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. 2024 Feb 14;4(3):306–318. doi: 10.1021/acsorginorgau.3c00061

Manipulating the Rate and Overpotential for Electrochemical Water Oxidation: Mechanistic Insights for Cobalt Catalysts Bearing Noninnocent Bis(benzimidazole)pyrazolide Ligands

Yu-Ting Wu †, Sharad V Kumbhar †, Ruei-Feng Tsai †, Yung-Ching Yang ‡, Wan-Qin Zeng ‡, Yu-Han Wang †, Wan-Chi Hsu †, Yun-Wei Chiang †,*, Tzuhsiung Yang †,*, I-Chung Lu ‡,*, Yu-Heng Wang †,*
PMCID: PMC11157513  PMID: 38855334

Abstract

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Electrochemical water oxidation is known as the anodic reaction of water splitting. Efficient design and earth-abundant electrocatalysts are crucial to this process. Herein, we report a family of catalysts (1–3) bearing bis(benzimidazole)pyrazolide ligands (H2L1–H2L3). H2L3 contains electron-donating substituents and noninnocent components, resulting in catalyst 3 exhibiting unique performance. Kinetic studies show first-order kinetic dependence on [3] and [H2O] under neutral and alkaline conditions. In contrast to previously reported catalyst 1, catalyst 3 exhibits an insignificant kinetic isotope effect of 1.25 and zero-order dependence on [NaOH]. Based on various spectroscopic methods and computational findings, the L3Co2III(μ-OH) species is proposed to be the catalyst resting state and the nucleophilic attack of water on this species is identified as the turnover-limiting step of the catalytic reaction. Computational studies provided insights into how the interplay between the electronic effect and ligand noninnocence results in catalyst 3 acting via a different reaction mechanism. The variation in the turnover-limiting step and catalytic potentials of species 1–3 leads to their catalytic rates being independent of the overpotential, as evidenced by Eyring analysis. Overall, we demonstrate how ligand design may be utilized to retain good water oxidation activity at low overpotentials.

Keywords: water oxidation, molecular electrocatalyst, homogeneous catalysis, noninnocent ligand, rate–overpotential correlation

Introduction

Energy security and the mitigation of climate change are both known to be at the core of energy policies worldwide.1−3 Due to the ongoing population growth and industrialization of the developing world, the use of renewable energy sources to reduce our dependence on fossil fuels is critical to address these energy-related issues.4,5 In addition to renewable energy sources, such as solar power, hydropower, and biomass, H2 has gradually gained global attention as an alternative to conventional fossil fuels.6,7 Due to its cleanliness and high energy density, H2 is considered a promising energy carrier to store energy in the form of chemical H–H bonds.8−12 Although H2 can be generated via water splitting, efficient hydrogen evolution can only be achieved using efficient, inexpensive, and robust (electro)catalysts for the water oxidation reaction (WOR, eq 1).13−18 However, the involvement of multielectron/multiproton reactions along with both bond-forming and bond-breaking processes results in sluggish reaction kinetics and limits the overall performance of the WOR.19−22

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With the above considerations in mind, exploring low-cost and earth-abundant catalysts to promote practical application of the WOR in energy conversion is imperative. Therefore, precious metal catalysts composed of noble elements,23 such as Ru and Ir, must be replaced. To date, various first-row transition metal catalysts,24−27 such as Co-based complexes, have been considered suitable candidates for use as molecular water oxidation catalysts (MWOCs) because of the similarity between CoIII and RuII (i.e., a low-spin electronic structure and an octahedral geometry).28−32 Although Ru-based WOCs have been extensively studied, systematic investigations have yet to be carried out on Co-based WOCs.

Homogeneous molecular (electro)catalysts are of particular interest due to the ease of probing their reaction mechanisms by spectroscopic methods, and the fact that their reaction sites can be adjusted using alternating ligand scaffolds.15,33 With a well-defined active site that can be characterized spectroscopically, the rational design of molecular catalysts toward efficient (electro)catalysis becomes more feasible. For example, in the case of metal-based catalysts, the reaction rate and overpotential (η) could be improved by modifying the ligand primary and secondary coordination spheres to optimize the electronic and steric effects, proton relay, hemilability, and noninnocent characteristics.34−41 One vital feature of the noninnocent ligand motif is that the ligand can participate in both proton- and electron-transfer (PT and ET) steps (i.e., (de)protonation and redox events), which could lead to improved catalytic efficiencies for energy-related multielectron, multiproton reactions by facilitating the PT and ET rates.42−46

Recently, our group investigated dimeric cobalt complexes, 1 and 1-Me, bearing bis(benzimidazole)pyrazolide (bbp) ligands, as homogeneous catalysts for the electrocatalytic oxidation of water (Figure 1).30 Preliminary mechanistic studies of the WOR catalyzed by 1 indicated that the rate exhibited a first-order dependence on [1] and [H2O], and a kinetic isotope effect (KIE) of 2.67 was observed when the reaction was conducted under alkaline conditions. Notably, in the case of catalyst 1, which bears a noninnocent ligand that exhibited higher rates at lower overpotentials, a reverse trend in the rate–overpotential correlation was observed relative to 1-Me and the majority of previously reported dimeric Ru-based WOCs.30 This encouraging result suggested that both the kinetic and thermodynamic performances of an MWOC can be improved by introducing noninnocent ligands, thereby paving a route to developing efficient nonprecious metal catalysts for application in the WOR.

Figure 1.

Figure 1

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Structures of 1 and 1-Me dimeric Co-based WOCs previously synthesized by our group. Structures of catalysts 2 and 3 are the new dimeric Co-based WOCs described herein. The structures of catalysts 4–7 are also shown; these are previously reported dimeric Co-based WOCs.47−49

Therefore, in the current study, we focus on 3 having a similar framework to 1 but incorporating electron-donating substituents on the bbp ligand to examine the synergy between the electronic effect and the noninnocent behavior in terms of its effect on the catalyst performance (Figure 1). In addition, a comprehensive mechanistic study is carried out based on experimental and computational data, and the turnover-limiting steps (TLSs) and catalysis-initiating potentials (Ecat) of the prepared catalysts are determined. The results of this study are expected to provide valuable insights into developing competent MWOCs.

Results and Discussion

Preparation and Characterization of the Dimeric Cobalt Complexes

The cooperativity of (redox) noninnocent ligands on molecular catalysts for application in small molecule catalysis (e.g., water oxidation and carbon dioxide reduction) has manifested in kinetics and thermodynamics.34−36 Thus, to further explore the synergy between the electronic effect and the (redox) noninnocent nature of MWOCs, a series of dimeric cobalt complexes bearing bbp ligands was prepared, wherein these ligands contained electron-donating substituents on their benzimidazolyl rings (Figure 2, –Me (2) and –OMe (3)). Complexes 2–3 were synthesized via a metalation reaction between Co(tpy)Cl2 (tpy: terpyridine) and the corresponding bbp ligands (H2L2 and H2L3) under aerobic conditions, as described in Section S2b,c of the Supporting Information. The X-ray crystal structures of 2 and 3 revealed that the cobalt ions occupy a distorted octahedral geometry, with the μ-peroxo-bridged Co2III unit and each CoIII center being supported by a tridentate tpy ligand and two N-donor sites through the backbone bpp scaffold. The sharp 1H and 13C NMR spectra recorded for 2 and 3 were indicative of their low-spin 3d6 electronic configurations (Section S17 of the Supporting Information).30,47,50

Figure 2.

