The rigidity and chelation effect of ligands on the hydrogen evolution reaction catalyzed by Ni(II) complexes

Abstract

With increasing interest in nickel-based electrocatalysts, three heteroleptic Ni(II) dithiolate complexes with the general formula [Ni(II)L(L′)₂] (1–3), L = 2-(methylene-1,1′-dithiolato)-5,5′-dimethylcyclohexane-1,3-dione and L′ = triphenylphosphine (1), 1,1′-bis(diphenylphosphino)ferrocene (DPPF) (2), and 1,2-bis(diphenylphosphino)ethane (DPPE) (3), have been synthesized and characterized by various spectroscopic techniques (UV-vis, IR, ¹H, and ³¹P{¹H} NMR) as well as the electrochemical method. The molecular structure of complex 2 has also been determined by single-crystal X-ray crystallography. The crystal structure of complex 2 reveals a distorted square planar geometry around the nickel metal ion with a NiP₂S₂ core. The cyclic voltammograms reveal a small difference in the redox properties of complexes (ΔE° = 130 mV) while the difference in the catalytic half-wave potential becomes substantial (ΔE_cat/2 = 670 mV) in the presence of 15 mM CF₃COOH. The common S^S-dithiolate ligand provides stability, while the rigidity effect of other ligands (DPPE (3) > DPPF (2) > PPh₃ (1)) regulates the formation of the transition state, resulting in the Ni^III–H intermediate in the order of 1 > 2 > 3. The foot-of-the-wave analysis supports the widely accepted ECEC mechanism for Ni-based complexes with the first protonation step as a rate-determining step. The electrocatalytic proton reduction activity follows in the order of complex 1 > 2 > 3. The comparatively lower overpotential and higher turnover frequency of complex 1 are attributed to the flexibility of the PPh₃ ligand, which favours the easy formation of a transition state.

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Introduction

The high consumption rate of fossil fuel has led to severe environmental pollution and its limited availability may cause a global energy crisis.¹ Therefore, the development of alternative sustainable and economical energy resources is among the major challenges. However, as an alternative, dihydrogen (H₂) is a sustainable and eco-friendly carbon-free fuel for future energy generation with a high energy density of 120–142 MJ kg⁻¹. Industrial large-scale H₂ production is mainly achieved by the steam-methane reforming process which affects energy efficiency because of the low purity of hydrogen products (owing to the presence of carbon-containing residues).² Water electrolysis is an efficient sustainable approach for the production of carbon-free H₂ by electrochemical and photochemical methods.³ However, the development of efficient electrocatalysts for the large-scale production of H₂ by electrochemical proton reduction using earth-abundant metals is still challenging.⁴ Platinum-based electrocatalysts are promising candidates amongst all electrocatalysts and demonstrate efficient reduction of protons to H₂. However, the high cost and low natural abundance limit the applicability of platinum-based electrocatalysts for large-scale H₂ production. Therefore, alternative cost-effective electrocatalysts containing transition metals exhibiting variable oxidation states and labile coordination geometry are required to be designed and synthesized.⁵ To obtain cost-effective electrocatalysts, metal complexes based on Ni,^6–8 Co,^9–11 Zn,¹² Cu,¹³ Mn,¹⁴ Fe,¹⁵ and Pd¹⁶ have been reported as electrocatalysts showing substantial hydrogen evolution reaction (HER) activity.¹⁷ Dithiolene-based metal complexes are another class of complexes that are conductive¹⁸ and electrocatalytically active,¹⁹ and exhibit non-linear optical properties.^20,21 The MN₂S₂ chromophore promotes the ease of electron transfer, and therefore facilitates the electrocatalytic reactions and the strong solution luminescence²² with solvatochromism²³ (indicative of coordination unsaturation with changes in coordination geometry).

The bidentate 1,1-dithiolate complexes with various metal ions, exhibiting square planar geometry, have been reported and extensively studied for structural variation, bonding ability, and redox chemistry with their application as photocatalysts and electrocatalysts for energy generation processes.²⁴ Ni-based dithiolate complexes exhibit promising H₂ production ability and act as molecular catalysts offering significant advantages in terms of ease of structural modifications to tune the stability, redox properties, and electrocatalytic activity.^25,26 DuBois and co-workers have employed internal proton relays in the second coordination sphere of Ni phosphine-based complexes to design extremely fast catalysts and demonstrated a Ni(N₂P₂) electrocatalyst with a turnover frequency (TOF) above 10⁵ s⁻¹ for H₂ production.²⁷ Crabtree et al.²⁸ synthesized Ni-containing metalloenzymes and investigated their proton reduction ability with a functional model of hydrogenase. Furthermore, Nocera et al.²⁹ reported the HER using Ni hangman porphyrins and unraveled the role of pendant proton relays and the proton-coupled electron transfer process. The Ni-porphyrin-based complexes offer relatively better electrocatalytic activity compared to the cobalt analogue. Recently, several Ni-based coordination materials and complexes, viz., NiN₄, NiS₄, NiN₂S₂, NiS₂O₂,^5,30 an N₅-pentadentate ligand system,³¹ Ni-porphyrin,³² bipyridine derivatives,^33,34 a Ni-sulfur-based radical ligand complex as a functional model of hydrogenase,³⁵ Ni-dithiolates,³⁶ Ni with pendant amine,³⁷ enzyme-based ligands,^7,38 Ni with macrocyclic ligands,³⁹ phosphinopyridyl,⁴⁰ corrole-chelated Ni complexes⁴¹ and so forth, have also been found to exhibit substantial electrocatalytic activity in proton reduction to molecular hydrogen.

The rational design of an efficient electrocatalyst generally depends on (i) the electronic effects induced by coordinated ligands which regulate the redox properties^5,6,41–43 and (ii) the steric or geometric effect induced by the coordinated ligands^6,44,45 which controls the mechanistic path and regulates the formation of the transition state, i.e., the energy barrier of the rate-determining step in electrocatalytic hydrogen production. Orio et al.⁴³ highlighted the electronic effects of different ligands on electrocatalytic hydrogen production using Ni-thiosemicarbazone-based complexes. Furthermore, Gross et al.⁴¹ reported corrole-chelated Ni-based complexes, wherein the low oxidation state with an extra negative charge is stabilized by the electron-withdrawing substituents present on the corrole ligand. This results in a relatively enhanced electrocatalytic reactivity toward the HER compared to the neutral Ni(II) porphyrin analogue. Furthermore, Roy et al.⁵ explored the impact of alkyl side chain lengths of ligands on the HER in mononuclear Ni(II)-based complexes. Very recently, Qiu et al.⁴⁴ reported bent bis(dipyrrin)Ni(II) complexes as an efficient electrocatalyst for the HER. Cao et al.⁴⁵ highlighted that the steric effects of substituents on ligands decide the mechanistic pathway (homolytic versus heterolytic) for the HER.

