bullock, R. M.; Raugei, S.; Helm, M. L. Dalton Trans. 2014, 43, 2744−2754. (20) Stolley, R. M.; Darmon, J. M.; Das, P.; Helm, M. L. Organometallics 2016, 35, 2965−2974. (21) Smith, S. E.; Yang, J. Y.; DuBois, D. L.; Bullock, R. M. Angew. Chem., Int. Ed. 2012, 51, 3152−3155. (22) Dutta, A.; DuBois, D. L.; Roberts, J. A. S.; Shaw, W. J. Proc. Natl. Acad. Sci. U. S. A. 2014, 111, 16286−16291. (23) Priyadarshani, N.; Dutta, A.; Ginovska, B.; Buchko, G. W.; O’Hagan, M.; Raugei, S.; Shaw, W. J. ACS Catal. 2016, 6, 6037−6049. (24) Raugei, S.; Helm, M. L.; Hammes-Schiffer, S.; Appel, A. M.; O’Hagan, M.; Wiedner, E. S.; Bullock, R. M. Inorg. Chem. 2016, 55, 445−460. (25) Raugei, S.; Chen, S.; Ho, M.-H.; Ginovska-Pangovska, B.; Rousseau, R. J.; Dupuis, M.; DuBois, D. L.; Bullock, R. M. Chem. - Eur. J. 2012, 18, 6493−6506. (26) Stolley, R. M.; Darmon, J. M.; Helm, M. L. Chem. Commun. 2014, 50, 3681−3684. (27) Lense, S.; Ho, M.-H.; Chen, S.; Jain, A.; Raugei, S.; Linehan, J. C.; Roberts, J. A. S.; Appel, A. M.; Shaw, W. Organometallics 2012, 31, 6719−6731. (28) Ginovska-Pangovska, B.; Dutta, A.; Reback, M. L.; Linehan, J. C.; Shaw, W. J. Acc. Chem. Res. 2014, 47, 2621−2630. (29) Yang, J. Y.; Bullock, R. M.; Dougherty, W. G.; Kassel, W. S.; Twamley, B.; DuBois, D. L.; Rakowski DuBois, M. Dalton Trans. 2010, 39, 3001−3010.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.organomet.7b00103. Multinuclear and variable-temperature NMR spectra, cyclic voltammograms, kinetic plots for H2 binding and electrocatalytic H2 oxidation, details of NBO analysis (PDF) All computed molecule Cartesian coordinates (XYZ) Accession Codes

CCDC 1546104−1546105 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.

■

REFERENCES

AUTHOR INFORMATION

Corresponding Author

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

R. Morris Bullock: 0000-0001-6306-4851 Aaron M. Appel: 0000-0002-5604-1253 Eric S. Wiedner: 0000-0002-7202-9676 Notes

The authors declare no competing financial interest.

■

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, Office of Science, Office of Basic Energy Sciences. Computational resources were provided at the National Energy Research Scientific Computing Center (NERSC) at Lawrence Berkeley National Laboratory. Mass spectrometry experiments were performed in the William R. Wiley Environmental Molecular Sciences Laboratory, a DOE national scientific user facility sponsored by the DOE’s Office of Biological and Environmental Research and located at the Pacific Northwest National Laboratory (PNNL). The authors I