Figure 2

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Oak ridge thermal ellipsoid plot (ORTEP) diagrams of complexes 2 and 3 with 30% probability ellipsoids. The hydrogen atoms and counterions (PF6–) have been omitted for clarity.

Electrochemical Behavior of the Dimeric Cobalt Complexes

Complexes 2 and 3 were subjected to cyclic voltammetry (CV) in a mixture of acetonitrile (CH3CN) and 0.1 M tetrabutylammonium hexafluorophosphate (TBAP, a supporting electrolyte) under anaerobic conditions. Note that all potentials reported in this work are relative to the ferrocenium/ferrocene couple ( Fc+/0). Thus, as shown in Figure 3, the CV of 2 contains a quasi-reversible oxidation wave at 1.01 V, representing a slight cathodic potential shift of ∼80 mV compared to that of 1. Surprisingly, a significant difference in the redox behavior of 3, featuring two oxidation processes at 0.82 and 0.91 V, was observed by both CV and differential pulse voltammetry (DPV), indicating that the more electron-donating substituents (σp: –OCH3 = −0.27; –CH3 = −0.17) influence the redox reservoir properties of these Co-based complexes. In addition, the striking shifts in the redox potentials of 2 and 3 under alkaline conditions were attributed to the increased electron densities resulting from deprotonation of the acidic –NH protons on their bpp ligands (Figure 3b, Table 2, and Section S4b of the Supporting Information). Computational studies further validated the redox couple assignments of 2 and 3, as shown by the quasi-reversible/irreversible CV behaviors (Table 2, vide infra). The chemical reversibility of 2 and 3 under alkaline conditions was further investigated using ultraviolet–visible (UV–vis) spectroelectrochemical (SEC) experiments (Figure 3c,d), wherein the potential windows were selected based on the redox couples obtained from their cyclic voltammograms (Figure 3a,b). The appearance of a new ligand-to-metal charge transfer band centered at the UV region (270–290 nm) indicates that the oxidized forms of 2 and 3 were afforded. Although the recorded cyclic voltammograms show quasi-reversible/irreversible electrochemical processes of 2 and 3, the resulting UV–vis SEC spectra demonstrated the good chemical reversibility of 2 and 3 under electrochemical conditions.

Figure 3.

Figure 3

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(a) Cyclic voltammograms of 1 (black),302 (blue), and 3 (red) in CH3CN. (b) Cyclic voltammograms of 1, (black),302 (blue), and 3 (red) in CH3CN in the presence of 10 mM NaOH, 10 mM NaOH, and 10 mM NaOH containing 1 mM 15-crown-5 crown ether, respectively. 15-Crown-5 crown ether was added to increase the solubility of NaOH and the nucleophilicity of the hydroxide ion in anhydrous CH3CN.51 The redox potentials of 3 were further confirmed by DPV ([cat] 0.4 mM; scan rate: 100 mV/s; working electrode (WE): glassy carbon (GC) disk electrode). UV–vis SEC spectra were acquired during the potential sweep of (c) 2 and (d) 3 in MeCN under alkaline conditions ([2 or 3]: 0.04 mM, [NBu4OH]: 1 mM, WE: Pt gauze). Scan rate: 1 mV/s; initial and final potential (mV): 450 (2) and 100 (3); switching potential (mV): 850 (2) and 500 (3). CVs were plotted by using the IUPAC convention. See the Supporting Information for the general considerations for the cyclic voltammetric measurements.

Table 2. Calculated Redox Potentials of 2+ and 3+ under Alkaline Conditions (V vs. Fc/Fc+).

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  2
 3
redox couplesa 2+/22+ 22+/23+ 3+/32+ 32+/33+
expt 0.59 0.73 0.22 0.37
calcd 0.68 0.80 0.27 0.29
assignment L2 L2 L3 L3
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a

3+ = 3-I, 32+ = 3-II, 33+ = 3-III.

Kinetic Studies and Stability Determination

Under the catalytic conditions employed herein, the CV results show an increase in the anodic wave when complexes 2 and 3 were present. In addition, the value of the catalytic current (ic) exhibited a linear, first-order dependence on [3] and a half-order dependence on [H2O] (Figure 4). These results may be rationalized using eq 2 as follows:52,53

Figure 4.

Figure 4

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(a) Cyclic voltammograms of 3 (0.1–1 mM) in solutions containing 5 M H2O in CH3CN. Inset: plot of ic/ip at 1.5 V vs [3]. (b) Cyclic voltammograms of 3 (0.4 mM) in solutions containing 0–10 M H2O in CH3CN. Inset: plot of ic/ip at 1.5 V vs [H2O]1/2. General conditions: [TBAP]: 0.1 M, WE: GC disk electrode, RE: Ag+/Ag (0.01 M AgNO3/0.1 M TBAP in MeCN), CE: Pt wire, scan rate: 100 mV/s. CVs were plotted using the IUPAC convention.

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where nc is the number of electrons transferred in the catalytic wave, F is the Faraday constant (96,485 C/mol), A is the surface area of the WE (2.25 × 102 π cm2), Ccat is the catalyst concentration (mol/cm3); Dcat is the diffusion coefficient of the catalyst (Dcat, determined by diffusion ordered spectroscopy (DOSY), Section S4d of the Supporting Information), kobs is defined as pseudo-first-order rate constant or the turnover frequency (TOF, s–1), Csubstrate is the substrate concentration (mol/cm3), and kcat is the catalytic rate constant (M–1 s–1). Equation 3 accounts for the first-order dependence of the catalytic current on [3] (Ccat) and the half-order dependence on [H2O] (Csubstrate), which together yield the second-order rate law presented in eq 4.

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The kinetic studies of 3 were performed under alkaline conditions to gain further mechanistic insight into the catalytic process. The WOR rates of 3 were observed to be invariable with respect to the NaOH concentration, suggesting a zero-order dependence on [NaOH] (Figure 5a). Upon CV investigation of the KIE for the WOR (Figure 5b), an insignificant KIE of 1.25 was derived using eq 5 (kcat(H) and kcat(D), rate constants under nondeuterated and deuterated conditions; ic(H) and ic(D), catalytic currents under nondeuterated and deuterated conditions), suggesting that PT is not the TLS.

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Figure 5.

Figure 5

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(a) Cyclic voltammograms of 3 (0.4 mM) in solutions containing 0–40 mM NaOH and 5 M H2O in CH3CN. Inset: plot of ic/ip at 1.5 V vs [NaOH]1/2. (b) Cyclic voltammograms of 3 (0.4 mM) in solutions containing 5 M H2O and 10 mM NaOH (red) and 5 M D2O and 10 mM NaOD (black). General conditions: TBAP: 0.1 M, WE: GC disk electrode, RE: Ag+/Ag (0.01 M AgNO3/0.1 M TBAP in MeCN), CE: Pt wire, scan rate: 100 mV/s. CVs were plotted using the IUPAC convention.

The results of these two kinetic experiments are contrary to those obtained for complex 1 under alkaline conditions (Table 1), indicating the possibility that the WOR follows a different reaction pathway due to modulation of the electronic effect and the noninnocence of the ligand scaffold. Subsequently, the electrochemical O2 evolution during controlled-potential water electrolysis by 3 was measured using a PreSens Microx 4 fiber-optic oxygen meter, and the Faraday efficiency was determined to be ≥90% over 3 h (Section S6b of the Supporting Information). Moreover, a TiIV(O)SO4 colorimetric assay confirmed the negligible formation of H2O2 in the reaction mixture (1% H2O2, nc ≈ 4),54,55 suggesting that O2 is the main product (Section S6c of the Supporting Information). The data presented above indicate that 3 is a competent WOC.

