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Article Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

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Mechanistic Insights into Homogeneous Electrocatalytic and Photocatalytic Hydrogen Evolution Catalyzed by High-Spin Ni(II) Complexes with S2N2‑Type Tetradentate Ligands Dachao Hong,†,‡ Yuto Tsukakoshi,† Hiroaki Kotani,† Tomoya Ishizuka,† Kei Ohkubo,∥,# Yoshihito Shiota,¶ Kazunari Yoshizawa,¶,§ Shunichi Fukuzumi,#,◊ and Takahiko Kojima*,† †

Department of Chemistry, Faculty of Pure and Applied Sciences, University of Tsukuba, CREST, Japan Science and Technology Agency, 1-1-1 Tennoudai, Tsukuba, Ibaraki 305-8571, Japan ‡ Interdisciplinary Research Center for Catalytic Chemistry, National Institute of Advanced Industrial Science and Technology, 1-1-1 Higashi, Tsukuba, Ibaraki 305-8565, Japan ∥ Institute for Advanced Co-Creation Studies and Institute for Academic Initiatives, Osaka University, Suita, Osaka 565-0871, Japan # Department of Chemistry and Nano Science, Ewha Womans University, Seoul 120-750, Korea ¶ Institute for Materials Chemistry and Engineering, Kyushu University, CREST, Japan Science and Technology Agency, Motooka, Nishi-Ku, Fukuoka 819-0395, Japan § Elements Strategy Initiative for Catalysts & Batteries, Kyoto University, Nishikyo-ku, Kyoto 615-8520, Japan ◊ Faculty of Science and Technology, Meijo University, SENTAN, Japan Science and Technology Agency, Nagoya, Aichi 468-8502, Japan S Supporting Information *

ABSTRACT: We report homogeneous electrocatalytic and photocatalytic H2 evolution using two Ni(II) complexes with S2N2-type tetradentate ligands bearing two different sizes of chelate rings as catalysts. A Ni(II) complex with a five-membered SC2S−Ni chelate ring (1) exhibited higher activity than that with a six-membered SC3S−Ni chelate ring (2) in both electrocatalytic and photocatalytic H2 evolution despite both complexes showing the same reduction potentials. A stepwise reduction of the Ni center from Ni(II) to Ni(0) was observed in the electrochemical measurements; the first reduction is a pure electron transfer reaction to form a Ni(I) complex as confirmed by electron spin resonance measurements, and the second is a 1e−/1H+ proton-coupled electron transfer reaction to afford a putative Ni(II)-hydrido (NiII−H) species. We also clarified that Ni(II) complexes can act as homogeneous catalysts in the electrocatalytic H2 evolution, in which complex 1 exhibited higher reactivity than that of 2. In the photocatalytic system using [Ru(bpy)3]2+ as a photosensitizer and sodium ascorbate as a reductant, complex 1 with the five-membered chelate ring also showed higher catalytic activity than that of 2 with the sixmembered chelate ring, although the rates of photoinduced electron-transfer processes were comparable. The Ni−H bond cleavage in the putative NiII−H intermediate should be involved in the rate-limiting step as evidenced by kinetic isotope effects observed in both photocatalytic and electrocatalytic H2 evolution. Kinetic analysis and density functional theory calculations indicated that the difference in H2 evolution activity between the two complexes was derived from that of activation barriers of the reactions between the NiII−H intermediates and proton, which is consistent with the fact that increase of proton concentration accelerates the H2 evolution.

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INTRODUCTION Converting light energy into chemical energy such as H2 is an urgent issue to solve the confronting energy problem and to realize a sustainable society.1−4 H2 has the highest mass energy density and is a clean energy carrier to produce only water after combustion.5 Noble metals such as Pt and Ru have been reported to catalyze reduction of proton to generate H2 under electrocatalytic and photocatalytic conditions in early works.6−12 Recently, to reduce the use of precious metals, © XXXX American Chemical Society

many researchers have devoted their enormous efforts to developing H2 evolution catalysts with earth-abundant metals in homogeneous and heterogeneous forms.13−23 In nature, [NiFe] hydrogenases containing Ni−Fe active sites are known to catalyze proton reduction to produce H2.24 The structure of the active site of [NiFe] hydrogenases has Received: April 2, 2018

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DOI: 10.1021/acs.inorgchem.8b00881 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

4.0) by using carbon papers (CP) as working electrodes. In the photocatalytic H2 evolution, we used [Ru(bpy)3]Cl2 (bpy = 2,2′-bipyridine), whose one-electron reduced [Ru(bpy)3]2+, [Ru(bpy●−)(bpy)2]+, shows an oxidation potential of −1.33 V (vs standard calomel electrode (SCE) in MeCN),35 as a photosensitizer and sodium ascorbate (NaAsc) as a reductant, together with the Ni(II) catalysts 1 and 2. Mechanistic insights into catalytic H2 evolution by the Ni(II) complexes will be discussed on the basis of kinetic, electrochemical, and spectroscopic investigation, together with density functional theory (DFT) calculations.

been determined to be a S-bridged dinuclear Ni−Fe complex coordinated by CO, CN−, and cysteine-derived thiolato ligands by X-ray crystallography.25 Inspired by the hydrogenase clusters, many studies have focused on molecular catalysts based on transition-metal complexes, especially Ni complexes, for H2 evolution catalysts under photocatalytic or electrocatalytic conditions.26,27 Recently, DuBois and co-workers have reported an efficient Ni catalyst with cyclic diphosphine ligands forming a sevenmembered chelate ring, which should show large flexibility in conformational freedom, for electrocatalytic H2 evolution.28 A series of mononuclear NiII bis(diphosphine) complexes was also reported to catalyze proton reduction efficiently at −0.90 V (vs Fc/Fc+, Fc = ferrocene) in acetonitrile.29 Additionally, a Ni complex with two nitrogen and phosphorus-type ligands exhibited a turnover frequency (TOF) of 8400 s−1 in electrocatalytic H2 evolution with use of acetic acid as a proton source.30 On the other hand, homogeneous systems for photocatalytic H2 evolution have been developed to gain mechanistic insights into the formation of H2. Photocatalytic H2 evolution has been achieved by using Ni(II) complexes with a 2-mercapto-benzimidazole ligand as catalysts, fluorescein (Fl) as a photosensitizer, and triethylamine (TEA) as a sacrificial electron donor under basic conditions.31 An efficient catalyst of [Ni(pyS)3] (pyS = pyridine-2-thiolate) can afford a high turnover number (TON) of 5500 at 40 h (250 h−1 as a TOF) under photocatalytic conditions with Fl and TEA.32 In photocatalytic H2 evolution, [Ir(ppy)3] (ppy = 2-phenylpyridine) was also used as a photosensitizer. Several Ni complexes having tetradentate macrocyclic ligands have been also employed for photocatalytic H2 evolution in the presence of an [Ir(ppy)3] derivative.33 In those early works, Ni complexes bearing S and P atoms as coordination sites of ligands showed high activity in both electrocatalytic and photocatalytic H2 evolution. Herein, we report mechanistic insights into electrocatalytic and photocatalytic H2 evolution using two Ni(II) complexes with tetradentate ligands, namely, [Ni(bpet)(CH3CN)2](ClO4)2 (1·(ClO4)2)34 and [Ni(bppt)(CH3CN)2](BPh4)2 (2· (BPh4)2), (bpet = bis(2-pyridylmethyl)-1,2-ethanedithiol, bppt = bis(2-pyridylmethyl)-1,3-propanedithiol), as catalysts (Chart 1). The tetradentate ligands of bpet (five-membered SC2S−Ni