The earlier synthesis efforts were more focused on improving the electronic effects and, consequently, the redox properties of complexes, while focusing less on the geometric and steric effects of coordinated ligands on HER kinetics. Earlier, Jones et al.⁶ revealed that the S₂P₂-coordinated Ni-based distorted square planar chelated complex [Ni(bdt)(DPPF)] is a promising candidate for enhanced and efficient production of hydrogen, where bdt = 1,2-benzenedithiolate and DPPF = 1,1′-bis(diphenylphosphino)ferrocene. On the other hand, the other analogous compound [Ni(bdt)(DPPE)] was not efficient towards the HER, where DPPE = 1,2-bis(diphenylphosphino)ethane. The synthesis of stable Ni(II) phosphine-based complexes is still challenging and only a few complexes are reported.^6,46–50 Singh et al. have synthesized functionalized heteroleptic Ni(II) dithiolate-phosphine-based complexes and studied their ability to catalyse water oxidation (oxygen evolution reaction)⁴⁷ as well as investigated their photosensitizing properties⁵¹ and semiconductor behaviour.⁵²

In this study, we have synthesized and characterized three heteroleptic Ni(II) complexes with the general formula [Ni(II)L(L′)₂], where L = 2-(methylene-1,1′-dithiolato)-5,5′-dimethylcyclohexane-1,3-dione is a common component in all complexes, while L′ varies as PPh₃ (1), DPPF (2), and DPPE (3). The chelation effect of the ligand K₂L provides stability, while the other ligands, L′, are chosen to explore the steric or geometric effect on the HER activity of the complexes. Complex 3 with L′ = DPPE is considered as rigid, 2 with DPPF as semi-flexible, while 1 as flexible with two PPh₃ ligands. Furthermore, we have electrochemically investigated the redox and electrocatalytic properties of these complexes toward the HER in acidic media. Finally, we have investigated the contribution of electronic, chelation, and steric or geometric effects of coordinated ligands to the HER activity of complexes 1–3.

Experimental section

Materials and methods

All reactions were carried out in open air at ambient temperature and pressure. The solvents were purified by standard procedures and dried before use where necessary. Dichloromethane (DCM) and methanol solvents were dried over calcium chloride and calcium carbonate, respectively, and freshly distilled prior to use. All other chemicals were purchased from Sigma-Aldrich, with the highest purity available, and used as received without further purification. The melting points of the complexes were determined in open capillaries using a Gallanlamp apparatus. Thin-layer chromatography (TLC) was performed on Merck 60 F254 silica gel, precoated on an aluminum plate.

Instrumentation

¹H and ³¹P{¹H} NMR spectra were recorded in deuterated solvents (CDCl₃/DMSO-d₆) containing 0.1% TMS (tetramethylsilane) using a JEOL ECZ500 MHz FT NMR spectrometer. The NMR chemical shift (δ) values are reported in parts per million (ppm) downfield from internal standard TMS and coupling constant (J) values are given in hertz (Hz). The UV-vis measurements were performed on a Shimadzu UV-1800 spectrophotometer using quartz cuvettes with a 1 cm path length. Single crystal XRD analysis was performed using a Brucker APEX-II CCD diffractometer and the details are provided in the crystal discussion section. Electrochemical experiments were performed at room temperature using a glassy carbon (GC) electrode (2 mm diameter) as a working electrode, a platinum coil as a counter electrode, and Ag/AgNO₃ (10 mM) in CH₃CN containing 0.1 M TBAPF₆ as a reference electrode. Ferrocene is used as an internal reference. Prior to use, the working GC electrode was mechanically cleaned with alumina powder of 0.5- and 0.03-micron sizes, followed by sonication in ethanol and triple distilled water for 15 minutes. The Pt coil counter electrode was boiled in nitric acid at 45 to 50 °C temperature for 10 minutes followed by sonication in ethanol and triple distilled water for 15 minutes each.⁵³ The electrodes were properly dried and sealed in argon-filled test tubes. The 1 mM solution of complexes was prepared in de-aerated anhydrous CH₃CN containing 0.1 M tetra-butylammonium tetrafluoroborate (TBABF₄). The cyclic and linear scan voltammograms were recorded using the Metrohm Autolab M204 instrument in the presence of argon gas at 25 °C.

Synthesis and characterization of complexes

The heteroleptic Ni(II) dithiolate complexes 1–3 were prepared by adopting the following general procedure. The characterization data of complex 3 well matched with the reported literature data.⁴⁷

Synthesis of [bis((triphenylphosphine)-2-(methylene-1,1′-dithiolato)-5,5′-dimethylcyclohexane-1,3-dione)Ni(II)] (1). Triphenylphosphine (0.262 g, 0.998 mmol) was dissolved in 2 mL of freshly distilled DCM, giving a clear solution, and further dropwise addition of NiCl₂·6H₂O (0.118 g, 0.496 mmol prepared in methanol) resulted in the formation of a white precipitate. The reaction mixture was then stirred at room temperature for 1 h followed by dropwise addition of the potassium salt of the dithiolate ligand (0.146 g, 0.499 mmol). Further stirring for 12 h resulted in the formation of an orange solid residue. This solid residue was filtered off using a G-4 sintered funnel, washed with methanol (3 × 2 mL) and dried under vacuum. Yield: ∼64%. M.p. 163–166 °C. IR data (KBr, cm⁻¹): 1603 (ν_C=O), 1397 (ν_C=CS2), 1080 (ν_C–S). ¹H NMR (500 MHz, DMSO-d₆, δ ppm): 7.68–7.54 (30H, m, aromatic on phenyl), 2.10 (4H, s, (–CH₂–)₂), 0.90 (6H, s, (CH₃–)₂). ³¹P{¹H} NMR (202 MHz, DMSO-d₆, δ ppm): 43.16. UV-vis [CH₃CN, λ_max (nm), ε (M⁻¹ cm⁻¹)]: 340 nm (ε = 0.91 × 10⁴), 377 nm (0.33 × 10⁴ M⁻¹ cm⁻¹) and 490 nm (ε = 0.19 × 10⁴).

Synthesis of [1,1′-bis((diphenylphosphino)ferrocene-2-(methylene-1,1′-dithiolato)-5,5′-dimethylcyclohexane-1,3-dione)Ni(II)] (2). Complex 2 was synthesized following the procedure used in the synthesis of complex 1 by utilizing the phosphine ligand DPPF (0.100 g, 0.180 mmol), NiCl₂·6H₂O (0.042 g, 0.176 mmol) and the potassium salt of the dithiolato ligand (0.052 g, 0.177 mmol), resulting in the formation of an orange solid residue. This solid residue was filtered off using a G-4 sintered funnel, washed with methanol (3 × 2 mL) and dried under vacuum. Yield: ∼66%. M.p. 114 °C. IR data (KBr, cm⁻¹): 1603 (ν_C=O), 1397 (ν_C=CS2), 1088 (ν_C–S). ¹H NMR (500 MHz, CDCl₃, δ ppm): 7.87–7.31(m, 20H, Ar–H), 4.35–4.30 (d, 8H, Cp-H), 2.22 (s, 4H, –(CH₂)₂−), 0.93 (s, 6H, –(CH₃)₂). ³¹P{¹H} NMR (202 MHz, CDCl₃, δ ppm): 29.11. UV-vis [CH₃CN, λ_max (nm), ε (M⁻¹ cm⁻¹)]: 334 nm (ε = 1.33 × 10⁴) and 415 nm (ε = 0.64 × 10⁴).