Table 1. Comparison of the Kinetic Parameters for the WOR Catalyzed by 1 and 3 under Neutral and Alkaline Conditionsa.

  1b 3
rate law R ∝ [cat]1[H2O]1 R ∝ [cat]1[H2O]1
base assistance yes no
KIE neutral 1.32(1) 1.26(2)
alkaline 2.67(10) 1.25(6)
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a

General conditions: [cat]: 0.4 mM, [NaOH]: 10 mM, TBAP: 0.1 M, solvent: MeCN. [H2O]: 10 M (18%, v/v) for 1; 5 M (9%, v/v).

b

Obtained from ref (38).

The stability of 3 was then accessed using a range of techniques.56 First, the rinse test was performed according to a previously described procedure,57 and it was found that the current obtained using the rinsed WE was similar to the background current (Section S5a of the Supporting Information). In addition, successive CV scans of 3 were carried out for 3 h under catalytic conditions, giving stable currents with no significant increases or decreases observed (Section S5b of the Supporting Information). In addition, after subjecting complex 3 to CPE for 8 h at 1.5 V, energy-dispersive X-ray spectroscopy (EDXS) did not detect the formation of cobalt-containing nanoparticle deposits on the electrode surface above the detection limits (Section S5c of the Supporting Information). These results provide evidence of the homogeneity of the cobalt complexes employed herein and exclude the involvement of heterogeneous CoOx in the catalytic process.

Identify the Reaction Intermediates and the Catalyst Resting State

Additional experiments were subsequently performed to investigate the Co species involved in the reaction under alkaline conditions. It was found that upon the addition of 2 equiv of sodium hydride to a solution of 3 in anhydrous CH3CN, the doubly deprotonated cobalt species 3+ (3-I) was formed, which was crystallographically characterized (Figure 6a). More specifically, the unit cell of 3-I consists of a single molecule of the cobalt complex and a single molecule of the PF6– anion, which differs from the neutral cobalt complexes 1–3, which contain three molecules of PF6– in their unit cells (Section S18 of the Supporting Information). Note that the Co2IIIO2 core unit of 3-I was retained, indicating that neither CoIII–O nor O–O bond cleavage takes place upon the deprotonation of 3. The 1H NMR spectrum recorded for 3-I confirms its diamagnetic nature, and the upfield shift of the 1H signals may be attributed to the shielding effect arising from the negatively charged bpp ligand (Section S12 of the Supporting Information). Subsequently, the one-electron oxidized cobalt species of 3-I was characterized by in situ electron paramagnetic resonance (EPR) spectroscopy by the addition of 1 equiv of [NBu4][IO4] to a solution of 3 under alkaline conditions. The resulting EPR spectrum recorded at 8 K exhibited a distinctive hyperfine splitting that correlated best with the low spin (S = 1/2) μ-1,2-peroxo dimeric cobalt species, 32+ (3-II).58 In addition, the simulated EPR spectrum was in good agreement with the experimental data of 3-II, yielding the principal g values of gx = 2.168, gy = 2.147, and gz = 1.965, along with the A-tensors of AxN, AyN, AzN = 26, 30, 35, and AxCo, AyCo, and AzCo = 102, 95, and 109 in units of Gauss (Figure 6b and Section S13 of the Supporting Information).

Figure 6.

Figure 6

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(a) ORTEP diagram of 3-I with 50% probability ellipsoids. (b) Experimental (black) and simulated (red) EPR spectra of 3-II in frozen CH3CN at 8 K. The EPR spectrum of 3-II was simulated using EasySpin (version 5.2.35). (c) cold-spray ionization mass spectrometry (CSI-MS) spectrum of 3-III (m/z 344.42) in MeCN at −5 °C. The theoretical isotopic distribution is shown in the blue bar. Inset: time-dependent CSI-MS ion intensities were obtained for 3-I and 3-III.

The cobalt species participating in the redox process under alkaline conditions were then analyzed by cold-spray ionization mass spectrometry (CSI-MS). Note that the CSI-MS platform prepared in-house provides the capability of monitoring the time-dependent progression of the reaction intermediates. Therefore, a reaction mixture containing 0.1 mM 3, 2.5 mM [NBu4][OH], and 1 mM [NBu4][IO4] in anhydrous CH3CN was introduced into the positive mode CSI source (Figure 6c and Section S14 of the Supporting Information). The time-dependent CSI-MS spectra indicated that the ion intensity at a mass–charge ratio (m/z) of 1033.33 (3-I, [3–2H–3PF6]+) decreased as a function of time. In addition, the ion signal at m/z 344.42 was attributed to the formation of a doubly oxidized species of 3-I (33+, 3-III), as illustrated in Figure 6c.

To probe the catalyst resting state of 3 under the catalytic conditions examined herein, the reaction mixture containing 5 mM 3, 20 mM NaOH, 100 mM NBu4IO4, and 5 M H2O was monitored by 1H NMR spectroscopy. The 1H NMR time-course analysis revealed that a new diamagnetic cobalt species with a characteristic signal at δ 4.7 ppm began to form after ∼5 min, as shown in Figure 7a. This cobalt species exhibits distinct chemical shifts from those of 3 and 3-I, and the integration of all protons suggests that this is a dimeric cobalt species that preserves the identical polypyridyl ligands as 3 (Section S12 of the Supporting Information). In addition, Figure 7b shows the time-dependent UV–vis SEC traces of the WOR catalyzed by 3 under alkaline conditions, wherein the disappearance of the broad transition centered at 500 nm indicates cleavage of the O–O bond from the CoIII–O–O–CoIII core,30,50,59,60 which differs from those obtained under noncatalytic conditions (Figure 3c,d). Considered together, the NMR and UV–vis SEC findings indicate that the catalyst resting state is not 3-I, 3-II, or 3-III. Therefore, to further evaluate the catalyst resting state, the catalytic reaction mixture was monitored in situ by CSI-MS for 40 min. The results showed an ion signal at m/z 509.25, which may be assigned to an [L3Co2III(μ-OH)]2+ species, as confirmed by isotopic distribution fitting (see Figure 7c).48,49,61 Importantly, the time-dependent analysis of this system demonstrates that the concentration of this intermediate increases during the steady-state reaction period, which suggests that the catalyst resting state may be attributed to a dimeric cobalt μ-hydroxide species.

Figure 7.

Figure 7

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(a) Time-course 1H NMR spectra of WOR catalyzed by 3 under alkaline conditions in CD3CN and the corresponding 1H NMR spectra of 3 and 3-I. See Figures S78–83 for the one-dimensional and two-dimensional NMR spectra of 3. (b) Time-course UV–vis SEC spectra of the WOR catalyzed by 3 in CH3CN at 1.5 V ([H2O] = 5 M, [NBu4OH] = 1 mM). An Eapp of 1.5 V was selected because the value of ic at 1.5 V is independent of the scan rate (ν) when ν is ≥800 mV/s (Section S9 of the Supporting Information). (c) Time-dependent ion intensity of [L3Co2III(μ-OH)]2+ obtained by monitoring the WOR catalyzed by 0.1 mM 3 in CH3CN at −5 °C with [NBu4IO4] = 5 mM, [NBu4OH] = 2.5 mM, and [H2O] = 10 M. Inset: CSI-MS spectra and isotopic distribution fitting (blue bars) of [L3Co2III(μ-OH)]2+.