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RESULTS AND DISCUSSION Preparations and Crystal Structures of Ni(II) Complexes. The ligands bpet and bppt and the Ni(II) complexes were prepared based on reported methods for Cu(II) complexes.36 The synthetic procedure is described in the Experimental Section in detail. The Ni(II) complexes were characterized by 1H NMR and mass spectrometry (MS) as shown in Figures S1 and S2 in the Supporting Information. The crystal structures of 134,37 and 238 were successfully determined by X-ray crystallography. The ORTEP drawings of the cationic moieties of 1 and 2 are depicted in Figure 1. The Ni−S bond

Chart 1. Chemical Structures of Ni Complexes with S2N2Type Tetradentate Ligands

Figure 1. ORTEP drawings and selected bond lengths and angles of the cationic moieties of 134 and 2 with 50% probability thermal ellipsoids. Hydrogen atoms are omitted for clarity.

lengths in 1 were 2.398 Å (for Ni−S1)39 and 2.394 Å (for Ni− S2),39 which were slightly shorter than those of 2 (Ni−S1 = 2.4467(6) Å). The angles of S1−Ni−S2 and N1−Ni−N2 are 89.61°39 and 171.6°39 for 1 and 89.91(3)° and 178.57(8)° for 2. The structural analysis indicates that 1 is in a tenser structure than 2, suggesting that 1 should be more reactive than 2. Redox Behavior and Electrolysis of Ni(II) Complexes. Electrochemical properties of the two Ni(II) complexes were investigated by cyclic voltammetry (CV) in MeCN as shown in Figure 2a, in which the redox potentials were calibrated relative to that of the ferrocene/ferrocenium (Fc/Fc+) redox couple as 0 V in MeCN. The first reduction waves of 1 and 2 were observed at −1.21 and −1.19 V (vs Fc/Fc+), respectively, corresponding to −0.81 and −0.79 V (vs SCE).40,41 The second reduction waves appeared at the same potential of

chelate ring) and bppt (six-membered SC3S−Ni chelate ring) provide S2N2-type coordination environments with different chelate-ring sizes designed to study dependence of H2 evolution activity on the chelate-ring flexibility. The redox property of the Ni(II) complexes was investigated to elucidate the reaction mechanism of H2 evolution. In the electrocatalytic H2 evolution, we conducted the reaction using the Ni(II) complexes as homogeneous catalysts in an ascorbate buffer (pH B


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Inorganic Chemistry

Figure 2. CVs of an MeCN solution (2.0 mL) containing a Ni complex (1.0 mM, 1: red line, 2: blue line) and 0.10 M TBAPF6 in the scan range (a) from 0 to −2.5 V and (b) from 0 to −1.4 V by using glassy carbon (A = 0.071 cm2) as a working electrode, a Pt wire as a counter electrode, and an Ag/AgNO3 reference electrode at the scan rate of 0.10 V s−1.

−1.56 V (vs Fc/Fc, −1.16 V vs SCE) for both complexes. The reduction waves of 1 and 2 were well-reproduced in the multiple scans under the same conditions (Figure S3). The first reduction waves can be assigned to the reduction of the NiII/ NiI process, and the second can be assigned to the NiI/Ni0 one. The CVs clearly demonstrate stepwise reduction of Ni(II) complexes, referred to as [LNiII]2+ (L = bpet, bppt), to [LNiI]+, followed by [LNi0]0 as shown in eq 1. 1e−

1e−

[LNi II]2 + Xoooo−Y [LNi I]+ ⎯→ ⎯ [LNi0]0 −1e

Figure 3. (a) CVs of MeCN solutions (2.0 mL) containing 1 (1.0 mM), AcOH (0 mM: black line, 0.03 mM: purple line, 0.05 mM: green line, 0.10 mM: yellow line, 0.5 mM: blue line, 1.0 mM: red line), and 0.10 M TBAPF6 by using a glassy carbon (A = 0.071 cm2) as a working electrode, a Pt wire as a counter electrode, and an Ag/AgNO3 reference electrode at the scan rate of 0.10 V s−1. The potential was referenced relative to the Fc/Fc+ couple as 0 V. (b) An enlarged view of reduction waves of (a) in the range from −1.1 to −1.7 V. (c) Plots of peak potentials of the first (red ●) and second (black ■) reduction waves relative to the concentration of AcOH. (d) CVs of MeCN solutions (2.0 mL) containing AcOH (43.7 mM) with 1 (1.0 mM, red line) and without 1 (gray line).

(1)

When the negative scan was reversed before the second wave, the first redox waves for 1 and 2 exhibited a quasireversible feature with a peak separation of 160 mV, which was larger than 57 mV (Figure 2b). The irreversibility of the second wave can be ascribed to a large structural change of a plausibly octahedral NiI complex (d9 electronic configuration) to a tetrahedral Ni0 complex in the d10 electronic configuration in the course of electrochemical reduction.34 An irreversible wave around −0.25 V was also observed in the CVs of both 1 and 2. The CVs of 1 with scan rate dependence on peak current was conducted to identify the irreversible wave (Figures S4 and S5). The peak currents of NiII/NiI, NiI/Ni0, and the irreversible wave were proportional to the square root of scan rates, representing the feature of irreversible redox waves rather than an absorption wave. The irreversible wave disappeared when the scan was reverted at the first reduction wave. Thus, the peak around −0.2 V can be attributed to a reoxidation wave of a Ni0 species with a large structural change or that due to a Ni0 species that underwent an electrochemical−chemical process. Effect of Proton Concentration on the Redox Potentials of 1. First, we examined the impact of proton on the complexes in MeCN by UV−vis spectroscopy. No spectral change was observed for UV−vis spectra of the Ni(II) complexes 1 and 2 by the addition of AcOH up to 291 mM (Figure S6), indicating no protonation occurs on the complexes under the conditions. Considering no protonation occurs to the Ni(II) complexes under the condition, we performed CV on 1 in MeCN solutions in the presence of various concentrations of acetic acid (AcOH) as shown in Figure 3a,b. The first reduction wave remained the same, but the second one positively shifted by 67 mV in accordance with the increase of the AcOH concentration ([AcOH]) up to 1.0 mM as shown in Figure 3c. The result indicates the proton concentration has influenced the reduction of NiI to Ni0 rather than NiII to NiI. The initial slop of the second reduction wave was calculated to be −54 mV (Ered/−log[AcOH]), suggesting proton-coupled electron trans-

fer (PCET),42,43 in which one-electron reduction coupled with one-proton acceptance occurred after one-electron reduction of [LNiII]2+ as shown in eq 2. The further increase of AcOH (up to 17.5 mM) resulted in disappearance of the second reduction waves and appearance of catalytic current (Figure S7). The phenomenon was also observed in CV measurements on 2 with AcOH (Figure S8). The overpotential for 1 and 2 in MeCN/ AcOH was calculated to be η = 0.55 and 0.56 V, respectively, based on the midpoint potentials of catalytic waves (Figure S9). 44 The overpotentials can be greater when the homoconjugation of acetic acid was taken into account.44a When higher concentration of AcOH was used, catalytic current was observed as shown in Figure 3d. The onset potential of catalytic current for 1 is positively shifted by 0.28 V as compared to that without 1, indicating the catalysis of 1. Electrocatalytic H2 evolution using 1 and 2 in MeCN/AcOH under the applied potential of −1.7 V (vs SCE) was confirmed by GC as shown in Figure S10. 1e−