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics (30) Cardenas, A. J. P.; Ginovska, B.; Kumar, N.; Hou, J.; Raugei, S.; Helm, M. L.; Appel, A. M.; Bullock, R. M.; O’Hagan, M. Angew. Chem., Int. Ed. 2016, 55, 13509−13513. (31) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: Hoboken, NJ, 2001. (32) Miedaner, A.; Haltiwanger, R. C.; DuBois, D. L. Inorg. Chem. 1991, 30, 417−427. (33) Wilson, A. D.; Newell, R. H.; McNevin, M. J.; Muckerman, J. T.; Rakowski DuBois, M.; DuBois, D. L. J. Am. Chem. Soc. 2006, 128, 358−366. (34) Galan, B. R.; Schoffel, J.; Linehan, J. C.; Seu, C.; Appel, A. M.; Roberts, J. A. S.; Helm, M. L.; Kilgore, U. J.; Yang, J. Y.; DuBois, D. L.; Kubiak, C. P. J. Am. Chem. Soc. 2011, 133, 12767−12779. (35) Seu, C. S.; Appel, A. M.; Doud, M. D.; DuBois, D. L.; Kubiak, C. P. Energy Environ. Sci. 2012, 5, 6480−6490. (36) Barrière, F.; Geiger, W. E. J. Am. Chem. Soc. 2006, 128, 3980− 3989. (37) Geiger, W. E.; Barrière, F. Acc. Chem. Res. 2010, 43, 1030−1039. (38) Jain, A.; Lense, S.; Linehan, J. C.; Raugei, S.; Cho, H.; DuBois, D. L.; Shaw, W. J. Inorg. Chem. 2011, 50, 4073−4085. (39) Brookhart, M.; Green, M. L. H.; Parkin, G. Proc. Natl. Acad. Sci. U. S. A. 2007, 104, 6908−6914. (40) Becke, A. D. J. Chem. Phys. 1993, 98, 5648−5652. (41) Perdew, J. P. Phys. Rev. B: Condens. Matter Mater. Phys. 1986, 33, 8822−8824. (42) Andrae, D.; Haeussermann, U.; Dolg, M.; Stoll, H.; Preuss, H. Theor. Chim. Acta 1990, 77, 123−141. (43) Francl, M. M.; Pietro, W. J.; Hehre, W. J.; Binkley, J. S.; Gordon, M. S.; DeFrees, D. J.; Pople, J. A. J. Chem. Phys. 1982, 77, 3654−65. (44) Rassolov, V. A.; Ratner, M. A.; Pople, J. A.; Redfern, P. C.; Curtiss, L. A. J. Comput. Chem. 2001, 22, 976−984. (45) Scherer, W.; Dunbar, A. C.; Barquera-Lozada, J. E.; Schmitz, D.; Eickerling, G.; Kratzert, D.; Stalke, D.; Lanza, A.; Macchi, P.; Casati, N. P. M.; Ebad-Allah, J.; Kuntscher, C. Angew. Chem., Int. Ed. 2015, 54, 2505−2509. (46) Zhang, Y.; Lewis, J. C.; Bergman, R. G.; Ellman, J. A.; Oldfield, E. Organometallics 2006, 25, 3515−3519. (47) Yao, W.; Eisenstein, O.; Crabtree, R. H. Inorg. Chim. Acta 1997, 254, 105−111. (48) Glendening, E. D.; Reed, A. E.; Carpenter, J. E.; Weinhold, F. NBO, Version 3.1. (49) Lein, M. Coord. Chem. Rev. 2009, 253, 625−634. (50) Thakur, T. S.; Desiraju, G. R. J. Mol. Struct.: THEOCHEM 2007, 810, 143−154. (51) Lein, M.; Harrison, J. A.; Nielson, A. J. Dalton Trans. 2013, 42, 10939−10951. (52) Cossi, M.; Rega, N.; Scalmani, G.; Barone, V. J. Comput. Chem. 2003, 24, 669−681. (53) Barone, V.; Cossi, M. J. Phys. Chem. A 1998, 102, 1995−2001. (54) Rauk, A.; Andose, J. D.; Frick, W. G.; Tang, R.; Mislow, K. J. Am. Chem. Soc. 1971, 93, 6507−6515. (55) Franz, J. A.; O’Hagan, M.; Ho, M.-H.; Liu, T.; Helm, M. L.; Lense, S.; DuBois, D. L.; Shaw, W. J.; Appel, A. M.; Raugei, S.; Bullock, R. M. Organometallics 2013, 32, 7034−7042. (56) Wilson, A. D.; Shoemaker, R. K.; Miedaner, A.; Muckerman, J. T.; DuBois, D. L.; Rakowski DuBois, M. Proc. Natl. Acad. Sci. U. S. A. 2007, 104, 6951−6956. (57) Wiedner, E. S.; Yang, J. Y.; Chen, S.; Raugei, S.; Dougherty, W. G.; Kassel, W. S.; Helm, M. L.; Bullock, R. M.; Rakowski DuBois, M.; DuBois, D. L. Organometallics 2012, 31, 144−156. (58) O’Hagan, M.; Shaw, W. J.; Raugei, S.; Chen, S.; Yang, J. Y.; Kilgore, U. J.; DuBois, D. L.; Bullock, R. M. J. Am. Chem. Soc. 2011, 133, 14301−14312. (59) O’Hagan, M.; Ho, M.-H.; Yang, J. Y.; Appel, A. M.; Rakowski DuBois, M.; Raugei, S.; Shaw, W. J.; DuBois, D. L.; Bullock, R. M. J. Am. Chem. Soc. 2012, 134, 19409−19424. (60) Ho, M.-H.; O’Hagan, M.; Dupuis, M.; DuBois, D. L.; Bullock, R. M.; Shaw, W. J.; Raugei, S. Dalton Trans. 2015, 44, 10969−10979.