Results and discussion

Synthesis and characterization

The syntheses of the potassium salt of the dithiolato ligand (K₂L) and heteroleptic Ni(II) complexes 1–3 are summarized in Scheme 1. The synthesis of the ligand began with the abstraction of the methylene acidic protons of 2-(methylene-1,1′-dithiolato)-5,5′-dimethylcyclohexane-1,3-dione derivatives using KOH, followed by an in situ reaction with CS₂ similar to the reported literature procedure.^51,52 The heteroleptic complexes 1–3 were prepared by the reaction of a mixture containing equivalent amounts of the potassium salt of the dithiolato ligand and NiCl₂·6H₂O in methanol with PPh₃, DPPF, and DPPE in DCM (Scheme 1), respectively. All complexes were obtained as auburn solids which were stable in air and moisture and melted in the temperature range of 110–190 °C. These complexes were characterized by UV-vis, FT-IR, and NMR spectroscopy techniques. Furthermore, the molecular structure of complex 2 was also determined by single crystal X-ray crystallography.

Scheme 1 Schematic illustration for the synthesis of complexes 1–3.

Electronic absorbance studies

UV-vis absorption spectra of the potassium salt of the dithiolate ligand (K₂L) in methanol and complexes 1–3 in acetonitrile were collected at room temperature at 10⁻⁴ molar concentration. The UV-vis spectra are displayed in Fig. 1. The electronic absorption spectrum of the ligand displays two bands at ∼281 nm (ε = 1.42 × 10⁴ M⁻¹ cm⁻¹) and 385 nm (ε = 0.19 × 10⁴ M⁻¹ cm⁻¹) arising from the intra-ligand charge transfer (ILCT) transitions. The absorption bands of the heteroleptic complexes 1, 2 and 3 are observed at 340 nm (ε = 0.91 × 10⁴ M⁻¹ cm⁻¹), 377 nm (0.33 × 10⁴ M⁻¹ cm⁻¹) and 490 nm (ε = 0.19 × 10⁴ M⁻¹ cm⁻¹); 334 nm (ε = 1.33 × 10⁴ M⁻¹ cm⁻¹) and 415 nm (ε = 0.64 × 10⁴ M⁻¹ cm⁻¹); and 295 nm (ε = 1.03 × 10⁴ M⁻¹ cm⁻¹), 343 nm (ε = 1.78 × 10⁴ M⁻¹ cm⁻¹) and 380 nm (ε = 1.32 × 10⁴ M⁻¹ cm⁻¹), respectively. The higher energy absorptions in the ligand and complexes are assigned to the intra/interligand charge transfer or metal-perturbed intraligand charge transfer (ILCT).^54,55 However, in the case of complex 2 containing 1,1′-bis(diphenylphosphino)ferrocene, the absorption band near 415 nm cannot be unequivocally assigned to the presence of the ferrocene moiety.^56,57 Furthermore, the low energy absorption band at 490 nm in complex 1 is associated with the metal to ligand charge transfer (MLCT) from nickel to phosphorus and nickel to sulfur.⁶

Fig. 1 UV-vis spectra of ligand (K₂L) and complexes 1–3 in acetonitrile solution at 10⁻⁴ molar concentration.

IR and NMR spectral studies

The IR spectra of complexes 1 and 2 show stretching frequencies for C=O, C=CS₂, and CS₂ functional groups at ∼1603, ∼1397, and ∼1080 cm⁻¹, respectively. In addition, the presence of the Cp group in the DPPF ligand results in an absorption band at 482 cm⁻¹. These stretching frequency values are in line with earlier reported values.^{47,51,52,58,59} IR spectra of complexes 1 and 2 are shown in Fig. S1 and S2, ESI.†

All complexes are well characterized by NMR spectral studies and depicted in Fig. S3–S6, ESI.† The ¹H NMR spectra of complexes showed resonance signals at δ ∼1.00 and ∼2.00–2.23 ppm attributed to the two methyl (–CH₃) and methylene protons of the dithiolate ligand, respectively. The aromatic protons of the PPh₃, DPPF, and DPPE groups shifted to the downfield region, observed at δ ∼7.20–7.93 ppm, which indicates the bonding of the phosphine ligand to the metal center. The signals for the cyclopentadienyl protons in complex 2 were shifted downfield (δ 4.35 and 4.30 ppm) compared to free DPPF (δ 4.24 and 3.95 ppm). This confirms the coordination of the DPPF ligand with the Ni(II) center, which decreases the electron density at the cyclopentadienyl rings of the ferrocene moiety.

We have also recorded the ³¹P{¹H} NMR spectra of 1 and 2 which showed a signal for phosphorus at chemical shift values of 43.16 and 29.11 ppm, respectively. This indicated the symmetrical bidentate bonding behavior of the DPPF ligand with a Ni(II) center, whereas both PPh₃ ligands in complex 1 are equivalent.

Single crystal X-ray structure

The single crystals of complex 2 suitable for X-ray crystallographic studies were obtained by slow evaporation of the solvent from a dichloromethane solution of the complex layered with methanol. The crystallographic data and other structure refinement details are provided in Table S1,† whereas selected bond distances and bond angles are listed in Table 1. The molecular structure of complex 2 is shown in Fig. 2.

Fig. 2 Molecular structure of complex 2. Solvent and hydrogen atoms are deleted for clarity. Thermal ellipsoids are shown at 50% probability.

Table 1 Selected bond lengths (Å) and bond angles (°) for complex 2

Bond length/Å		Bond angle/°
Ni(1)–S(1)	2.2009(5)	S(1)–Ni–S(2)	76.93(2)
Ni(1)–S(2)	2.2105(6)	S(1)–Ni–P(1)	94.50(2)
Ni(1)–P(1)	2.2108(6)	S(1)–Ni–P(2)	161.95(2)
Ni(1)–P(2)	2.2082(6)	S(2)–Ni–P(1)	166.02(2)
S(1)–C(1)	1.733(2)	S(2)–Ni–P(2)	90.02(2)
S(2)–C(1)	1.7317(19)	P(2)–Ni–P(1)	100.54(2)
C(1)–C(2)	1.386(3)