Computational Findings and Overall Catalytic Mechanism

Quantum chemical studies using density functional theory (DFT) were subsequently performed to further understand the experimentally observed redox potentials of 3, the reaction intermediates, and the different KIE and [OH–]-dependent results obtained for the WORs catalyzed by 1 and 3. Initially, the redox potentials of 3+/32+ and 32+/33+ were calculated, and the results are summarized in Table 2. Excellent agreements can be observed between the calculated and experimental redox potentials under alkaline conditions (Figure 3b). The first two oxidations at ∼0.2 and 0.4 V were found to be centered on L3, differing from the observations made for complex 1, in which only one accessible oxidation occurred on L1.30 This was attributed to tetrasubstitution of the bbp ligand by electron-donating methoxy groups.

To investigate the WOR catalyzed by 3, computational studies were performed, and the calculated catalytic cycle is summarized in Figure 8. The precatalyst was selected to be the doubly oxidized dideprotonated cobalt species (3-III) because the WOR is initiated at the potential of 32+/33+ (Ecat) according to the CV results (c.f., Figure 4a). Hydrolysis of 3-III leads to the generation of the end-on superoxo CoIIICoIII complex 3a via the transition state 3-IIITS featuring an intramolecular ET from the peroxo subunit to the attacked Co center. This elementary step has a Gibbs free energy of reaction (ΔG) computed as 24.2 kcal/mol. The resulting intermediate, 3a, then undergoes proton-coupled electron transfer (PCET) to form a CoIIICoIII hydroxide superoxide species bearing a diradical ligand (3b) with a ΔG of 34.8 kcal/mol. Subsequently, 3b proceeds through an intramolecular SN2-like attack of the hydroxide ligand on the other Co center to release O2 via the facile transition state structure 3bTS to generate the CoIIICoIII μ-hydroxide intermediate (3c, ΔG: −7.8 eV) that the bpp scaffold features radical character. The subsequent evolution of the O2 from 3c to generate 3d is exergonic by −13.8 kcal/mol, and 3-IV with diradical character on L3 is formed via the one-electron oxidation of 3d at a potential of 0.64 V.

Figure 8.

Figure 8

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Calculated catalytic cycle for the WOR catalyzed by complex 3.62 See Section S16 of the Supporting Information for the considerations of standard Gibbs free energy.

A turnover-limiting transition state occurs by the hydrolysis of 3-IV via the attack of the H2O molecule on one of the CoIII centers (3-IVTS) to generate the corresponding CoIIICoIII aqua hydroxide complex 3e with a ΔG of −1.2 kcal/mol (3-IV → 3e).62 This finding is consistent with the experimental observation of 3-IV in CSI-MS (Figure 7c) and the second-order rate law from CV (Table 1). The deprotonation of 3e produces the dihydroxide species 3f with a corresponding ΔG of 34.4 kcal/mol (equivalent to a pKa of 24). 3f then undergoes stepwise ET/PT to yield 3h with ΔG values of 43.4 and 24.4 kcal/mol, respectively (pKa of 17). Lastly, 3-III is regenerated from 3h via a PCET with a ΔG of 11.5 kcal/mol to complete the catalytic cycle.

As suggested by the KIE and [OH–] dependence results as well as the CSI-MS evidence (Section S14 of the Supporting Information), the reaction pathways of 1 and 3 were found to diverge from the [(L2•+)Co2III(μ-OH), L = L3 or L1]4+ intermediate (3-IV: L = L3). For 3, the second H2O molecule attacks one of the Co centers is proposed as the TLS (3-IV → 3e), as abovementioned. In the case of 1, conversely, the deprotonation of [(L12•+)Co2III(μ-OH)]4+ (1-IV) occurs first to generate [(L12•+)Co2III(μ-O)]3+ (1-V), followed by the attack of H2O to afford the corresponding dihydroxo intermediate, [(L12•+)Co2III(OH)2]3+ (1f). The computational result suggests that 3-IV possesses an unfavorable calculated pKa of 33.0 (equivalent to the standard Gibbs free energy of deprotonation of 43.2 kcal/mol), higher than that of [(L12•+)Co2III(μ-OH)]4+ (1-IV, pKa = 27.3) by nearly 6 orders of magnitude (Section S16 of the Supporting Information). This difference in the pKa values is expected due to the electron-donating methoxy substituents of L3, rendering the CoIII centers less Lewis acidic and, thus, the μ-hydroxy motif less acidic. The increased acidity of 1-IV may lead to the formation of the deprotonated species 1-V, indicated by CSI-MS ([L1Co2III(μ-O)]+, m/z 897.25; section S14 of the Supporting Information), leading to the divergent mechanism of 1 and 3.

Mechanistic Insights into the log(TOFmax)–η Relationship

The TOF and overpotential (η) are regarded as normalized descriptors for representing the performance of a catalyst in terms of its kinetics and thermodynamics, respectively.27,63−65 To understand the effectiveness of the WOCs examined in this study, their TOF and η values were measured and assessed using the log(TOFmax)–η relationship.27,36 For this purpose, the TOFmax values under different conditions were calculated using eq 2 (see Table 3), and the analysis was performed at a scan rate where the ic exhibits a scan-rate independent regime (i.e., ≥800 mV/s, Section S9 of the Supporting Information).52,53,66 Within this regime, the ic value represents the kinetics of the chemical steps instead of substrate diffusion to the electrode surface; therefore, the resulting TOF can indicate the intrinsic catalytic activity of the molecular (electro)catalyst.27 Additionally, the WOR overpotential was estimated by following the literature protocol (Section S10 of the Supporting Information).67 The open-circuit potential (OCP) of H+/H2 was first measured under the experimental conditions of interest, and subsequently, the corresponding thermodynamic potential of H2O/O2 (EH2O/O2) was estimated by adding 1.23 V to the H+/H2 OCP.68,69 The η were then evaluated using eq 6, where Ecat is the catalytic potential initiating the WOR (i.e., 22+/23+ and 32+/33+):53

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Table 3. Kinetic and Thermodynamic WOR Data for Complexes 1–3.

entry catalytic condition catalyst log(TOFmax, s–1) Ecat (mV)a,b EH2O/O2 (mV)a η (mV)a
a 5 M H2O, 0 mM NaOH 1c 1.31 1090 390 700
2 1.11 1010 390 620
3 1.23 890 390 500
b 5 M H2O, 10 mM NaOH 1c 1.54 820 20 800
2 1.23 730 20 710
3 1.33 370 20 350
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a

Potentials are given relative to Fc+/0.

b

Redox couple: [cat]3+/[cat]4+.

c

The ic values used to calculate the TOFmax of 1 were retrieved from ref (30), but the value of Dcat was determined from DOSY experiments (Section S4d of the Supporting Information).

Correlating log(TOFmax) with η for catalysts 1–3 in neutral and alkaline conditions shows a plateau of TOFmax values across the range of ∼400 mV overpotential (Scheme 1a). Although the TLS of complexes 1 and 3 differ under neutral and alkaline conditions,30 Eyring analysis revealed that 1–3 coincidentally features nearly identical Gibbs energies of activation (ΔG‡, 17.5–17.9 kcal/mol, Section S15 of the Supporting Information). Consequently, TOFmax becomes independent of the driving force under these catalytic conditions, allowing a low-overpotential WOR to take place (i.e., <400 mV) without compromising the catalytic turnover. Although the η of complexes 1 and 2 under alkaline conditions are higher than those under neutral conditions (1b vs 1a, 2b vs 2a), intriguingly, complex 3 behaves oppositely (3b vs 3a). This result is due to the cathodic shift in Ecat of 3 under alkaline conditions, which is greater than that of EH2O/O2, making the low-overpotential WOR accessible (Scheme 1b).

