Bioinspired Copper(II) Complexes With N,S‐Donor Ligands: Structure–Activity Relationships and Mechanistic Insights for Aqueous Hydrogen Evolution Reaction

Abstract

The design and development of bioinspired coordination compounds are crucial for laying the foundation to deliver low‐cost, energy‐efficient materials that promote energy sustainability. In this work, isostructural copper complexes, Cu‐L_N2S2 and Cu‐L_N3S, in tune with the donor centers in the ligand backbone, (L_N2S2 = 1,2‐bis((pyridin‐2‐ylmethyl)thio)ethane and L_N3S =N‐(pyridin‐2‐ylmethyl)‐2‐((pyridine‐2‐ylmethyl)thio)ethane‐1‐amine) have been prepared, structurally characterized, and evaluated for their electrocatalytic fate towards the sustainable hydrogen production activities in water. Single‐crystal X‐ray crystallography reveals that the copper centers in both complexes adopt a distorted square pyramidal geometry. Cu‐L_N2S2 exhibits superior electrocatalytic performance over Cu‐L_N3S in acidic aqueous media, achieving an outstanding turnover frequency (TOF) of 2.90 × 10³ s⁻¹ with 96% Faradaic efficiency. Detailed mechanistic insights, supported by spectroscopic, analytical, and DFT calculations, reveal divergent HER pathways: Cu‐L_N3S follows an ECEC mechanism, while Cu‐L_N2S2 operates via a CECE sequence. Protonation and reduction site analyses highlight the critical role of mixed hard–soft donor environments in modulating redox behavior and promoting hydride formation for efficient hydrogen production. The greater structural distortion arises from the higher number of sulfur donors in the ligand environment of Cu‐L _N2S2 , which in turn promotes the formation of the observed coordination geometry feasible for sustainable electrocatalytic evolution of hydrogen in water.

This work elucidates the bioinspired design of a water‐stable square pyramidal copper complex that exhibits excellent electrocatalytic activity in water through proton‐electron coupled catalytic pathways for sustainable hydrogen production.

1. Introduction

A sustainable future rest on the foundation of clean and reliable energy. Amid growing global energy demand, electrocatalytic hydrogen production stands out as a premier green fuel solution [1, 2, 3, 4, 5]. While platinum remains a benchmark [6], its rarity and cost underscore the urgent need for efficient, robust, and economical alternatives [7, 8, 9]. Drawing inspiration from nature's hydrogenase enzymes with [Ni‐Fe [10] Fe–Fe [11, 12], and [Fe [11] active centers, researchers are designing earth‐abundant metal‐based electrocatalysts, enhanced by versatile and tuneable ligand frameworks. Since then, significant progress has been made using coordination‐driven strategies involving (N, O, S, P)‐donor ligands [12, 13, 14, 15] and iron [16, 17, 18, 19], nickel [20, 21, 22, 23, 24, 25], cobalt [26, 27, 28, 29, 30, 31, 32, 33], and molybdenum [34] metal ions. Researchers such as those led by Du [35], Sally Brooker group [36], Inke Siewert [37], Pannee Leeladee [38], and Sumanta Kumar Padhi [13] have advanced HER catalysis using N₂, N₃, or N₄ donor units such as amines, pyridines, and imidazoles. Despite the significant progress, the development of copper‐based catalysts featuring (N, S)‐donor ligands remains largely untapped, presenting a promising frontier for designing efficient, robust, and sustainable systems for hydrogen evolution.

Copper, a bio‐essential and earth‐abundant metal, is both cost‐effective and highly versatile in energy and organic transformations [39, 40, 41, 42] While many copper complexes show HER activity in organic solvents, their reliance on environmentally unfriendly organic solvents [13, 43] stems from the challenge of accessing low‐valent copper states at high reduction potentials. However, rational ligand and donor atom design can finely tune their electronic and structural properties, unlocking progress in the catalytic performance even in aqueous media. Notably, the blend of hard (N) and soft (S) donors enables fine‐tuning of redox behavior and electronic structure, making it a powerful strategy in catalyst design.

Haddad et al. demonstrated the HER activity of a Cu(II) complex incorporating a redox‐active (N,S)‐donor ligand, reporting an overpotential of 0.80 V in organic solvent [44]. However, the lack of performance assessment in aqueous media limits its applicability for sustainable hydrogen production. Likewise, Eisenberg and coworkers explored HER using Ni(II) [45] and Co(II) complexes [46, 47], but their studies were confined to organic or mixed aqueous‐organic systems [48]. These insights emphasize the importance of advancing molecular HER catalysts that are not only efficient but also compatible with aqueous environments, paving the way for greener and more scalable hydrogen generation technologies.

Recently, Zhan and coworkers introduced a series of robust copper complexes bearing N₄‐amine or N₂O₂‐oxamato ligands, which showcased notable HER activity under both acidic and neutral aqueous conditions [49]. However, despite their promising performance, copper complexes featuring (S,S)‐ or (N,S)‐donor ligands remain scarcely explored for homogeneous HER [44, 50, 51]. The significance of N,S‐type coordination arises from its prevalence in biological systems, particularly in metalloenzymes such as hydrogenases, iron–sulfur clusters, and nitrile hydratases. The M—S bond, in contrast to the more common biorelevant M—N bonds (typically found in carboxylates, phenols, and imidazoles), exerts a stronger influence on the properties of the metal center and/or the sulfur atom because of its higher degree of covalency. The synergy of hard (N) and soft (S) donor atoms in a ligand framework allows for precise modulation of redox properties and electronic structure. Building upon these advancements, our group has recently reported, for the first time, a copper complex supported by an N₂S₂ ligand framework that exhibits efficient HER activity specifically in acidic aqueous media [52]. This design not only leverages the redox versatility and proton‐responsive nature of the (N,S)‐donor environment but also offers a promising platform for further mechanistic exploration and structural optimization toward sustainable hydrogen evolution.