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The Ni(II) metal center shows a distorted square planar coordination geometry and is surrounded by bidentate S^S-dithiolate and P^P-chelating DPPF ligands. The root mean square (r.m.s.) deviation of the four donor atoms (P1P2S1S2) in the equatorial plane from the least-squares plane is 0.204 Å, whereas the deviation of the nickel metal center from the plane is 0.0094(4) Å. The distorted square planar geometry around the nickel center has also been supported by a calculated geometrical index value (τ) of 0.227.⁶⁰ The Ni–S bond distances (2.2009(5) and 2.2105(6) Å) and Ni–P distances (2.2108(6) and 2.2082(6) Å) are well within the range expected for square planar Ni complexes.^51,52,61 Similarly, the distance of iron from nickel, 4.157 Å, is within the expected range. The S(1)–Ni–P(1) and S(2)–Ni–P(2) angles show less deviation from 90° than the S(1)–Ni–S(2) and P(1)–Ni–P(2) angles. It is interesting to note the S(1)–Ni–S(2) bite angle is 76.93(2)°, which is significantly smaller than the ideal value of 90° due to the formation of a strained four-membered chelate ring by the 1,1-dithiolate ligand. The S–Ni–S bite angle reported for complex 3 is 77.80(3)°.⁴⁷ Another notable structural parameter in complex 2 is the P(1)–Ni–P(2) bond angle of 100.54(2)° which is remarkably larger than the P–Ni–P bond angle (87.76(3)°) reported in complex 3. This could be attributed to the fact that the DPPF ligand is known to be highly sterically demanding as compared to the DPPE ligand. The C(1)–S(1) and C(1)–S(2) bond distances are 1.733(2) and 1.7317 (19) Å in complex 2. These distances are shorter than that reported for a typical C–S single bond length (1.81 Å) due to partial π-conjugation delocalization over the S–C–S group. The C1–C2 bond length in 2 (1.386(3) Å) is well in conformity with the C(sp²)=C(sp²) distance. The C–S and C1–C2 bond distances are very similar to the distances reported for complex 3.

The crystal structure of complex 3 was earlier reported by Singh et al.⁴⁷ Similarly, complex 1 is more likely to have a distorted square planar geometry due to the bidentate chelation effect of the dithiolato ligand. The distorted square planar geometry was earlier reported by Bellis et al.⁶² for an analogous series of complexes having Ni with (PPh₃)₂ and sulfonyldithiocarbimate as a bidentate chelating ligand.

The phase purity of the bulk sample of complex 2 was also confirmed by a comparison of the experimental powder X-ray diffraction (PXRD) patterns with the simulated PXRD patterns calculated from single crystal X-ray data. We observed that both the PXRD patterns (experimental and simulated) match well, indicating the phase purity of the bulk samples (Fig. S7, ESI†).

Electrochemical studies

Redox properties of synthesized complexes. The redox behaviour of all synthesized complexes 1–3 were electrochemically investigated by recording the cyclic voltammograms (CVs) using a glassy carbon (GC) electrode in nitrogen-degassed acetonitrile (CH₃CN) containing 0.1 M tetrabutylammonium tetrafluoroborate (TBABF₄) in the potential range of −2.5 V to +1.0 V versus Fc^+/0 (Fig. 3a). The CVs of complexes 1–3 exhibited quasi-reversible redox processes of Ni^II(L)/Ni^I(L) with cathodic peak currents (I_pc) appearing at potentials (E_pc) of −1.26, −1.32 V, and −1.39, while the anodic peak currents (I_pa) at potentials (E_pa) of −1.06, −1.14, and −1.18 V versus Fc^+/0, respectively (Table 2). Considering the reduction sweep, after the Ni^II(L)/Ni^I(L) reduction wave, a broad peak is observed near −1.8 and −2.0 V versus Fc^+/0 for complexes 1 and 2, respectively. These are attributed to the electrochemically quasi-reversible reductive process associated with the coordinated K₂L ligand, while in the anodic scan of 1–3, the peak observed near −0.2 to −0.5 V is due to the electrochemical oxidation of the K₂L ligand (Fig. S8, ESI†). Furthermore, a quasi-reversible redox wave is also observed for complex 2 near 0.7 V with a prominent oxidation peak of the DPPF ligand.⁶³

Fig. 3 Cyclic voltammograms (CVs) of complexes 1 (blue), 2 (purple), and 3 (red): (a) 1 mM 1, 2, and 3 with 0.1 M TBABF₄ in acetonitrile at a 200 mV s⁻¹ scan rate, (b) 1 mM 1, 2, and 3 with 15 mM TFA in acetonitrile containing 0.1 M TBABF₄ at a 200 mV s⁻¹ scan rate, and (c) the linear scan voltammograms (LSVs) at 50 mV s⁻¹.

Table 2 Comparison of the redox properties of complexes 1–3 (1 mM) and 0.1 M TBABF₄ in acetonitrile at a scan rate of 200 mV s⁻¹

Complex	E pc/V (Ni^II/Ni^I)	E pa/V (Ni^I/Ni^II)	ΔEp/V	E°/V		D cat (cm^2 s^−1)
1	−1.26	−1.06	0.20	−1.16	0.56	2.27 × 10^−7
2	−1.32	−1.14	0.18	−1.23	0.54	5.29 × 10^−7
3	−1.39	−1.18	0.21	−1.29	0.42	—

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The varying electronic properties and the ligand-induced geometric distortions are likely to influence the reaction intermediate energy at the transition state, consequently regulating the redox behavior of the complexes. The magnitude of the reduction potential for Ni^II(L)/Ni^I(L), i.e., E_pc of these complexes, follows the trend 3 > 2 > 1 (Table 2). The comparatively lower magnitude of the reduction potential of complex 1 is due to the poor sigma donor ability of the PPh₃ ligand compared to that of DPPF (2) and DPPE (3). Hence, the lowest unoccupied molecular orbital (LUMO) is less destabilized in complex 1, making it more readily accept electrons at a less negative reduction potential.

The linear correlation of peak currents of complexes 1 and 2 with the square root of scan rates revealed the diffusion-controlled electron transfer behavior (Fig. S9, ESI†). The diffusion coefficients for 1 and 2 were estimated as 2.27 × 10⁻⁷ and 5.29 × 10⁻⁷ cm² s⁻¹, respectively, from the slope of the peak current versus square root of scan rate plot and using eqn (S1) (ESI†). The characteristic peak currents and the corresponding peak potentials for 1–3 are listed in Table 2. Furthermore, to establish the best and most efficient electrocatalyst for the electrochemical proton reduction out of three complexes 1–3, CVs and linear scan voltammogram (LSVs) were recorded using 15 mM trifluoroacetic acid (TFA) containing 1 mM 1–3 with 0.1 M TBABF₄, as shown in Fig. 3b and c, respectively. The experimental observation revealed the expected phenomena (vide infra) for 1 and 2 because of the strong reducing ability of these complexes, while no such peak is observed for 3. Hence, complexes 1 and 2 were further extensively evaluated as electrocatalysts for electrochemical hydrogen generation.

Electrocatalytic hydrogen production by complexes 1 and 2. The performance of these heteroleptic mononuclear (1) and binuclear (2) complexes as effective electrocatalysts for HER application was tested using TFA as a proton source. To determine the activity of 1 and 2 as electrocatalysts, initially, CVs and LSVs of 1 and 2 were recorded in the presence of TFA (pK_a = 12.7) in CH₃CN, as shown in Fig. 4. The addition of TFA (5 mM to 25 mM) triggers the catalytic process, resulting in the appearance of a new catalytic peak with half-wave potentials (E_cat/2) at −1.25 and −1.44 V vs. Fc^+/0 for 1 and 2, respectively. The intensity of the new catalytic peak became more significant with the subsequent addition of TFA (Fig. 4). This rise in catalytic peak current represents the generation of H₂ from the electrocatalytic reduction of protons. The dependence of the catalytic peak current on the scan rate and acid concentration indicates a diffusion-limited electrocatalytic process (Fig. 4 and Fig. S10, ESI†). The catalytic peak potentials of 1 and 2 shifted towards more negative values with the sequential increase in the TFA concentration.