Scheme 1. (a) Correlations between log(TOFmax) versus η for the WOR catalyzed by 1–3. (b) Under alkaline conditions, the cathodic shift in Ecat of 3 is more significant than that of EH2O/O2, indicating that WOR can be conducted at a lower η. (c) Turnover-limiting and overpotential-determining steps in an EC′ mechanism are depicted. (d) Comparison of turnover-limiting and overpotential-determining steps of 1 and 3.

Scheme 1

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The catalyst resting states, 1-IV and 3-IV, are proposed on the basis of CSI-MS evidence (Section S14 of the Supporting Information).

A possible rationale for this log(TOFmax)–η plot could be obtained based on the typical EC′ mechanism, as illustrated in Scheme 1c.66,70 In this mechanism, E represents a catalysis-initiating redox couple (Ecat) related to the determination of η, and C′ represents an irreversible chemical step, namely, the TLS. For an MWOC containing electron-donating substituents in its primary coordination sphere, Ecat is expected to exhibit a cathodic shift due to the increased electron density on the MWOC. Although only a lower η is required to drive the WOR in this instance, an offsetting effect occurs on the successive C′ step. Catalyst regeneration then becomes more challenging because of the more reducing nature of the MWOC counterpart, leading to lower kcat and TOFmax values. By altering the TLS and Ecat in the catalytic cycle (Scheme 1d), the present dimeric cobalt complexes are capable of retaining commendable activities at low overpotentials (Scheme 1a). Overall, it was demonstrated that the trade-off between the TOF and η can be circumvented through the use of noninnocent ligands, such as H2L3, with assistance from the electronic effect.

Conclusions

A family of catalysts (1–3) bearing bis(benzimidazole)pyrazolide ligands (H2L1–H2L3) is reported for use in the electrochemical WOR and to demonstrate how ligand design can be utilized to retain good water oxidation activity at low overpotentials. Collectively, the obtained data suggest that the ligand scaffold incorporated into catalysts 1–3 exhibits a noninnocent behavior, in which ligand-centered oxidation is an essential component in the 4e–/4H+ oxidation of H2O. Under alkaline conditions, the kinetic studies reveal that 3 has a different reaction pathway, TLS, and overpotential-determining step compared with complex 1. The reaction intermediates and catalyst resting states in the WOR catalyzed by 3 were evaluated using X-ray crystallography, electrochemical, and spectroscopic methods. The spectroscopic evidence indicated that multielectron/multiproton transfer occurs during the reaction, and the characterized cobalt species provided a vision of the reaction mechanism. Supplemented by density functional theory investigations, we deduced that the TLS involved the nucleophilic attack of water on [(L32•+)Co2III(μ-OH)]4+ (3-IV), which differs from the base-assisted deprotonation of dimeric [(L12•+)Co2III(OH)]4+ (I–IV), which was identified as the TLS of the WOR catalyzed by 1.

The present results highlight that the mutual influence between the electronic effect and the noninnocent character of the ligand on the MWOCs can manipulate the PT and ET steps, leading to different reaction mechanisms. Furthermore, Eyring analysis demonstrates that complexes 1–3 exhibit similar Gibbs free energies of activation. Moreover, a plot of log(TOFmax) against η showed a plateau region wherein the TOFmax is independent of η over 400 mV, which was attributed to the different overpotential-determining steps of 1–3. These observations provide valuable insights into catalyst design and are expected to lead to catalysts suitable for operation at low overpotentials.

Experimental Section

General Methods and Materials

All commercially purchased reagents were utilized in their original state. The required precursors and metal salts were purchased from Sigma-Aldrich, Acros Organics, and TCI. 1H NMR and 13C NMR spectra were recorded on a Bruker Avance III spectrometer (400 and 500 MHz for 1H NMR, 100 and 125.7 MHz for 13C NMR). 1H diffusion-ordered spectroscopy (DOSY) NMR spectra were recorded on a JEOL ECZ500R/S1 (500 MHz). Crystal evaluation and data collection were carried out on a Bruker X8 APEX Quazar SMART APEXII diffractometer with Mo Kα (λ = 0.71073 Å) radiation and the diffractometer and Rigaku XtaLAB Synergy R, DW system, HyPix-Arc 150 with Cu Kα (λ = 1.54184 Å) radiation and the diffractometer. ESI and EI mass spectra were recorded on a VARIAN 901-MS and a JEOL JMS-700 mass spectrometer, respectively. SEM images and EDX analysis were obtained with a JEOL JSM-7000 FESEM instrument equipped with an EDX detector. EPR spectra were measured on a Bruker EPR-plus.

Synthesis of Ligands and Cobalt Complexes

The required precursors, ligands (H2L1, H2L2, H2L3), and complexes (1, 2, 3) were synthesized according to previously reported literature procedures with necessary modification. The details are given in the Supporting Information with characterization data. Complex 1 was obtained as a purple solid, yielding 140 mg (55%). Complex 2 was synthesized according to the procedure employed to synthesize complex 1. The crude was purified on an alumina column and eluted with 50 mM KPF6 in acetone to obtain a purple solid, yielding 48 mg (23%). Complex 3 was synthesized using the procedure employed to synthesize complex 1, but precursors were added reversely. The crude product could be purified on an alumina column and eluted with 50 mM KPF6 in acetone to obtain the purple solid, yielding 46 mg (40%).

Electrochemistry

Electrochemical experiments were conducted by using a PalmSens4 potentiostat connected to a computer with PSTrace software, employing a three-electrode setup. The working electrode was a 3.0 mm diameter glassy carbon disk with a platinum wire as the auxiliary electrode and a silver wire pseudoreference as the reference electrode. The pseudoreference for MeCN solvents included 100 mM [NBu4][PF6] as a supporting electrolyte and 10 mM AgNO3. The working electrode was polished with 0.05 μm alumina on a wetted Buehler felt pad between each CV experiment. All voltammograms were internally referenced to the redox potential of Fc+/0. The laboratory temperature was maintained at 25 ± 2 °C.

Cold-Spray Ionization Mass Spectrometry (CSI-MS)

The catalytically relevant cobalt intermediates were detected by a CSI-MS system, combining a home-built cold-spray ionization source and a linear ion trap mass spectrometer (LTQ XL, Thermo Fisher Scientific) to detect the unstable intermediates. The entire sample transport line was constructed with a dual-layered structure to regulate the temperature of the sample.

Acknowledgments

This research was supported by the Ministry of Science and Technology (Grant No. MOST 109-2113-M-007-024-MY3). The authors thank Hui-Chi Tan, Pei-Lin Chen, Hsin-Ru Wu, and Chu-Ting Han from the Instrumentation Centre of National Tsing Hua University for their help with the NMR measurements, X-ray data collection and structural determination, ESI-MS measurements, and EDXS measurements, respectively. The authors would also like to thank Hsiu-Ni Huang from the Instrumentation Centre of the National Taiwan Normal University for carrying out the EPR analysis and You-Song Cheng from the Instrumentation Centre of National Yang-Ming Chiao Tung University for NMR measurements.

Data Availability Statement

The data supporting this study can be accessed in the published article and its accompanying Supporting Information.

Supporting Information Available

The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/acsorginorgau.3c00061.

  • General instrumentation methods and experimental details, including characterization data such as electrochemical analysis of the turnover frequencies, H2O2 quantitation, OCP measurements for the determination of EH/H2+, thermodynamic potentials for EO2/H2O2, half-wave potentials of the cobalt catalysts, kinetic data, 1H and 13C NMR shift values, ESI-MS/HRMS, 1H and 13C NMR spectra of the synthesized compounds, UV–visible spectra and single-crystal XRD data for complexes (PDF)

The authors declare no competing financial interest.