(61) Yang, J. Y.; Bullock, R. M.; Shaw, W. J.; Twamley, B.; Fraze, K.; Rakowski DuBois, M.; DuBois, D. L. J. Am. Chem. Soc. 2009, 131, 5935−5945. (62) Nicholson, R. S.; Shain, I. Anal. Chem. 1964, 36, 706−723. (63) Saveant, J. M.; Vianello, E. Electrochim. Acta 1965, 10, 905−920. (64) Coetzee, J. F.; Padmanabhan, G. R. J. Am. Chem. Soc. 1965, 87, 5005−5010. (65) Roberts, J. A. S.; Bullock, R. M. Inorg. Chem. 2013, 52, 3823− 3835. (66) Wiedner, E. S.; Yang, J. Y.; Dougherty, W. G.; Kassel, W. S.; Bullock, R. M.; Rakowski DuBois, M.; DuBois, D. L. Organometallics 2010, 29, 5390−5401. (67) Hulley, E. B.; Helm, M. L.; Bullock, R. M. Chem. Sci. 2014, 5, 4729−4741. (68) Yang, J. Y.; Chen, S.; Dougherty, W. G.; Kassel, W. S.; Bullock, R. M.; DuBois, D. L.; Raugei, S.; Rousseau, R.; Dupuis, M.; Rakowski DuBois, M. Chem. Commun. 2010, 46, 8618−8620. (69) Kilgore, U. J.; Stewart, M. P.; Helm, M. L.; Dougherty, W. G.; Kassel, W. S.; Rakowski DuBois, M.; DuBois, D. L.; Bullock, R. M. Inorg. Chem. 2011, 50, 10908−10918. (70) G. Märkl, V.; Jin, G. Y.; Schoerner, C. Tetrahedron Lett. 1980, 21, 1409−1412. (71) Heintz, R. A.; Szalay, P. S.; Weisgerber, A.; Dunbar, K. R. Useful Reagents and Ligands. In Inorganic Syntheses; Wiley: Hoboken, NJ, 2002; Vol. 33, pp 75−121. (72) Spinney, H. A.; Clough, C. R.; Cummins, C. C. Dalton Trans. 2015, 44, 6784−6796. (73) Ramsay, W. J.; Nitschke, J. R. J. Am. Chem. Soc. 2014, 136, 7038−7043. (74) APEX3, SAINT, and SADABS; Bruker AXS, Inc., Madison, WI, 2015. (75) Sheldrick, G. M. SHELXTL, Version 6.14; Bruker AXS Inc., Madison, WI, 2004. (76) Budzelaar, P. H. M. gNMR, 5.0 ed.; 1995−2006. (77) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S. A.; Daniels, D.; Farkas, O.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09, Revision D.01; Gaussian, Inc., Wallingford, CT, 2013. (78) Bondi, A. J. Phys. Chem. 1964, 68, 441−451.

J

DOI: 10.1021/acs.organomet.7b00103 Organometallics XXXX, XXX, XXX−XXX