Fig. 4 Variation of the catalytic peak current of complexes 1 (a and b) and 2 (c and d) on increasing the concentration of TFA. CVs and LSVs were recorded for 1 mM 1 and 2 in CH₃CN containing 0.1 M TBABF₄ at a scan rate of 200 mV s⁻¹.

The onset potential for the HER is calculated from the LSV curve, recorded at 50 mV s⁻¹ scan rate in 15 mM TFA. The onset potentials (E^onset) for the HER are observed at −0.59 V and −0.65 V for complexes 1 and 2, respectively, as shown in Fig. 3(c). The performance of the catalysts was evaluated by comparing their i_c/i_p values, where i_c is the catalytic peak current and i_p is the peak current in the absence of acid. In the presence of 25 mM TFA, the i_c/i_p value of 18 for 1 is significantly larger than that of 11 observed of 2, as shown in Fig. 5(a). Moreover, i_c/i_p = 10 and <30 were obtained under identical conditions for the corrole-chelated nickel complexes.^41,44 The catalytic peak currents exhibit acid-independence at 250 and 150 mM TFA for complexes 1 and 2, respectively. Above these acid concentrations, no substantial enhancement in the catalytic peak current was observed (Fig. S11, ESI†). Such behavior is observed because of a sufficiently high concentration of acid (H⁺); hence, the catalytic process is not much faster to further deplete the acid during the course of the experiment. In this limit, the i_c/i_p value becomes independent of both the catalyst concentration and the scan rate. Therefore, using this acid-independent i_c/i_p value, the rate constant (k) for the hydrogen evolution process can be calculated using the formula⁶⁴

where n corresponds to the number of electrons transferred (two for the HER), R is the universal gas constant, T is the experimental temperature, ν is the potential scan rate, and F is Faraday's constant. The values of i_c/i_p = 92 and 52 for complexes 1 and 2, shown in Fig. 5a, give the turnover frequencies (TOFs) using eqn (1) of 3.28 × 10³ and 1.05 × 10³ s⁻¹, respectively.

Fig. 5 (a) Dependence of the relative catalytic peak current (i_c/i_p) on the increasing concentration of TFA using complexes 1 and 2 at 200 mV s⁻¹, (b) Tafel plot representing the relationship between the overpotential and the logarithmic current density of 1–3. Plots (c) and (d) represent the charge and current build-up vs time in the CPE experiment recorded at −1.6 V vs. Fc^+/0 for 1 mM of the complex in acetonitrile containing 15 mM TFA and 0.1 M TBABF₄.

The overpotential (η) and Tafel slope for the HER are other important characteristic parameters to evaluate the performance of catalysts. η is estimated by applying the method given by Fourmond et al.⁶⁵ using the theoretical half-wave potential (E^T_1/2) and the experimental catalytic half-wave potential (E_cat/2) as:

η = |E^T_1/2 − E_cat/2| (2)

The E^T_1/2 is taken as −0.80 V versus Fc^+/0 for TFA in CH₃CN, calculated using the Fourmond method (eqn (S2), ESI†).⁶⁵ The overpotential for 1 was obtained to be 450 mV, which is comparable with that of corrole-chelated Nickel complexes⁴¹ and much lower than that of 2 (640 mV). The lower value of the Tafel slope as 112 mV per decade for complex 1 compared to 152 mV per decade for complex 2 shown in Fig. 5(b) indicates the fast and efficient electrocatalytic hydrogen production ability of complex 1. The amount of charge and current produced during the constant potential electrolysis (CPE) experiment for a 30-minute duration follows the order 1 > 2 > 3, as shown in Fig. 5(c and d). This obtained charge is directly proportional to the amount of H₂ generated per mole of the catalyst and related to the turnover number (TON) for the catalysts. The physical characteristic parameters for 1 and 2 are listed in Table 3.

Table 3 Electrochemical characteristic properties of complexes 1–3 (1 mM) in 15 mM TFA

Complex	E ^onset/V	E cat/2/V	η/mV		Tafel slope (mV dec^−1)	TOF/s^−1 ×10^3
1	−0.59	−1.25	450	92	112	3.28
2	−0.65	−1.44	640	52	152	1.05
3	−0.81	−1.92	1012	—	467	—

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The foot-of-the-wave analysis (FOWA)^64,66,67 was further applied for the mechanistic elucidation via the identification of the rate-determining step and the estimation of the useful HER kinetics constant from CV data. The use of low concentrations of the acid in FOWA is usually preferred to directly analyze the electrochemical reduction process.⁶⁶ The low acid concentration generally increases the lifetime of the catalytic intermediates (Ni^III/II–H) and helps to detect them using electrochemical experiments. For the widely accepted ECEC mechanism for Ni-based complexes,^{6,41–43,68,69} the generation of the foot of wave axis (1/{1 + exp[(F/RT)(E − E_cat/2)]}) vs. i/i_p plot (Fig. S12, ESI†) exhibits a linear behavior for both 1 and 2 in 10 and 15 mM TFA. For mechanistic insight, the estimation of the rate-determining step is an important perspective. Even for a fast ET process, the kinetic limitations usually arise because of the slow first (k₂ ≫ k₁) or the second (k₁ ≫ k₂) protonation step. For 1 and 2, the E_cat/2 values are more negative than their E° values (without TFA), suggesting that the first protonation step (k₁), resulting in the formation of Ni^III–H, is the rate-determining step. For 1 and 2, the ratio of k₁/k₂ (eqn (S6), ESI†) is 0.026 and 0.00028, respectively.

In addition to the facile kinetics and low overpotential requirements, the stability factor is another important parameter for an electrocatalyst performance evaluation. Both catalysts (1 and 2) exhibit good stability in the presence of TFA after conducting the CPE experiment for 30 minutes, as shown in Fig. 5(d). UV-vis spectroscopy also reveals similar peak characteristics and intensity after the 4 h of addition of 25 equivalents of acid, as shown in Fig. S13 (ESI†). In addition, the standard rinse experiment was performed which exhibits no significant current difference and additional redox waves in CVs, as shown in Fig. S14 (ESI†). This confirms that the HER process occurs through a molecular electrocatalyst present in the solution rather than an ill-defined nano-particulate material deposited at the working electrode (GC) surface during the electrochemical measurements.