1e−,1H+

[LNi II]2 + Xoooo−Y [LNi I]+ ⎯⎯⎯⎯⎯⎯⎯→ [LNi II − H]+ −1e

(2)

Electrocatalytic H2 Evolution Using 1 and 2 as Catalysts. Electrocatalytic H2 evolution was also conducted by bulk electrolysis at the applied potential of −0.90 V (vs SCE) in an ascorbate buffer solution (pH 4, 15 mL) containing a catalyst (1, 2, or Ni(ClO4)2) and without a catalyst, using a CP (A = 2.0 cm2) electrode as a working electrode in the twocomponent cell (Figure S11). The overpotential of the bulk electrolysis was calculated to be 0.42 V on the basis of the applied potential (−0.90 V) and E°(H2/2H+) at pH 4.0 (−0.48 V vs SCE). The amount of catalysts used in the electrocatalytic H2 evolution was optimized to be 80 nmol by independent CV measurements with electrodes, where catalysts were immobiC


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Inorganic Chemistry lized as shown in Figure S12. H2 evolution was observed, and the amount of H2 evolved was quantified by gas chromatography (GC). Time courses of H2 evolution in the bulk electrolysis are shown in Figure 4. Complex 1 afforded the

Figure 5. Time course of H2 evolution under (a) bulk electrolysis of a various buffer (pH 4.0−7.0, 15 mL) containing 1 (80 nmol) by a pristine CP working electrode. (b) Bulk electrolysis of H2O or D2O solution (0.10 M sodium ascorbate) containing 1 using a pristine CP working electrode.

proton is involved in the rate-determining step. Deuterium oxide (D2O) was also used for the bulk electrolysis with 1 as shown in Figure 5b. The TOF in D2O (TOFD2O) was determined to be 50 h−1, which was smaller than that in H2O (TOFH2O: 2.3 × 102 h−1). The deuterium kinetic isotope effect (KIE = TOFH2O/TOFD2O) on the H2 evolution reaction was determined to be 4.6, which was larger than that (2.2) in the photocatalytic H2 evolution (vide infra). The KIE value indicates that the NiII−H bond cleavage to produce H2 is involved in the rate-determining step in the electrocatalytic H2 evolution. Mechanism of Electrocatalytic H2 evolution. On the basis of the CV results in Figures 2 and 3, the mechanism of electrocatalytic H2 evolution is proposed as depicted in Scheme 1.46 The Ni(II) complexes were reduced stepwise by two

Figure 4. Time course of H2 evolution under bulk electrolysis of an ascorbate buffer (pH 4.0, 15 mL) at −0.90 V (vs SCE) containing 1 (80 nmol, red ●), 2 (80 nmol, blue ■), and Ni(ClO4)2 (80 nmol, orange ▲) or without a catalyst (black ◆) by using a pristine CP electrode (blue ■), a Pt wire, and a Ag/AgCl (saturated KCl solution) electrode as a working electrode, a counter electrode, and a reference electrode, respectively.

largest amount of H2 evolved (90 μmol) with an apparent TON of 1.1 × 103 {apparent TON = (mole of evolved H2)/ (mole of Ni complex; 80 nmol)} compared to the apparent TON for 2 (TON = 19) at 5 h (Table 1). The reason for the Table 1. Summary of Electrocatalytic H2 Evolution with 1 and 2 as Catalystsa cat

H2, μmol

FE (%)

apparent TONb

apparent TOF,c h−1

1 2

90 1.5

95 80

1.1 × 103 19

2.3 × 102 3.6

Scheme 1. Proposed Catalytic Mechanism of Electrocatalytic H2 Evolution with 1 and 2

Bulk electrolysis at −0.90 V (vs SCE) in ascorbate buffer (pH 4.0) at 298 K. bApparent TON in 5 h electrolysis: Apparent TON = (mole of evolved H2)/(mole of Ni complex; 80 nmol). cApparent TOFs were calculated based on the slop of H2 evolution for 5 h. a

low TON of 2 should be the low solubility of 2 in water as the hydrophobic BPh4− salt. The Faraday efficiency (FE) of H2 evolution for 1 and 2 reached 95% and 80% at 5 h, respectively. Only little H2 evolution was observed in control bulk electrolysis using Ni(ClO4)2 as a catalyst and without a catalyst at 5 h reaction time. After the electrolysis, the CP electrode in the solution of 1 was investigated by scanning electron microscopy−energy dispersive X-ray spectroscopy (SEMEDX) to check whether NiS film formed or not;45 no Ni atoms was detected on the electrode, indicating that 1 does not undergo decomposition to form catalytically active materials as reported by Dempsey and co-workers.45 Those results indicate 1 acts as an efficient homogeneous catalyst for electrochemical H2 evolution in an aqueous solution. Dependence of Electrocatalytic H2 Evolution of 1 on pH. We performed homogeneous bulk electrolysis in aqueous buffer solutions of 1 at various pH values to investigate dependence of the catalytic activity on the proton concentration. When pH was raised, the amount of evolved H2 was significantly decreased. The amount of H2 evolution at 5 h for pH 5, 6, and 7 reached only 32, 9.7, and 1.7 μmol, respectively (Figure 5a). Thus, the proton concentration should affect directly the efficiency of the H2 evolution, suggesting that

electrons to form [LNi0]0 as evidenced by the observation of two reduction waves in MeCN. As demonstrated in Figure 3b, the second reduction step should be PCET involving one electron and one proton. It suggests that the [LNiI]+ complex undergoes PCET to form a NiII−H intermediate, [LNiII−H]+, in the presence of proton (eq 2). Although the formation of [LNiII−H]+ species through direct two-electron reduction has been also proposed in literature,46 in our system, the formation should proceed in a stepwise mechanism including the 1e−/ 1H+ PCET step. H2 was evolved by a subsequent reaction of [LNiII−H]+ with a proton. At a higher proton concentration, the reduction wave due to the process from [LNiI]+ to [LNi0]0 could not be observed due to the overlap of a large catalytic current wave, although it was observed at lower proton concentrations. The observations of the PCET process at lower proton concentrations and overlap of the catalytic current over the second PCET reduction wave in the CV suggest that the D