Supplementary Material

gg3c00061_si_001.pdf (10.2MB, pdf)

References

  1. Liu J.; Mooney H.; Hull V.; Davis S. J.; Gaskell J.; Hertel T.; Lubchenco J.; Seto K. C.; Gleick P.; Kremen C.; Li S. Systems integration for global sustainability. Science 2015, 347, 1258832 10.1126/science.1258832. [DOI] [PubMed] [Google Scholar]
  2. Berrang-Ford L.; Biesbroek R.; Ford J. D.; Lesnikowski A.; Tanabe A.; Wang F. M.; Chen C.; Hsu A.; Hellmann J. J.; Pringle P.; et al. Tracking global climate change adaptation among governments. Nat. Clim. Change 2019, 9, 440–449. 10.1038/s41558-019-0490-0. [DOI] [Google Scholar]
  3. Kaufman D. S.; Broadman E. Revisiting the Holocene global temperature conundrum. Nature 2023, 614, 425–435. 10.1038/s41586-022-05536-w. [DOI] [PubMed] [Google Scholar]
  4. Chow J.; Kopp R. J.; Portney P. R. Energy resources and global development. Science 2003, 302, 1528–1531. 10.1126/science.1091939. [DOI] [PubMed] [Google Scholar]
  5. Davis S. J.; Lewis N. S.; Shaner M.; Aggarwal S.; Arent D.; Azevedo I. L.; Benson S. M.; Bradley T.; Brouwer J.; Chiang Y.-M.; et al. Net-zero emissions energy systems. Science 2018, 360, eaas9793 10.1126/science.aas9793. [DOI] [PubMed] [Google Scholar]
  6. Turner J. A. Sustainable hydrogen production. Science 2004, 305, 972–974. 10.1126/science.1103197. [DOI] [PubMed] [Google Scholar]
  7. Stern A. G. A new sustainable hydrogen clean energy paradigm. Int. J. Hydrogen Energy 2018, 43, 4244–4255. 10.1016/j.ijhydene.2017.12.180. [DOI] [Google Scholar]
  8. Megía P. J.; Vizcaíno A. J.; Calles J. A.; Carrero A. Hydrogen production technologies: From fossil fuels toward renewable sources. A mini review. Energy Fuels 2021, 35, 16403–16415. 10.1021/acs.energyfuels.1c02501. [DOI] [Google Scholar]
  9. Osman A. I.; Mehta N.; Elgarahy A. M.; Hefny M.; Al-Hinai A.; Al-Muhtaseb A. a. H.; Rooney D. W. Hydrogen production, storage, utilisation and environmental impacts: a review. Environ. Chem. Lett. 2022, 20, 1–36. 10.1007/s10311-021-01322-8. [DOI] [Google Scholar]
  10. Usman M. R. Hydrogen storage methods: Review and current status. Renew. Sust. Energ. Rev. 2022, 167, 112743 10.1016/j.rser.2022.112743. [DOI] [Google Scholar]
  11. Jena P. Materials for hydrogen storage: past, present, and future. J. Phys. Chem. Lett. 2011, 2, 206–211. 10.1021/jz1015372. [DOI] [Google Scholar]
  12. Schlapbach L.; Züttel A. Hydrogen-storage materials for mobile applications. Nature 2001, 414, 353–358. 10.1038/35104634. [DOI] [PubMed] [Google Scholar]
  13. Li J.; Güttinger R.; Moré R.; Song F.; Wan W.; Patzke G. R. Frontiers of water oxidation: the quest for true catalysts. Chem. Soc. Rev. 2017, 46, 6124–6147. 10.1039/C7CS00306D. [DOI] [PubMed] [Google Scholar]
  14. Ye S.; Ding C.; Liu M.; Wang A.; Huang Q.; Li C. Water oxidation catalysts for artificial photosynthesis. Adv. Mater. 2019, 31, 1902069 10.1002/adma.201902069. [DOI] [PubMed] [Google Scholar]
  15. Li J.; Triana C. A.; Wan W.; Saseendran D. A.; Zhao Y.; Balaghi S. E.; Heidari S.; Patzke G. Molecular and heterogeneous water oxidation catalysts: recent progress and joint perspectives. Chem. Soc. Rev. 2021, 50, 2444–2485. 10.1039/D0CS00978D. [DOI] [PubMed] [Google Scholar]
  16. Shih A. J.; Monteiro M. C.; Dattila F.; Pavesi D.; Philips M.; da Silva A. H.; Vos R. E.; Ojha K.; Park S.; van der Heijden O.; et al. Water electrolysis. Nat. Rev. Methods Primers 2022, 2, 84. 10.1038/s43586-022-00164-0. [DOI] [Google Scholar]
  17. Qi J.; Zhang W.; Cao R. Solar-to-hydrogen energy conversion based on water splitting. Adv. Energy Mater. 2018, 8, 1701620 10.1002/aenm.201701620. [DOI] [Google Scholar]
  18. Eisenberg R. Rethinking water splitting. Science 2009, 324, 44–45. 10.1126/science.1172247. [DOI] [PubMed] [Google Scholar]
  19. Gagliardi C. J.; Vannucci A. K.; Concepcion J. J.; Chen Z.; Meyer T. J. The role of proton coupled electron transfer in water oxidation. Energy Environ. Sci. 2012, 5, 7704–7717. 10.1039/c2ee03311a. [DOI] [Google Scholar]
  20. Brudvig G. W.; Reek J. N.; Sakai K.; Spiccia L.; Sun L. Catalytic systems for water splitting. ChemPlusChem 2016, 81, 1017–1019. 10.1002/cplu.201600436. [DOI] [PubMed] [Google Scholar]
  21. Barber J. Photosynthetic energy conversion: natural and artificial. Chem. Soc. Rev. 2009, 38, 185–196. 10.1039/B802262N. [DOI] [PubMed] [Google Scholar]
  22. Tahir M.; Pan L.; Idrees F.; Zhang X.; Wang L.; Zou J.-J.; Wang Z. L. Electrocatalytic oxygen evolution reaction for energy conversion and storage: a comprehensive review. Nano Energy 2017, 37, 136–157. 10.1016/j.nanoen.2017.05.022. [DOI] [Google Scholar]
  23. Roger I.; Shipman M. A.; Symes M. D. Earth-abundant catalysts for electrochemical and photoelectrochemical water splitting. Nat. Rev. Chem. 2017, 1, 0003. 10.1038/s41570-016-0003. [DOI] [Google Scholar]
  24. Kärkäs M. D.; Åkermark B. Water oxidation using earth-abundant transition metal catalysts: opportunities and challenges. Dalton Trans. 2016, 45, 14421–14461. 10.1039/C6DT00809G. [DOI] [PubMed] [Google Scholar]
  25. Zhang L. H.; Mathew S.; Hessels J.; Reek J. N.; Yu F. Homogeneous Catalysts Based on First-Row Transition-Metals for Electrochemical Water Oxidation. ChemSusChem 2021, 14, 234–250. 10.1002/cssc.202001876. [DOI] [PMC free article] [PubMed] [Google Scholar]
  26. Kondo M.; Tatewaki H.; Masaoka S. Design of molecular water oxidation catalysts with earth-abundant metal ions. Chem. Soc. Rev. 2021, 50, 6790–6831. 10.1039/D0CS01442G. [DOI] [PubMed] [Google Scholar]
  27. Hsu W. C.; Wang Y. H. Homogeneous Water Oxidation Catalyzed by First-Row Transition Metal Complexes: Unveiling the Relationship between Turnover Frequency and Reaction Overpotential. ChemSusChem 2022, 15, e202102378 10.1002/cssc.202102378. [DOI] [PubMed] [Google Scholar]