Electronic and geometric/steric effect-induced enhanced proton coupled electron transfer. The large value of E_onset − E_cat/2 for the HER by 1 (660 mV) and 2 (790 mV) in 15 mM TFA indicates the dominant contribution of proton transfer (PT) during the electron transfer (ET) process. The absence of any change in the position of peaks in the UV-vis spectrum of 1 and 2 in the presence of 25 equivalents of TFA opposed the possibility of complete PT prior to the ET. On the other hand, the only ET prior to the PT is also inconsistent as the onset of the catalysis process in Fig. 3(b and c) occurs at a relatively lower potential compared to the reduction of 1 and 2 in the absence of acid, as shown in Fig. 3(a). A more feasible phenomenon is the proton-coupled electron transfer that leads to the formation of a Ni^III–H intermediate,⁴¹via reduction of Ni^II(L) followed by PT. The CVs in Fig. 3a and b show a small difference in redox potentials, ΔE° = |E°(3) − E°(1)| = 130 mV from Table 2, and their electrocatalytic onset potentials, ΔE^onset = |E^onset(3) − E^onset(1)| = 220 mV from Table 3. This suggests the weak electronic effect of the coordinated ligands. On the other hand, the large difference in the catalytic half-wave potential between complexes 1 and 3, ΔE_cat/2 = |E_cat/2(3) − E_cat/2(1)| = 670 mV from Table 3, suggests an enhanced rigidity effect of the coordinated ligands.

The enhanced electrocatalytic proton reduction by 1 compared to 2 is not only because of the electronic effects of coordinated ligands but also pertains to the spatial flexibility of the PPh₃ ligand in 1 compared to 2. The addition of a proton to the reduced form (electrocatalytically active) of complexes 1 and 2 forms a five-coordinated reaction intermediate before hydrogen production. The chelation effect in 2 by both DPPF and dithiolate ligands results in the only possibility of a square pyramidal intermediate and restricts the further coordination of other H⁺ followed by electron transfer. However, 1 has several possibilities to obtain a stable configuration between trigonal bipyramidal and square pyramidal geometries at the transition state.⁶ These possible geometric fluctuations due to free PPh₃ ligands cause enhanced entropic contribution, which lowers the activation free energy. Therefore, 1 is more likely to follow the lower activation energy route and consequently become a fast and robust catalyst for proton reduction compared to complexes 2 and 3.

Furthermore, we have listed and compared the recently synthesized Ni(II) chelate complexes in Table S2 (ESI†) to decipher the electronic and geometric effects in HER catalysis. Mirica and coworkers described a bioinspired electrocatalyst, (N₂S₂)Ni(II), which produces hydrogen from TFA with a TOF of ∼1250 cm⁻¹ at an acid concentration of <43 mM in CH₃CN.⁷⁰ This work highlights the role of a pendant hemilabile pyridyl group within the N₂S₂ ligand in high HER kinetics. Furthermore, an interesting work by Lau et al.³⁹ showed that the incorporation of a phosphorous atom in place of the nitrogen donor site in the tetradentate chelating macrocyclic ligand enhances the electrocatalytic HER activity of the nickel(II) molecular complexes. They have reported a turnover frequency (TOF) of 220 s⁻¹ at 1007 mV overpotential using the Ni^II(N₃P ligand) electrocatalyst in the presence of HClO₄ acid in CH₃CN. Recently, Zhang et al.⁷¹ reported Ni[P₂S₂C₂(C₆H₄R-P)₂] complexes which exhibit variation in η by 110 mV due to change in the substituent (R) from an electron donating group (–OCH₃) to an electron withdrawing (–Br) group. Generally, we observe a weak electronic effect of substituents on chelate Ni(II)S₂P₂/S₂N₂ complexes. Mitsopoulou et al.⁷² also reported heteroleptic oxothiolate Ni(II) complexes to investigate the role of electronic properties of substituents in the Ni(II)/Ni(I) redox potential and a high overpotential was observed in the electrocatalytic HER with an electron donating methyl substituent in the 2,2′-bipyridine ligand. Also, Qiu et al.⁴⁴ reported bent bis(dipyrrin) Ni(II) complexes which exhibit a lower overpotential compared to planar bis(dipyrrin)Ni(II). In our work, complex 3 shows an η value comparable to that of Ni[P₂S₂C₂(C₆H₄OCH₃P)₂], while complexes 1 and 2 (with a distorted square planar geometry) show comparatively lower η values, as listed in Table S2.† The variation in redox potentials (electronic effect) of 1–3 (Fig. 3a) originates from the poor electron donating ability of PPh₃⁷³ compared to those of DPPF and DPPE ligands, whereas the significant effect on the electrocatalytic behavior arises due to the rigidity and chelation effect of PPh₃, DPPF and DPPE ligands. The difference in η value in complexes 1 and 3 is 562 mV which is because of the flexibility of the PPh₃ ligand in 1 compared to the chelating ligand (DPPE) in 3. Therefore, this suggests that distorted or bent square planar Ni(II)-based complexes with flexible ligands exhibit enhanced electrocatalytic activity, resulting in a low overpotential and high turnover frequency. The detailed schematic mechanistic pathway is shown in Fig. 6 and discussed in the following section.

Fig. 6 Proposed mechanistic scheme (based on experimental data) for H₂ production by complexes 1 and 2.

Proposed mechanism

The experimentally observed most probable mechanistic pathway for complexes 1 and 2 follows the ECEC mechanism, widely accepted for several Ni-based complexes,^{6,41–43,68,69} where E indicates the electron transfer process, while C represents the chemical change due to H⁺ binding. The initial protonation of the ligand causes inhibition of electron donation from the ligand to the central Ni^II. This makes the reduction easier and results in a shift of the reduction peak to a relatively lower potential in the presence of acid. Ni^I(L) is the active catalyst that binds a H⁺ to Ni^I, resulting in the formation of Ni^IIIH(L). Mostly, the two generic pathways for the release of hydrogen gas by metal hydrides are (i) the acid–base (heterolytic) and (ii) the homolytic bimetallic route.⁴⁵ The preferred pathway depends on the steric effect of coordinated ligands along with the strength or the local concentration of the acid and catalyst. The steric effect by the bulky PPh₃ group in 1 and 2 excludes the possibility of homolytic hydrogen production.⁴⁵ Also, the linear variation of the catalytic peak current with the concentration of the acid, as shown in Fig. 5(a), and the complex (Fig. S15 and S16†) additionally supports the heterolytic HER dominance for 1 and 2 under this particular reaction condition.^74,75 Furthermore, this proton coupling promotes the electron transfer process and causes the emergence of a new catalytic peak with the formation of Ni^IIH(L) followed by the coupling of another H⁺, resulting in the generation of H₂ and regeneration of the catalyst.

Conclusion

The present work reported the synthesis and characterization of three heteroleptic Ni(II) dithiolate complexes. The single-crystal X-ray structure revealed that heterobimetallic complex 2 has a distorted square planar geometry about the Ni(II) center, which is coordinated with bidentate S^S-dithiolate and P^P-chelating DPPF ligands. All complexes are investigated as an electrocatalyst for proton reduction to hydrogen and their catalytic activity follows the order 1 > 2 > 3. The best catalytic activity has been observed with complex 1, which shows a low overpotential (450 mV) and high TOF (3.28 × 10³ s⁻¹). The FOWA of CV data indicates that the H⁺ reduction by the Ni complex follows the ECEC mechanism. The electrochemical studies show that the rigidity and chelation of the phosphine ligand play a major role in the electrocatalytic properties of complexes. Furthermore, the rinse experiment studies confirm that the molecular integrity of complexes remains preserved during the electrocatalytic proton reduction.