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Inorganic Chemistry

obtained by 1 (Figure S15). The highest efficiency in the photocatalytic H2 evolution was attained when a DMA/ ascorbate buffer (1:1 v/v) was employed as a mixed solution. The amount of H2 evolution using 1 and 2 was significantly decreased, when the ascorbate buffer was replaced by H2O in the mixed solution (Figure S16), suggesting that the proton concentration should be an important factor for the photocatalytic H2 evolution, which is consistent with the arguments described in the electrochemical H2 evolution section. The dependence of the H2 evolution rates on the concentration of 1 was also examined, as shown in Figure S17. The amount of H2 evolution increased linearly in accordance with an increase of the concentration of 1. This result suggests that H2 evolved by the electrophilic attack of proton to a NiII−H intermediate derived from 1 to release H2 (Scheme 2). The deuterium KIE on the photocatalytic H2

rate-determining step in the present H2 evolution reaction should be the reaction of NiII−H species with proton rather than the formation of the NiII−H intermediate. Considering the activities of 1 and 2 in the bulk electrolysis, a small difference in chelating ring members of ligands between 1 including a five-membered chelate ring for the dithioether moiety and 2 having a six-membered chelate ring for the moiety in Chart 1 resulted in remarkable difference in TOFs; TOF for 1 was more than 50 times larger than that for 2 (Table 1), despite the two complexes showing similar reduction potentials (Figure 2). We assume that the difference in the catalytic activity between 1 and 2 would be derived from that of the reactivity of putative [LNiII−H]+ toward proton to form H2 in the rate-limiting step as mentioned above. Photocatalytic H2 Evolution Catalyzed by Ni Complexes. To gain mechanistic insights into H2 evolution by 1 and 2, we examined photocatalytic H2 evolution by using [Ru(bpy)3]2+ as a photosensitizer and sodium ascorbate (NaAsc) as a sacrificial electron donor. The reaction was performed under photoirradiation (λ = 450 nm) of a N,Ndimethylacetamide (DMA)/ascorbate buffer (pH 4, 4.0 mL, 1:1 v/v) containing 1 or 2 (0.10 mM), [Ru(bpy)3]2+ (0.25 mM), and NaAsc (0.10 M). The amount of H2 evolved was quantified by GC analysis. The time courses of H2 evolution with 1 and 2 are displayed in Figure 6. The amount of H2 evolved by 1 (14

Scheme 2. Proposed Catalytic Cycle of Photocatalytic H2 Evolution with Ni Complexes

evolution with 1 and 2 was investigated as shown in Figure S18. The KIE values were determined to be 2.2 for 1 and 1.4 for 2 under the photocatalytic conditions, which are smaller than that of 1 (4.6) in the electrocatalytic H2 evolution as described above. Those results indicate that the Ni−H bond cleavage should be also involved in the rate-determining step of the photocatalytic H2 evolution. Catalytic Cycle of Photocatalytic H2 Evolution. In light of the redox potentials determined by electrochemical measurements, the Ni(II) complexes 1 and 2 can be reduced by a one-electron-reduced species of the photosensitizer, [Ru(bpy●−)(bpy)2]+, which shows Eox = −1.33 V versus SCE.35 In addition, photoinduced electron transfer from ascorbate to the triplet metal-to-ligand charge transfer (MLCT) excited state of the photosensitizer ( 3 {[Ru(bpy)3]2+}*) has been well-established: The oxidation potential of Asc− has been determined to be +0.15 V (vs SCE), and the reduction potential of 3{[Ru(bpy)3]2+}* has been reported to be +0.84 V (vs SCE),35 indicating the electron transfer (ET) reaction should be downhill to occur due to the positive driving force of ET. Together with experimental results mentioned above, we propose a plausible catalytic cycle for the photocatalytic H2 evolution using the Ni(II) complex 1 or 2 as catalysts and [Ru(bpy)3]2+ as the photosensitizer as shown in Scheme 2.47 Upon photoirradiation to [Ru(bpy)3]2+ in the presence of an electron donor, the one-electron reduced complex ([Ru(bpy●−)(bpy)2]+) was produced by reductive quenching of 3{[Ru(bpy)3]2+}* by an electron donor (D, ascorbate in this case). Stepwise two-electron reduction of the Ni(II) complex by 2 equiv of [Ru(bpy●−)(bpy)2]+ yields a NiII−H intermediate, which is able to react with protons to generate H2.

Figure 6. Time courses of H2 evolution under photoirradiation (λ = 450 nm) of a DMA/ascorbate buffer (pH 4, 4.0 mL, 1:1 v/v) containing a Ni(II) catalyst {(0.10 mM), 1 for red ●, 2 for blue ■, and Ni(ClO4)2 for green ◆}, [Ru(bpy)3]2+ (0.25 mM), and NaAsc (0.10 M).

μmol) in 4 h is ∼10 times larger than that of 2 (1.5 μmol). The TONs for 1 and 2 at 4 h were determined to be 36 and 3.8, respectively. The quantum yield of photocatalytic H2 evolution with 1 was determined to be 0.78% (see Figure S13). Only a negligible amount of H2 was evolved in a control experiment by using Ni(ClO4)2 that can be a source of Ni nanoparticles as an actual catalyst under the same conditions, suggesting complexes 1 and 2 act as homogeneous catalysts in the photocatalytic H2 evolution. In addition, no particles were observed by dynamic light scattering (DLS) measurements in the reaction solution with 1 after 4 h of irradiation. Moreover, no influence on the catalytic activity of 1 was observed with the addition of a drop of Hg (Figure S14). The results also support the homogeneous catalysis of 1 under the present photocatalytic conditions. Photocatalytic H2 evolution with 1 under several different conditions was also tested to optimize the catalytic conditions. A MeCN/ascorbate buffer solution was employed in the photocatalytic reaction, and only a little amount of H2 was E