  28. Sala X.; Romero I.; Rodríguez M.; Escriche L.; Llobet A. Molecular Catalysts that Oxidize Water to Dioxygen. Angew. Chem., Int. Ed. 2009, 48, 2842–2852. 10.1002/anie.200802659. [DOI] [PubMed] [Google Scholar]
  29. Wasylenko D. J.; Palmer R. D.; Berlinguette C. P. Homogeneous water oxidation catalysts containing a single metal site. Chem. Commun. 2013, 49, 218–227. 10.1039/C2CC35632E. [DOI] [PubMed] [Google Scholar]
  30. Hsu W. C.; Zeng W. Q.; Lu I. C.; Yang T.; Wang Y. H. Dinuclear Cobalt Complexes for Homogeneous Water Oxidation: Tuning Rate and Overpotential through the Non-Innocent Ligand. ChemSusChem 2022, 15, e202201317 10.1002/cssc.202201317. [DOI] [PubMed] [Google Scholar]
  31. Sala X.; Romero I.; Rodríguez M.; Escriche L.; Llobet A. Molekulare Katalysatoren für die Oxidation von Wasser zu Disauerstoff. Angew. Chem. 2009, 121, 2882–2893. 10.1002/ange.200802659. [DOI] [Google Scholar]
  32. Blakemore J. D.; Crabtree R. H.; Brudvig G. W. Molecular Catalysts for Water Oxidation. Chem. Rev. 2015, 115, 12974–13005. 10.1021/acs.chemrev.5b00122. [DOI] [PubMed] [Google Scholar]
  33. Zhang B.; Sun L. Artificial photosynthesis: opportunities and challenges of molecular catalysts. Chem. Soc. Rev. 2019, 48, 2216–2264. 10.1039/C8CS00897C. [DOI] [PubMed] [Google Scholar]
  34. Ghosh A.; Dasgupta S.; Kundu A.; Mandal S. The impact of secondary coordination sphere engineering on water oxidation reactivity catalysed by molecular ruthenium complexes: a next-generation approach to develop advanced catalysts. Dalton Trans. 2022, 51, 10320–10337. 10.1039/D2DT01124G. [DOI] [PubMed] [Google Scholar]
  35. Mondal R.; Guin A. K.; Chakraborty G.; Paul N. D. Metal–ligand cooperative approaches in homogeneous catalysis using transition metal complex catalysts of redox noninnocent ligands. Org. Biomol. Chem. 2022, 20, 296–328. 10.1039/D1OB01153G. [DOI] [PubMed] [Google Scholar]
  36. Ma H.-C.; Hsiao S.-C.; Wang Y.-H. Tuning primary and secondary coordination spheres of ruthenium complexes for the homogeneous water oxidation reaction: a perspective from catalytic activity and overpotential. Catal. Sci. Technol. 2023, 13, 1598–1622. 10.1039/D2CY02109A. [DOI] [Google Scholar]
  37. Thammavongsy Z.; Mercer I. P.; Yang J. Y. Promoting proton coupled electron transfer in redox catalysts through molecular design. Chem. Commun. 2019, 55, 10342–10358. 10.1039/C9CC05139B. [DOI] [PubMed] [Google Scholar]
  38. Li W.; Li F.; Yang H.; Wu X.; Zhang P.; Shan Y.; Sun L. A bio-inspired coordination polymer as outstanding water oxidation catalyst via second coordination sphere engineering. Nat. Commun. 2019, 10, 5074 10.1038/s41467-019-13052-1. [DOI] [PMC free article] [PubMed] [Google Scholar]
  39. Grützmacher H. Cooperating ligands in catalysis. Angew. Chem., Int. Ed. 2008, 47, 1814–1818. 10.1002/anie.200704654. [DOI] [PubMed] [Google Scholar]
  40. Wang P.; Liang G.; Reddy M. R.; Long M.; Driskill K.; Lyons C.; Donnadieu B.; Bollinger J. C.; Webster C. E.; Zhao X. Electronic and steric tuning of catalytic H2 evolution by cobalt complexes with pentadentate polypyridyl-amine ligands. J. Am. Chem. Soc. 2018, 140, 9219–9229. 10.1021/jacs.8b05108. [DOI] [PubMed] [Google Scholar]
  41. Garrido-Barros P.; Funes-Ardoiz I.; Drouet S.; Benet-Buchholz J.; Maseras F.; Llobet A. Redox non-innocent ligand controls water oxidation overpotential in a new family of mononuclear Cu-based efficient catalysts. J. Am. Chem. Soc. 2015, 137, 6758–6761. 10.1021/jacs.5b03977. [DOI] [PubMed] [Google Scholar]
  42. Lyaskovskyy V.; de Bruin B. Redox Non-Innocent Ligands: Versatile New Tools to Control Catalytic Reactions. ACS Catal. 2012, 2, 270–279. 10.1021/cs200660v. [DOI] [Google Scholar]
  43. Singh B.; Indra A. Role of redox active and redox non-innocent ligands in water splitting. Inorg. Chim. Acta 2020, 506, 119440 10.1016/j.ica.2020.119440. [DOI] [Google Scholar]
  44. Sutradhar M.; Pombeiro A. J. L.; da Silva J. A. L. Water oxidation with transition metal catalysts with non-innocent ligands and its mechanisms. Coord. Chem. Rev. 2021, 439, 213911 10.1016/j.ccr.2021.213911. [DOI] [Google Scholar]
  45. Anferov S. W.; Czaikowski M. E.; Anderson J. S. Metal–ligand cooperative transfer of protons and electrons. Trends Chem. 2021, 3, 993–996. 10.1016/j.trechm.2021.10.002. [DOI] [PMC free article] [PubMed] [Google Scholar]
  46. Nakada A.; Matsumoto T.; Chang H.-C. Redox-active ligands for chemical, electrochemical, and photochemical molecular conversions. Coord. Chem. Rev. 2022, 473, 214804 10.1016/j.ccr.2022.214804. [DOI] [Google Scholar]
  47. Rigsby M. L.; Mandal S.; Nam W.; Spencer L. C.; Llobet A.; Stahl S. S. Cobalt analogs of Ru-based water oxidation catalysts: overcoming thermodynamic instability and kinetic lability to achieve electrocatalytic O2 evolution. Chem. Sci. 2012, 3, 3058–3062. 10.1039/c2sc20755a. [DOI] [Google Scholar]
  48. Davenport T. C.; Ahn H. S.; Ziegler M. S.; Tilley T. D. A molecular structural analog of proposed dinuclear active sites in cobalt-based water oxidation catalysts. Chem. Commun. 2014, 50, 6326–6329. 10.1039/c3cc46865h. [DOI] [PubMed] [Google Scholar]
  49. Ishizuka T.; Watanabe A.; Kotani H.; Hong D.; Satonaka K.; Wada T.; Shiota Y.; Yoshizawa K.; Ohara K.; Yamaguchi K.; et al. Homogeneous photocatalytic water oxidation with a dinuclear CoIII–pyridylmethylamine complex. Inorg. Chem. 2016, 55, 1154–1164. 10.1021/acs.inorgchem.5b02336. [DOI] [PubMed] [Google Scholar]
  50. Fukuzumi S.; Mandal S.; Mase K.; Ohkubo K.; Park H.; Benet-Buchholz J.; Nam W.; Llobet A. Catalytic four-electron reduction of O2 via rate-determining proton-coupled electron transfer to a dinuclear cobalt-μ-1, 2-peroxo complex. J. Am. Chem. Soc. 2012, 134, 9906–9909. 10.1021/ja303674n. [DOI] [PubMed] [Google Scholar]