Conflicts of interest

There are no conflicts to declare.

Acknowledgements

We gratefully acknowledge financial support from the University Grants Commission (UGC), New Delhi, for the award of the UGC-BSR Research Start-Up-Grant (No. F. 30-431/2018) (BSR) and Banaras Hindu University-Seed Grant Under the IoE Scheme. Anjali acknowledges the CSIR, India, for the JRF. The Department of Chemistry, Institute of Science, Banaras Hindu University, Varanasi, is gratefully acknowledged for its infrastructural facilities. R. K. is thankful to SERB (CRG/2021/005388) for financial support. G. K. M. thanks UGC-India for providing the SRF.

References

1. X. L. Gu, J. R. Li, Q. L. Li, Y. Guo, X. B. Jing, Z. B. Chen and P. H. Zhao, J. Inorg. Biochem., 2021, 219, 111449.
2. L. Huo, C. Jin, K. Jiang, Q. Bao, Z. Hu and J. Chu, Adv. Energy Sustainability Res., 2022, 3, 2100189.
3. W.-X. Liu, C.-L. Wang, J.-M. Lei, S.-Z. Zhan and S.-P. Wu, Inorg. Nano-Met. Chem., 2022, 52, 533–541.
4. G. G. Luo, Y. H. Wang, J. H. Wang, J. H. Wu and R. B. Wu, Chem. Commun., 2017, 53, 7007–7010.
5. A. Barma, M. Chakraborty, S. K. Bhattacharya, P. Ghosh and P. Roy, Mater. Adv., 2022, 3, 7655–7666.
6. L. Gan, T. L. Groy, P. Tarakeshwar, S. K. Mazinani, J. Shearer, V. Mujica and A. K. Jones, J. Am. Chem. Soc., 2015, 137, 1109–1115.
7. V. Vij, S. Sultan, A. M. Harzandi, A. Meena, J. N. Tiwari, W.-G. Lee, T. Yoon and K. S. Kim, ACS Catal., 2017, 7, 7196–7225.
8. D. Jana, H. K. Kolli, S. Sabnam and S. K. Das, Chem. Commun., 2021, 57, 9910–9913.
9. T. Fogeron, J. P. Porcher, M. Gomez-Mingot, T. K. Todorova, L. M. Chamoreau, C. Mellot-Draznieks, Y. Li and M. Fontecave, Dalton Trans., 2016, 45, 14754–14763.
10. J. K. Yadav, B. Singh, S. K. Pal, N. Singh, P. Lama, A. Indra and K. Kumar, Dalton Trans., 2023, 52, 936–946.
11. J. K. Yadav, A. Mishra, G. K. Mishra, S. K. Pal, K. U. Narvekar, A. Rahaman, N. Singh, P. Lama and K. Kumar, New J. Chem., 2023, 47, 20583–20593.
12. A. Z. Haddad, B. D. Garabato, P. M. Kozlowski, R. M. Buchanan and C. A. Grapperhaus, J. Am. Chem. Soc., 2016, 138, 7844–7847.
13. A. Z. Haddad, S. P. Cronin, M. S. Mashuta, R. M. Buchanan and C. A. Grapperhaus, Inorg. Chem., 2017, 56, 11254–11265.
14. M. D. Sampson and C. P. Kubiak, Inorg. Chem., 2015, 54, 6674–6676.
15. H. Tang and M. B. Hall, J. Am. Chem. Soc., 2017, 139, 18065–18070.
16. T. Straistari, R. Hardré, J. Massin, M. Attolini, B. Faure, M. Giorgi, M. Réglier and M. Orio, Eur. J. Inorg. Chem., 2018, 2018, 2259–2266.
17. K. E. Dalle, J. Warnan, J. J. Leung, B. Reuillard, I. S. Karmel and E. Reisner, Chem. Rev., 2019, 119, 2752–2875.
18. A. Dutta, A. K. Chatterjee, K. Bhar, A. Banerjee and R. Ghosh, J. Chem. Sci., 2019, 131, 1–7.
19. W. R. McNamara, Z. Han, P. J. Alperin, W. W. Brennessel, P. L. Holland and R. Eisenberg, J. Am. Chem. Soc., 2011, 133, 15368–15371.
20. D. Espa, L. Pilia, L. Marchiò, M. L. Mercuri, A. Serpe, A. Barsella, A. Fort, S. J. Dalgleish, N. Robertson and P. Deplano, Inorg. Chem., 2011, 50, 2058–2060.
21. V. Madhu and S. Das, Inorg. Chem., 2008, 47, 5055–5070.
22. J. A. Zuleta, M. S. Burberry and R. Eisenberg, Coord. Chem. Rev., 1990, 97, 47–64.
23. J. A. Zuleta, J. M. Bevilacqua and R. Eisenberg, Coord. Chem. Rev., 1991, 111, 237–248.
24. S. K. Pal, B. Singh, J. K. Yadav, C. L. Yadav, M. G. B. Drew, N. Singh, A. Indra and K. Kumar, Dalton Trans., 2022, 51, 13003–13014.
25. J.-W. Wang, W.-J. Liu, D.-C. Zhong and T.-B. Lu, Coord. Chem. Rev., 2019, 378, 237–261.
26. A. Das, Z. Han, W. W. Brennessel, P. L. Holland and R. Eisenberg, ACS Catal., 2015, 5, 1397–1406.
27. M. L. Helm, M. P. Stewart, R. M. Bullock, M. R. DuBois and D. L. DuBois, Science, 2011, 333, 863–866.
28. L. L. Efros, H. H. Thorp, G. W. Brudvig and R. H. Crabtree, Inorg. Chem., 1992, 31, 1722–1724.
29. D. K. Bediako, B. H. Solis, D. K. Dogutan, M. M. Roubelakis, A. G. Maher, C. H. Lee, M. B. Chambers, S. Hammes-Schiffer and D. G. Nocera, Proc. Natl. Acad. Sci. U. S. A., 2014, 111, 15001–15006.
30. J. P. Collin, A. Jouaiti and J. P. Sauvage, Inorg. Chem., 1988, 27, 1986–1990.
31. P. Zhang, M. Wang, Y. Yang, D. Zheng, K. Han and L. Sun, Chem. Commun., 2014, 50, 14153–14156.
32. B. H. Solis, A. G. Maher, D. K. Dogutan, D. G. Nocera and S. Hammes-Schiffer, Proc. Natl. Acad. Sci. U. S. A., 2016, 113, 485–492.
33. P. H. A. Kankanamalage, S. Mazumder, V. Tiwari, K. K. Kpogo, H. B. Schlegel and A. C. N. Verani, Chem. Commun., 2016, 52, 13357.
34. D. Xue, Q.-X. Peng and S.-Z. Zhan, Inorg. Chem. Commun., 2017, 82, 11–15.
35. A. Begum, G. Moula and S. Sarkar, Chemistry, 2010, 16, 12324–12327.