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Inorganic Chemistry Detection of One-Electron Reduced Ni Intermediates. The reduction of the Ni(II) complexes by [Ru(bpy●−)(bpy)2]+ is expected to produce one-electron reduced Ni species as observed in the electrochemical measurements, in which the one-electron reduction step is quasi-reversible (Figure 2b). Thus, we performed electron spin resonance (ESR) measurements to detect the Ni(I) state of 1 in the S = 1/2 spin state, which is ESR active, by using CoII(Cp)2 (Eox = −1.34 V vs Fc/ Fc+)40,48 as a reductant. CoII(Cp)2 (1 equiv) was added to 1 (1.0 mM) in MeCN under Ar, and the mixture solution was investigated by ESR spectroscopy at 150 K. As a result, we observed an ESR signal at giso = 2.15 with a peak-to-peak separation of 150 G (Figure S19a), assignable not to an organic radical but to a Ni(I) species.49 A characteristic absorbance at 520 nm was also observed by UV−vis spectroscopy for the Ni(I) species (Figure S19b), which was reminiscent of that reported for a reported Ni(I) complex coordinated by a similar ligand with two thioether sulfur atoms and two pyridines.50 The results suggest the formation of a Ni(I) species from the Ni(II) complex (1) in the 1e− chemical reduction and also confirm the possibility of the formation in the photocatalytic reaction with [Ru(bpy●−)(bpy)2]+ as a reductant. The observation of the one-electron reduced species, the Ni(I) intermediate, also allows us to conclude that the first reduction step is not ratelimiting. Kinetics Study on the Photocatalytic H2 Evolution. Reductive quenching of 3{[Ru(bpy)3]2+}* by NaAsc to generate [Ru(bpy●−)(bpy)2]+ as a responsible reductant was scrutinized in DMA/H2O (1:1 v/v) at 298 K. The rate constant of the electron transfer from NaAsc to 3{[Ru(bpy)3]2+}* was determined to be 8.2 × 108 M−1 s−1 from a Stern−Volmer plot (Figure S20) on the basis of the reported lifetime of the excited state in DMA/H2O mixed solvents (0.80 μs).51 The rate of electron transfer from [Ru(bpy●−)(bpy)2]+ to a Ni(II) complex was also determined by nanosecond transient absorption spectroscopy in a deaerated DMA/ascorbate buffer solution (pH 4; 1:1 v/v) upon photoexcitation at 450 nm (Figure 7). The transient absorption due to [Ru(bpy●−)(bpy)2]+ was observed at 530 nm (Figure S21) in the absence of the Ni catalysts. In the presence of 1 or 2, the decay of the absorbance derived from [Ru(bpy●−)(bpy)2]+ became faster than that in the blank test because of occurrence of electron transfer from [Ru(bpy●−)(bpy)2]+ to 1 or 2 as shown in Figure 7a,b, respectively. The decay rate constants increased linearly with increasing concentrations of Ni(II) complexes. The secondorder rate constants of electron transfer from [Ru(bpy●−)(bpy)2]+ to the Ni complexes were determined to be 3.1 × 106 M−1 s−1 (1) and 4.2 × 106 M−1 s−1 (2), respectively, from the slopes of dependence of the pseudo-first-order decay rate constants on concentrations of the Ni(II) complexes (Figure 7c,d). The ET rate constants determined for 1 and 2 are comparable, since the reduction potentials of both complexes are comparable, as can be seen in Figure 2. Rate-Determining Step of the Photocatalytic Reaction. Compared with the initial H2 evolution rate of 1 (9.1 μmol h−1) shown in Figure 6, the two electron-transfer processes occur expeditiously as demonstrated by the Stern− Volmer plot (7.8 × 108 M−1 s−1) for the formation of [Ru(bpy●−)(bpy)2]+ and the decays of the transient absorption (3.1 × 106 M−1 s−1 (1) and 4.2 × 106 M−1 s−1 (2)) for the reduction of the Ni(II) catalysts by [Ru(bpy●−)(bpy)2]+. These facts strongly suggest that the ET reactions to form the putative NiII−H intermediates are not rate-limiting, but one of the

Figure 7. Time profiles of absorbance at 530 nm due to the decay of [Ru(bpy●−)(bpy)2]+ in the presence of various concentrations of (a) 1 (red ●: 0.89 mM, blue ●: 2.2 mM, and green ●: 4.4 mM) and (b) 2 (red ●: 0.70 mM, blue ●: 1.4 mM, green ●: 3.1 mM, and pink ●: 6.1 mM) in a DMA/ascorbate buffer (pH 4, 4.0 mL, 1:1 v/v) containing [Ru(bpy)3]2+ (0.10 mM) upon photoexcitation at 450 nm. Plots of kobs against concentration of (c) 1 and (d) 2.

subsequent steps related to the H2 evolution should be the ratedetermining step. In the present case, the most plausible ratedetermining step is the reaction of the NiII−H intermediates with proton as discussed above. This proposal is also discussed on the basis of DFT calculations in the following section. DFT Calculations to Gain Mechanistic Insights into H2 Evolution. We performed DFT calculations to elucidate the mechanism of H−H bond formation through the reaction between the putative Ni−H intermediate and a proton of an oxonium ion (H3O+) to release H2 by taking solvent effects into account with the polarizable continuum model (water, ε = 78.3). Figure 8 shows DFT-optimized structures of the putative Ni−H complexes derived from 1 and 2, a plausible transition state, and Ni(II) species formed after H2 release together with their relative energies (see also Tables S1−S6). The activation barriers (G‡ at 298 K) to afford the transition states, [NiII−H··· (H3O+)]2+, were calculated to be 6.7 kcal mol−1 for 1 and 9.3 kcal mol−1 for 2. The values of activation barriers are consistent with previously reported computational results.52 Our results suggest that the reaction of H3O+ and [NiII(H)(L)(H2O)]+ occurs faster for 1 than for 2. The trend is consistent with the results of the electrocatalytic and photocatalytic H2 evolution as shown in Figures 4 and 6 and Table 1. In the course of the calculations on the transition state, facile coupling between the hydrido ligand and the proton of H3O+ was observed at 5.1 Å (1) and 4.4 Å (2) of atomic separations as depicted in Figure 8. This result suggests that the H−H bond formation proceeds easily through electrostatic attraction between the negatively charged hydrido ligand and the positively charged proton as the driving force. To examine the possibility of dissociation of a pyridine nitrogen atom in the putative Ni(II)-hydride intermediates, we conducted DFT calculations on plausible Ni(II)-hydrido intermediates with or without dissociation of one pyridine arm involving one water molecule. The results indicate that the η4-structures with one remote water molecule are more stable F


Inorganic Chemistry

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CONCLUSIONS We have demonstrated that two Ni(II) complexes with tetradentate S2N2 ligands 1 and 2 can catalyze H2 evolution in the electrochemical and photochemical setups. Complex 1 with the five-membered chelate ring exhibited higher catalytic activity than that of 2 with the six-membered chelate ring in the catalytic systems. The electrochemical and ESR measurements clearly demonstrated that the reduction of the Ni(II) complexes proceeds stepwise; the first step is a reversible 1e−-reduction reaction, and the second one is an irreversible 1e−/1H+ PCET reaction. The homogeneous electrocatalytic systems using 1 and 2 as catalysts in ascorbate buffer (pH 4.0) showed high Faraday efficiency in H2 evolution. The ratedetermining step was clarified to be Ni−H bond cleavage by the reaction with proton in the electrochemical H2 evolution on the basis of kinetic study demonstrating normal KIE (4.6 for 1). Under the photochemical conditions with the visible light irradiation at 450 nm, we also observed normal KIE in the H2 production rates, indicating that the Ni−H bond cleavage should also be involved in the rate-limiting step. A sequence of electron-transfer reactions proceeded in much faster rates to reduce the Ni(II) complexes than the rates of H2 evolution in the photocatalytic H2 evolution under the same conditions, confirming that the rate-determining step in the H2 evolution using 1 and 2 is not the electron-transfer cascade but the reaction of the Ni−H intermediates with proton to cleave the Ni−H bonds. The reactivity difference between 1 and 2 in the H2 evolution was attributed to the difference in the activation barrier in the reaction of putative NiII−H intermediates and proton as suggested by DFT calculations. Concerning the structural impact on the reactivity of the catalysts, as seen in the crystal structures of 1 and 2, the structural tension derived from three consecutive five-membered chelate rings should raise the potential of the intermediate to make it more reactive in comparison with that derived from a more flexible five-six-fivemembered chelate sequence to afford a more relaxed and stable structure at the ground state. Finally, the present study provides valuable information on the mechanism of the H2 evolution using Ni(II) complexes as catalysts toward the development of efficient molecular catalytic systems for both electrochemical and photochemical H2 evolution.