  51. More M. B.; Ray D.; Armentrout P. Intrinsic affinities of alkali cations for 15-crown-5 and 18-crown-6: bond dissociation energies of gas-phase M+–crown ether complexes. J. Am. Chem. Soc. 1999, 121, 417–423. 10.1021/ja9823159. [DOI] [Google Scholar]
  52. Costentin C.; Drouet S.; Robert M.; Savéant J.-M. Turnover numbers, turnover frequencies, and overpotential in molecular catalysis of electrochemical reactions. Cyclic voltammetry and preparative-scale electrolysis. J. Am. Chem. Soc. 2012, 134, 11235–11242. 10.1021/ja303560c. [DOI] [PubMed] [Google Scholar]
  53. Rountree E. S.; McCarthy B. D.; Eisenhart T. T.; Dempsey J. L. Evaluation of homogeneous electrocatalysts by cyclic voltammetry. Inorg. Chem. 2014, 53, 9983–10002. 10.1021/ic500658x. [DOI] [PubMed] [Google Scholar]
  54. Eisenberg G. Colorimetric Determination of Hydrogen Peroxide. Ind. Eng. Chem., Anal. Ed. 1943, 15, 327–328. 10.1021/i560117a011. [DOI] [Google Scholar]
  55. Lee Y.; Park G. Y.; Lucas H. R.; Vajda P. L.; Kamaraj K.; Vance M. A.; Milligan A. E.; Woertink J. S.; Siegler M. A.; Narducci Sarjeant A. A.; et al. Copper (I)/O2 Chemistry with Imidazole Containing Tripodal Tetradentate Ligands Leading to μ-1,2-Peroxo–Dicopper (II) Species. Inorg. Chem. 2009, 48, 11297–11309. 10.1021/ic9017695. [DOI] [PMC free article] [PubMed] [Google Scholar]
  56. Wu X.; Li F.; Zhang B.; Sun L. Molecular complexes in water oxidation: Pre-catalysts or real catalysts. J. Photochem. Photobiol., C 2015, 25, 71–89. 10.1016/j.jphotochemrev.2015.07.002. [DOI] [Google Scholar]
  57. Lee K. J.; McCarthy B. D.; Dempsey J. L. On decomposition, degradation, and voltammetric deviation: the electrochemist’s field guide to identifying precatalyst transformation. Chem. Soc. Rev. 2019, 48, 2927–2945. 10.1039/C8CS00851E. [DOI] [PubMed] [Google Scholar]
  58. McAlpin J. G.; Stich T. A.; Ohlin C. A.; Surendranath Y.; Nocera D. G.; Casey W. H.; Britt R. D. Electronic Structure Description of a [Co(III)3Co(IV)O4] Cluster: A Model for the Paramagnetic Intermediate in Cobalt-Catalyzed Water Oxidation. J. Am. Chem. Soc. 2011, 133, 15444–15452. 10.1021/ja202320q. [DOI] [PubMed] [Google Scholar]
  59. Givaja G.; Volpe M.; Edwards M. A.; Blake A. J.; Wilson C.; Schröder M.; Love J. B. Dioxygen reduction at dicobalt complexes of a Schiff base calixpyrrole ligand. Angew. Chem., Int. Ed. 2007, 46, 584–586. 10.1002/anie.200603201. [DOI] [PubMed] [Google Scholar]
  60. Louka F. R.; Nguyen L. T.; Albering J. H.; Mautner F. A.; Massoud S. S. Unprecedented formation of a doubly bridged μ-peroxo-μ-pyrazolyldicarboxylato-dicobalt (III) complex. Inorg. Chem. Commun. 2012, 15, 269–271. 10.1016/j.inoche.2011.10.042. [DOI] [Google Scholar]
  61. Kon H.; Nagata T. New ternary ligands consisting of an N4 bridging ligand and two terpyridines, and their Co (II) and Ni (II) dinuclear complexes. Structure, redox properties, and reaction with acid. Dalton Trans. 2013, 42, 5697–5705. 10.1039/c3dt33011g. [DOI] [PubMed] [Google Scholar]
  62. The calculated Gibbs free energy of activation was found to be 36.0 kcal/mol. The observed discrepancy between the calculated and experimental value is a well-known issue in the condensed phase due to the neglect of the explicit solvent–solute interactions, exacerbated by highly polar and H-bond donor solvent molecules such as H2O. See Section S16 of the Supporting Information for the considerations of Standard Gibbs Free Energy.
  63. Costentin C.; Savéant J.-M. Towards an intelligent design of molecular electrocatalysts. Nat. Rev. Chem. 2017, 1, 0087. 10.1038/s41570-017-0087. [DOI] [Google Scholar]
  64. Anantharaj S.; Karthik P. E.; Noda S. The Significance of Properly Reporting Turnover Frequency in Electrocatalysis Research. Angew. Chem., Int. Ed. 2021, 60, 23051–23067. 10.1002/anie.202110352. [DOI] [PMC free article] [PubMed] [Google Scholar]
  65. Stratakes B. M.; Dempsey J. L.; Miller A. J. Determining the Overpotential of Electrochemical Fuel Synthesis Mediated by Molecular Catalysts: Recommended Practices, Standard Reduction Potentials, and Challenges. ChemElectroChem 2021, 8, 4161–4180. 10.1002/celc.202100576. [DOI] [Google Scholar]
  66. Costentin C.; Savéant J. M. Multielectron, Multistep Molecular Catalysis of Electrochemical Reactions: Benchmarking of Homogeneous Catalysts. ChemElectroChem 2014, 1, 1226–1236. 10.1002/celc.201300263. [DOI] [Google Scholar]
  67. Roberts J. A. S.; Bullock R. M. Direct Determination of Equilibrium Potentials for Hydrogen Oxidation/Production by Open Circuit Potential Measurements in Acetonitrile. Inorg. Chem. 2013, 52, 3823–3835. 10.1021/ic302461q. [DOI] [PubMed] [Google Scholar]
  68. Mondal B.; Chattopadhyay S.; Dey S.; Mahammed A.; Mittra K.; Rana A.; Gross Z.; Dey A. Elucidation of Factors That Govern the 2e–/2H+ vs 4e–/4H+ Selectivity of Water Oxidation by a Cobalt Corrole. J. Am. Chem. Soc. 2020, 142, 21040–21049. 10.1021/jacs.0c08654. [DOI] [PubMed] [Google Scholar]
  69. Marenich A. V.; Cramer C. J.; Truhlar D. G. Universal solvation model based on solute electron density and on a continuum model of the solvent defined by the bulk dielectric constant and atomic surface tensions. J. Phys. Chem. B 2009, 113, 6378–6396. 10.1021/jp810292n. [DOI] [PubMed] [Google Scholar]
  70. Martin D. J.; McCarthy B. D.; Rountree E. S.; Dempsey J. L. Qualitative extension of the EC′ Zone Diagram to a molecular catalyst for a multi-electron, multi-substrate electrochemical reaction. Dalton Trans. 2016, 45, 9970–9976. 10.1039/C6DT00302H. [DOI] [PubMed] [Google Scholar]

Associated Data

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Supplementary Materials

gg3c00061_si_001.pdf (10.2MB, pdf)

Data Availability Statement

The data supporting this study can be accessed in the published article and its accompanying Supporting Information.

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