36. K. Koshiba, K. Yamauchi and K. Sakai, Angew. Chem., Int. Ed., 2017, 56, 4247–4251.
37. W. A. Hoffert, J. A. S. Roberts, R. M. Bullock and M. L. Helm, Chem. Commun., 2013, 49, 7767–7769.
38. M. L. Reback, B. Ginovska, G. W. Buchko, A. Dutta, N. Priyadarshani, B. L. Kier, M. L. Helm, S. Raugei and W. J. Shaw, J. Coord. Chem., 2016, 69, 1730–1747.
39. L. Chen, G. Chen, C.-F. Leung, S.-M. Yiu, C.-C. Ko, E. Anxolabéhère-Mallart, M. Robert and T.-C. Lau, ACS Catal., 2015, 5, 356–364.
40. R. Tatematsu, T. Inomata, T. Ozawa and H. Masuda, Angew. Chem., Int. Ed., 2016, 55, 5247–5250.
41. Q.-C. Chen, S. Fite, N. Fridman, B. Tumanskii, A. Mahammed and Z. Gross, ACS Catal., 2022, 12, 4310–4317.
42. R. Jain, A. A. Mamun, R. M. Buchanan, P. M. Kozlowski and C. A. Grapperhaus, Inorg. Chem., 2018, 57, 13486–13493.
43. M. Papadakis, A. Barrozo, T. Straistari, N. Queyriaux, A. Putri, J. Fize, M. Giorgi, M. Reglier, J. Massin, R. Hardre and M. Orio, Dalton Trans., 2020, 49, 5064–5073.
44. S. Xue, X. Lv, N. Liu, Q. Zhang, H. Lei, R. Cao and F. Qiu, Inorg. Chem., 2023, 62, 1679–1685.
45. X. Guo, N. Wang, X. Li, Z. Zhang, J. Zhao, W. Ren, S. Ding, G. Xu, J. Li, U.-P. Apfel, W. Zhang and R. Cao, Angew. Chem., Int. Ed., 2020, 59, 8941–8946.
46. E. Cerrada, A. Moreno and M. Laguna, Dalton Trans., 2009, 6825–6835.
47. Anamika, D. K. Yadav, K. K. Manar, C. L. Yadav, K. Kumar, V. Ganesan, M. G. B. Drew and N. Singh, Dalton Trans., 2020, 49, 3592–3605.
48. E. A. Standley, S. J. Smith, P. Muller and T. F. Jamison, Organometallics, 2014, 33, 2012–2018.
49. Ş. Güveli, J. Coord. Chem., 2020, 73, 137–153.
50. A. Yamamoto, T. Yamamoto, S. Komiya and F. Ozawa, Pure & Appl. Chem., 1984, 56, 1621–1634.
51. K. K. Manar, A. N. Gupta, A. K. Gupta, L. B. Prasad, P. Srivastava, M. G. B. Drew and N. Singh, ChemistrySelect, 2017, 2, 2655–2664.
52. A. N. Gupta, V. Kumar, V. Singh, K. K. Manar, A. K. Singh, M. G. B. Drew and N. Singh, Polyhedron, 2015, 101, 251–256.
53. M. Kumar, G. K. Mishra and R. Kant, Electrochim. Acta, 2019, 327, 135024.
54. K. K. Manar, Neetu, Anamika, P. Srivastava, M. G. B. Drew and N. Singh, ChemistrySelect, 2017, 2, 8301–8311.
55. A. B. P. Lever, Inorganic electronic spectroscopy, Elsevier, Amsterdam, 1984.
56. A. T. Armstrong, F. Smith, E. Elder and S. P. McGlynn, J. Chem. Phys., 1967, 46, 4321–4328.
57. Y. S. Sohn, D. N. Hendrickson and H. B. Gray, J. Am. Chem. Soc., 1971, 93, 3603.
58. K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, Wiley Interscience, New York, 4th edn, 1986.
59. K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, Part B, Wiley Interscience, New York, 6th edn, 2009.
60. L. Yang, D. R. Powell and R. P. Houser, Dalton Trans., 2007, 955–964.
61. J. D. Wrixon, J. J. Hayward, O. Raza and J. M. Rawson, Dalton Trans., 2014, 43, 2134–2139.
62. M. R. L. Oliveira, H. P. Vieira, G. J. Perpetuo, J. Janezak and V. M. D. Bellis, Polyhedron, 2002, 21, 2243–2250.
63. W. S. Han, Y.-J. Kim and S. W. Lee, Bull. Korean Chem. Soc., 2003, 24, 60–64.
64. V. C. C. Wang and B. A. Johnson, ACS Catal., 2019, 9, 7109–7123.
65. V. Fourmond, P. A. Jacques, M. Fontecave and V. Artero, Inorg. Chem., 2010, 49, 10338–10347.
66. E. S. Wiedner and R. M. Bullock, J. Am. Chem. Soc., 2016, 138, 8309–8318.
67. E. S. Rountree, B. D. McCarthy, T. T. Eisenhart and J. L. Dempsey, Inorg. Chem., 2014, 53, 9983–10002.
68. E. Ahmad, S. Rai and S. K. Padhi, Int. J. Hydrogen Energy, 2019, 44, 16467–16477.
69. Z. Chen, T. Wang, T. Sun, Z. Chen, T. Sheng, Y.-H. Hong, Z.-A. Nan, J. Zhu, Z.-Y. Zhou, H. Xia and S.-G. Sun, Chin. J. Chem., 2018, 36, 1161–1164.
70. S. Sinha, G. N. Tran, H. Na and L. M. Mirica, Chem. Commun., 2022, 58, 1143–1146.
71. L. Chen, T. Li, B. Xie, C. Lai, R.-W. Ji, J.-Y. He, J.-X. Cao, M.-N. Liu, W. Li and D.-L. Zhang, Catal. Sci. Technol., 2023, 13, 3655–3666.
72. F. Kamatsos, K. Bethanis and C. A. Mitsopoulou, Catalysts, 2021, 11, 401–418.
73. T. Thananatthanachon and M. R. Lecklider, J. Chem. Educ., 2017, 94, 786–789.
74. C. N. Valdez, J. L. Dempsey, B. S. Brunschwig, J. R. Winkler and H. B. Gray, Proc. Natl. Acad. Sci. U. S. A., 2012, 109, 15589–15593.
75. K. Kwak, W. Choi, Q. Tang, M. Kim, Y. Lee, D. E. Jiang and D. Lee, Nat. Commun., 2017, 8, 14723.

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Footnote

† Electronic supplementary information (ESI) available: FT-IR, ¹H, and ³¹P{¹H} NMR spectra of complexes 1–3, PXRD of the complex, homogeneous electrochemical analysis, and FOWA. CCDC 2259175. For ESI and crystallographic data in CIF or other electronic format see DOI: https://doi.org/10.1039/d3dt03932c

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