Figure 8. A calculated Gibbs free energy diagram for the H2 evolution and optimized structures of Ni(II)-hydrido intermediates, transition states of reactions, and resultant species with H2 at the UB3LYP/6311+G** level of theory. Units are in angstroms and kilocalories per mole.

(11.4 kcal mol−1 for both Ni(bpet) and Ni(bppt) than the corresponding η3-structures including one aqua ligand and one dissociated pyridine arm (Figure S22). Thus, we concluded that the Ni(II)-hydrido intermediates should have η4-bpet or η4bppt ligand without dissociation of the pyridine arm as depicted in Figure 8. As mentioned above, the comparable rate constants of electron transfer from [Ru(bpy●−)(bpy)2]+ to 1 and 2 can be explained on the basis of the fact that the reduction potentials of 1 and 2 exhibit no significant difference as depicted in Figure 2, indicating virtually the same driving force of ET for both complexes in the reduction to generate Ni(0) species. Thus, the reactivity difference between 1 and 2 in H2 evolution can be attributed to the Ni−H bond cleavage step, rather than the formation of Ni−H intermediates. The normal KIE values of the photocatalytic H2 evolution for 1 (2.2) and 2 (1.4) also support that Ni−H bond-cleaving step, not the Ni−H bondforming step, should be involved in the rate-determining step. The calculated activation energy of 1 in Ni−H bond cleavage is smaller than that of 2, as depicted in Figure 8 resulting the higher activity of 1 in the photocatalytic H2 evolution. On the basis of the experimental and theoretical results, the activity difference between 1 and 2 can be ascribed to the Ni−H bond cleavage step to react with a proton. The stability difference of the Ni−H intermediates between 1 and 2 could also affect the activity difference. According to the results of DFT calculations, the reaction of the Ni−H intermediate with proton (H3O+) occurs smoothly due to strong electrostatic attraction (dotted line in Figure 8) between a negatively charged hydrido ligand and a positively charged proton in certain separation (5.1 Å for 1 and 4.4 Å for 2) to release H2. Thus, the driving force of the H−H bond formation can be attributable to the electrostatic attractive force.

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EXPERIMENTAL SECTION

General. Chemicals and solvents were purchased from commercial sources and used as received unless otherwise mentioned. All synthetic reactions were performed under an argon atmosphere. [Ru(bpy)3]Cl2 was synthesized and purified based on a reported method.53 1H NMR measurements were performed on JEOL EX 400 and Bruker AVANCE 400 spectrometers. Electrospray ionization time-of-flight mass spectrometry (ESI-TOF-MS) spectra were measured on a JEOL JMS-T100CS spectrometer. Electrocatalytic measurements were performed on an ALS/CH Instruments Electrocatalytic Analyzer model 660A. SEM-EDX measurements were performed on Hitachi S4800 with an EDX (EMAX EX-320). Synthesis of Bis(2-pyridylmethyl)-1,2-ethanedithiol (bpet).36 To absolute methanol (12 mL) was added Na (0.3 g, 13 mmol) under Ar. When all the sodium metal dissolved, 1,2-ethanedithiol (255 μL, 3.04 mmol) was added to the solution and stirred at room temperature for 10 min. Then, the solution of 2-picolyl chloride hydrochloride (1.06 g, 6.46 mmol) in absolute methanol (7 mL) was added slowly to the mixture. The mixture was stirred for 18 h at room temperature under Ar. After removal of the solvent, the residue was extracted with chloroform and H2O. The organic layer was dried over anhydrous Na2SO4. The volatiles of the filtrate were evaporated and dried in vacuo to give pale yellow oil (0.78 g, 2.83 mmol, 93% yield). 1H NMR G


Article

Inorganic Chemistry (CDCl3): δ 2.67 (s, 4H), 3.81 (s, 4H), 7.13 (dd, J = 4.0, 7.8 Hz, 2H), 7.33 (d, J = 7.8 Hz, 2H), 7.62 (t, J = 7.8 Hz, 2H), 8.49 (d, J = 4.0 Hz, 2H). MS (MeOH): m/z = 299.02 (calcd. for [bpet + Na]+: 299.07). Synthesis of Bis(2-pyridylmethyl)-1,3-propanedithiol (bppt).36 To absolute methanol (15 mL) was added Na (0.23 g, 10 mmol) under Ar. When all the sodium metal was dissolved, 1,3propanedithiol (199 μL, 1.98 mmol) was added and stirred at room temperature for 10 min. Then, the solution of 2-picolyl chloride hydrochloride (0.65 g, 3.97 mmol) in absolute methanol (10 mL) was added slowly to the mixture. The mixture was stirred for 18 h at room temperature under Ar. After removal of the solvent, the residue was extracted with chloroform and H2O. The organic layer was dried over anhydrous Na2SO4. The volatiles of the filtrate were evaporated and dried in vacuo to give pale yellow oil (0.56 g, 1.91 mmol, 96% yield). 1 H NMR (CDCl3): δ 1.80 (q, J = 7.2 Hz, 2H), 2.53 (t, J = 7.2 Hz, 4H), 3.79 (s, 4H), 7.13 (dd, J = 4.0, 7.8 Hz, 2H), 7.33 (d, J = 8.0 Hz, 2H), 7.62 (t, J = 7.8 Hz, 2H), 8.50 (d, J = 4.0 Hz, 2H). MS (MeOH): m/z = 313.02 (calcd. for [bppt + Na]+: 313.08). Synthesis of [NiII(bpet)(MeCN)2](ClO4)2 (1·(ClO4)2).34 A solution of Ni(ClO4)2·6H2O (0.88 g, 2.4 mmol) in methanol (50 mL) was added to a solution of bpet (0.66 g, 2.4 mmol) in methanol (20 mL), and the mixture was stirred for 1 d at room temperature under Ar. The resultant blue precipitate was filtered and washed with MeOH. Recrystallization of the solid obtained from acetonitrile/ethyl acetate gave purple crystals (1.00 g, 1.62 mmol, 68% yield). 1H NMR (CD3CN): δ 9.88, 16.19, 46.00, 56.41, 81.77. Anal. Calcd for C14H20Cl2N2NiO10S2: C 29.50, H 3.54, N 4.91, found: C 29.74, H 3.33, N 4.96%. MS (MeCN): m/z = 432.93 (calcd. for [1 − 2MeCN + ClO4]+: 432.96). Synthesis of [NiII(bppt)(MeCN)2](BPh4)2 (2·(BPh4)2). A solution of Ni(ClO4)2·6H2O (0.88 g, 2.40 mmol) in methanol (50 mL) was added to a solution of bppt (0.66 g, 2.38 mmol) in methanol (20 mL), and the mixture was stirred for 1 d at room temperature under Ar. The mixture was evaporated and dried in vacuo. Sticky residue was dissolved in methanol, and a methanol solution of NaBPh4 was added to the solution to afford a purple solid. Recrystallization of the solid obtained from acetonitrile/ethyl acetate gave purple crystals (1.05 g, 0.98 mmol, 63% yield). 1H NMR (CD3CN): δ −22.47, 16.46, 46.75, 54.58, 121.07. Anal. Calcd for C72H76B2N4NiO3S2: C 72.68, H 6.44, N 4.71, found: C 72.75, H 6.50, N 4.65%. MS (MeCN): m/z = 424.98 (calcd. for [2 − 2MeCN + Ph]+: 425.07). X-ray Crystallography on 2. X-ray diffraction data were collected for 2 at 120 K on a Bruker APEXII Ultra diffractometer. The structure was solved by a direct method (SIR-97)54 and expanded with a differential Fourier technique. All non-hydrogen atoms were refined anisotropically, and the refinement was performed with full-matrix least-squares on F. All calculations were performed using the YadokariXG crystallographic software package.55 In the structure refinements, the contribution of severely disordered solvent molecules of crystallization was subtracted from the diffraction pattern by the “Squeeze” program.56 CCDC 1581124 contains the supplementary crystallographic data for 2. Preparation of Ni-Complexes-Loaded Carbon Paper (CP) Electrodes. CP electrodes loading complex 1 or 2 were prepared by a drop-casting method as follows. A certain amount of 5% alcoholic Nafion solution (50 μL, Sigma-Aldrich Co. LLC.) was added to an MeCN (2.0 mL) solution of a Ni(II) complex (1.0 mM), and the solution was sonicated for 30 min. Each 5.0 μL of the solution was taken and drop-casted on a CP (effective area: 1.0 cm × 1.0 cm), until the desired amount of the complex was loaded, and the CP was dried in vacuo. Electrocatalytic H2 Evolution. Bulk electrolysis was performed in a two-compartment electrocatalytic cell (42 mL) bridged by a Nafion perfluorinated membrane (Sigma-Aldrich Co. LLC.), where a CP was used as a working electrode, a Ag/AgNO3 electrode for nonaqueous solutions, a Ag/AgCl electrode for aqueous solutions as a reference electrode, and a Pt mesh as a counter electrode. Homogeneous bulk electrolysis was performed in an Ar-saturated mixed solvent of MeCN/ H2O (9:1 v/v, 15 mL) containing 1 and 2 (1.0 mM) and tetrabutylammonium hexafluorophosphate (TBAPF6; 0.10 M) as an

electrolyte at 298 K. Heterogeneous bulk electrolysis was performed in an Ar-saturated buffered solution (0.10 M, ascorbate buffer for pH 4 and pH 5, phosphate buffer for pH 7) containing 1 and 2 (1.0 mM) at 298 K. Typically, a 10 × 10 mm piece of CP was immersed in the solution. The evolved gas in the headspace was sampled by a gastight syringe (100 μL) and quantified by a Shimadzu GC-2014 gas chromatograph (GC: Ar carrier, a packed column with activated charcoal (3.0 mm × 3.0 m, 60−80 mesh) at 353 K) equipped with a thermal conductivity detector (TCD). Photocatalytic H2 Evolution. A mixed solution (4.0 mL) of DMA/ascorbate buffer (pH 4.0, 1:1 v/v) containing 1 or 2 (0.10 mM) and [Ru(bpy)3]Cl2 (0.25 mM) was flushed with Ar gas for 20 min. The solution was then irradiated with a 6 W light-emitting diode (LED) lamp (λ = 450 nm). After each reaction time (15, 30, 60, 120, 180, and 240 min), 100 μL of gas in the headspace of the reaction vessel was taken by using a gastight syringe and analyzed by GC to quantify H2 evolved. Quantum Yield Determination. A mixed solution (4.0 mL) of DMA/ascorbate buffer (pH 4.0; 1:1 v/v) containing 1 (0.10 mM) and [Ru(bpy)3]Cl2 (0.25 mM) was irradiated with a 6 W LED lamp (λ = 450 nm). The quantum yield (QY) of the photochemical H2 evolution under the conditions was calculated according to the following equation: QY(%) = {(2 × R/I) × 100}, where R (mol s−1) is the H2 evolution rate, and I (einstein s−1) is the rate of photon flux of the incident light. As two electrons produce one molecular hydrogen, two photons are required for this system. The total number of incident photons was determined by a standard method using an actinometer (potassium ferrioxalate, K3[FeIII(C2O4)]3) in H2O at room temperature. The rate of photon flux of the incident light (I) was determined to be 6.02 × 10−8 einstein s−1 under photoirradiation with the LED lamp at 450 nm. DLS Measurements. Dynamic light scattering (DLS) was measured on an Otsuka Electronics model FDLS-3000. The reaction solution (4.0 mL) of DMA/ascorbate buffer (pH 4.0; 1:1 v/v) containing 1 (0.10 mM) and [Ru(bpy)3]Cl2 (0.50 mM) under Ar atmosphere before and after photoirradiation was diluted by 10 times with DMA for the DLS measurements. Absorption and Emission Spectroscopies. UV−Vis spectroscopy was performed on Shimadzu UV-3600 and UV-2450 spectrometers at room temperature using quartz cells (light pass length = 1 cm). Emission spectroscopy was performed on a HORIBA FluoMax-4 spectrofluorometer at room temperature using quartz cells (light pass length = 1.0 cm). ESR Measurements. ESR spectra were taken on a Bruker X-band spectrometer (EMXPlus9.5/2.7) with a liquid nitrogen transfer system under nonsaturating microwave power conditions (3.0 mW), operating at 9.535 GHz. The magnitude of the modulation was chosen to optimize the resolution and the signal-to-noise (S/N) ratio of the observed spectrum (modulation amplitude, 10 G; modulation frequency, 100 kHz). CoII(Cp)2 (1 equiv) was added to 1 (2.0 mM) in MeCN under Ar. The solution was immediately frozen for the measurement at 150 K. Nanosecond Laser Flash Photolysis. Laser flash photolysis experiments were performed using a Panther OPO pumped by Nd:YAG laser (Continuum, SLII-10, 4−6 ns fwhm) at λ = 450 nm with the power of 0.5 mJ per pulse. The transient absorption measurements were made using a continuous-wave xenon lamp (150 W) and a photomultiplier (Hamamatsu 2949) as a probe light and a detector, respectively. The output from a photomultiplier was recorded on a digitizing oscilloscope (Tektronix, TDS3032, 300 MHz). Computational Method. We optimized local minima on the potential energy surfaces using the B3LYP method.57,58 For the Ni atom, we used the Wachters−Hay basis sets,59 and for the H, C, N, and O atoms, we used the 6-311+G** basis set.60 The harmonic vibrational frequencies were then calculated at the same level, which was used to estimate the zero-point energy (ZPE) and thermal contributions to the solvation Gibbs free energies (G) of the reaction components using the self-consistent reaction field (SCRF) calculations.61 Vibrational frequency calculations were performed to H


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Inorganic Chemistry confirm the transition states and local minima. There is no constraint to optimize stable structures and transition states. The thermal corrections to G were calculated for standard conditions (T = 298 K and p = 1.0 atm). The dielectric constant in the PCM calculations was set to ε = 78.3 to simulate the water solvent medium. All SCRF calculations were performed with Gaussian 09.62

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(5) Elgrishi, N.; McCarthy, B. D.; Rountree, E. S.; Dempsey, J. L. Reaction pathways of hydrogen-evolving electrocatalysts: electrochemical and spectroscopic studies of proton-coupled electron transfer processes. ACS Catal. 2016, 6, 3644−3659. (6) Kotani, H.; Ohkubo, K.; Takai, Y.; Fukuzumi, S. Viologenmodified platinum clusters acting as an efficient catalyst in photocatalytic hydrogen evolution. J. Phys. Chem. B 2006, 110, 24047− 24053. (7) (a) Ozawa, H.; Haga, M.; Sakai, K. A Photo-hydrogen-evolving molecular device driving visible-light-induced EDTA-reduction of water into molecular hydrogen. J. Am. Chem. Soc. 2006, 128, 4926− 4927. (b) Masaoka, S.; Mukawa, Y.; Sakai, K. Frontier orbital engineering of photo-hydrogen-evolving molecular devices: a clear relationship between the H2-evolving activity and the energy level of the LUMO. Dalton Trans. 2010, 39, 5868−5876. (c) Whang, D. R.; Sakai, K.; Park, S. Y. Highly efficient photocatalytic water reduction with robust iridium(III) photosensitizers containing arylsilyl substituents. Angew. Chem., Int. Ed. 2013, 52, 11612−11615. (8) Morales-Guio, C. G.; Stern, L. A.; Hu, X. Nanostructured hydrotreating catalysts for electrochemical hydrogen evolution. Chem. Soc. Rev. 2014, 43, 6555−6569. (9) Wang, C.; deKrafft, K. E.; Lin, W. Pt nanoparticles@photoactive metal−organic frameworks: efficient hydrogen evolution via synergistic photoexcitation and electron injection. J. Am. Chem. Soc. 2012, 134, 7211−7214. (10) Greeley, J.; Jaramillo, T. F.; Bonde, J.; Chorkendorff, I. B.; Norskov, J. K. Computational high-throughput screening of electrocatalytic materials for hydrogen evolution. Nat. Mater. 2006, 5, 909− 913. (11) Yamada, Y.; Miyahigashi, T.; Kotani, H.; Ohkubo, K.; Fukuzumi, S. Photocatalytic hydrogen evolution under highly basic conditions by using Ru nanoparticles and 2-phenyl-4-(1-naphthyl)quinolinium ion. J. Am. Chem. Soc. 2011, 133, 16136−16145. (12) Maeda, K.; Sahara, G.; Eguchi, M.; Ishitani, O. Hybrids of a Ruthenium(II) polypyridyl complex and a metal oxide nanosheet for dye-sensitized hydrogen evolution with visible light: effects of the energy structure on photocatalytic activity. ACS Catal. 2015, 5, 1700− 1707. (13) Thoi, V. S.; Sun, Y.; Long, J. R.; Chang, C. J. Complexes of earth-abundant metals for catalytic electrochemical hydrogen generation under aqueous conditions. Chem. Soc. Rev. 2013, 42, 2388− 2400. (14) (a) Queyriaux, N.; Jane, R. T.; Massin, J.; Artero, V.; ChavarotKerlidou, M. Recent developments in hydrogen evolving molecular cobalt(II)−polypyridyl catalysts. Coord. Chem. Rev. 2015, 304−305, 3−19. (b) 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. (c) Straistari, T.; Fize, J.; Shova, S.; Rég lier, M.; Artero, V.; Orio, M. A thiosemicarbazone−nickel(II) complex as efficient electrocatalyst for hydrogen evolution. ChemCatChem 2017, 9, 2262−2268. (d) Hu, Z.; Shen, Z.; Yu, J. C. Phosphorus containing materials for photocatalytic hydrogen evolution. Green Chem. 2017, 19, 588−613. (15) Fukuzumi, S.; Hong, D.; Yamada, Y. Bioinspired photocatalytic water reduction and oxidation with earth-abundant metal catalysts. J. Phys. Chem. Lett. 2013, 4, 3458−3467. (16) (a) Yamada, Y.; Miyahigashi, T.; Kotani, H.; Ohkubo, K.; Fukuzumi, S. Photocatalytic hydrogen evolution with Ni nanoparticles by using 2-phenyl-4-(1-naphthyl)quinolinium ion as a photocatalyst. Energy Environ. Sci. 2012, 5, 6111−6118. (b) Hong, D.; Yamada, Y.; Sheehan, M.; Shikano, S.; Kuo, C.-H.; Tian, M.; Tsung, C.-K.; Fukuzumi, S. Mesoporous nickel ferrites with spinel structure prepared by an aerosol spray pyrolysis method for photocatalytic hydrogen evolution. ACS Sustainable Chem. Eng. 2014, 2, 2588−2594. (17) Kaur-Ghumaan, S.; Schwartz, L.; Lomoth, R.; Stein, M.; Ott, S. Catalytic hydrogen evolution from mononuclear iron(II) carbonyl complexes as minimal functional models of the [FeFe] hydrogenase Active Site. Angew. Chem., Int. Ed. 2010, 49, 8033−8036.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.8b00881. 1 H NMR spectra, MS spectra, CVs, UV−vis spectra, time courses of H2 evolution under various conditions, a setup for bulk electrolysis, ESR and UV−vis spectra of Ni(I) species, Stern−Volmer plots, nanosecond transient absorption spectra, DFT-optimized structures of possible intermediates, and Cartesian coordinates of DFToptimized structures (PDF) Accession Codes

CCDC 1581124 contains 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 data_ [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: [email protected]. ORCID

Dachao Hong: 0000-0003-0581-1315 Hiroaki Kotani: 0000-0001-7737-026X Tomoya Ishizuka: 0000-0002-3897-026X Kazunari Yoshizawa: 0000-0002-6279-9722 Shunichi Fukuzumi: 0000-0002-3559-4107 Takahiko Kojima: 0000-0001-9941-8375 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by Grants-in-Aid (Nos. 24655044, 24245011, 25107508, 15H00861 to T.K., 16H02268 to S.F., 16K05725 to Y. S., JP2410914 to K.Y.) from Japan Society for the Promotion of Science (JSPS). T.K. also appreciates financial supports from The Asahi Glass Foundation and The Mitsubishi Foundation. D.H. gratefully acknowledges support from JSPS by Grant-in-Aid (No. 15J04635) for JSPS fellowship for young scientists.

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REFERENCES

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