Nevertheless, (N, S)‐based electrocatalysts’ literature still lacks a thorough understanding of structure–activity relationships and mechanistic pathways, particularly concerning the influence of sulfur incorporation within the ligand framework. In this context, we have synthesized and comprehensively characterized two structurally analogous bioinspired water‐stable Cu (II) complexes, [Cl−Cu−L_N2S2]ClO₄ (Cu‐L_N3S) and [Cl−Cu−L_N3S]ClO₄ (Cu‐L_N2S2), using combination of techniques including ultraviolet–visible spectroscopy (UV‐Vis), fourier transform infrared spectroscopy (FTIR), nuclear magnetic resonance (¹H and ¹³C NMR), electrospray ionization mass spectrometry (ESI‐MS), electron paramagnetic resonance (EPR), and field emission scanning electron microscopy (FESEM) along with by single‐crystal X‐ray diffraction (SCXRD) for definitive structural elucidation. Subsequently, we critically examined the electrocatalytic performance of these Cu(II) complexes in aqueous media, employing acetic acid (AcOH) as a proton source. Our investigation culminates in the establishment of a clear structure–activity relationship between the two complexes, one featuring an N₂S₂ donor set and the other an N₃S donor set. Furthermore, mechanistic insights reveal that the complexes operate via distinct catalytic cycles, attributed to their differing coordination environments. This study provides a foundational framework that will undoubtedly inspire and inform the future design of Cu‐based electrocatalysts employing N, S‐donor ligands.

2. Results and Discussion

2.1. Synthesis, Characterization, Crystal Structure Description, and Solution Phase Behavior of (Cu‐L_N2S2) and (Cu‐L_N3S)

The ligands, L_N2S2 and L_N3S, were synthesized by a modification to previously reported methods. L_N2S2 was synthesized by reacting 1,2‐ethanedithiol with sodium ethoxide in dry ethanol at 0–5°C, followed by the dropwise addition of 2‐picolyl chloride hydrochloride and subsequent reflux for 2 h. In contrast, L_N3S was synthesized by treating cysteamine hydrochloride with KOH in dry methanol, followed by the addition of 2‐picolyl chloride hydrochloride and refluxing for 4 h. Subsequently, copper(II) complexes, Cu‐L_N2S2 and Cu‐L_N3S, were synthesized by reacting the ligands L_N2S2 and L_N3S with hydrated Cu(ClO₄)₂ in methanol at room temperature. Suitable blue single crystals were obtained through slow evaporation of the reaction mixture. The synthetic route for the ligands and their corresponding copper complexes is illustrated in Scheme 1. Comprehensive synthetic procedures for both the ligands and their corresponding Cu(II) complexes including characterization (¹H, ¹³C NMR and ESI‐MS Figure S1−S6) are provided in the Supplementary Information (SI).

SCHEME 1.

Synthetic procedure of L_N3S, L_N2S2, Cu‐L_N3S, and Cu‐L_N2S2.

The single‐crystal X‐ray diffraction (SCXRD) analysis revealed that the complex Cu‐L_N3S crystallized in the monoclinic space group P2₁/c while Cu‐L_N2S2 crystallized in the triclinic space group P‐1. The structural refinement parameters of the complexes are provided in Table S1. The ORTEP diagram shows that Cu‐L_N2S2 and Cu‐L_N3S adopt distorted square pyramidal coordination geometries with respective geometry index tau (τ) values 0.0344 and 0.0289 [53]. Cu‐L_N3S adopts relatively lesser distortion in square pyramidal geometry where the sulfur S(1), nitrogen N(2A), N(3), and chlorine atom Cl(1) occupy the equatorial positions and the nitrogen atom (N1) resides in the apical position (Figure 1a). On the other hand, in Cu‐L_N2S2, sulfur (S1), nitrogen (N1, N2) and chlorine (Cl1) atoms occupy the equatorial positions, while sulfur (S2) atom resides at the apical position (Figure 1g). In both complexes, dicationic charge of Cu‐centers is compensated by the coordinated chloride and one counter perchlorate ion. The Cu—Cl bond distance of Cu‐L_N2S2 is found to be slightly longer (2.272 Å) than Cu‐L_N3S (2.258 Å). Moreover, for Cu‐L_N2S2, the dihedral angle between the two planes, Cu1‐S1‐C6 and Cu1‐S2‐C9, is 87.93°, whereas the dihedral angle of Cu‐L_N3S between the two planes, Cu1‐S1‐C6 and Cu1‐N2A‐C9, is 65.68° (Figure S7). In Cu‐L_N3S, the smaller dihedral angle (65.68°) allows greater π‐orbital overlap between the donor atoms and the copper center, particularly along the Cu—S and Cu—N bonds. This increases electronic delocalization, enhancing electronic communication through the ligand. In Cu‐L_N2S2, the larger angle (87.93°) results in poor alignment of donor p‐orbitals with the Cu center, reducing orbital overlap and limiting electronic coupling introducing more lability in the reaction conditions. Further, supramolecular framework study of Cu‐L_N3S demonstrates the involvement of strong π ^… π interactions between the pyridine centroids of L_N3S leading to the dimer of Cu‐L_N3S which extends its to grow a 1D supramolecular architecture via Cl ^… H interactions (Figure 1b). In contrast, the crystalline framework in Cu‐L_N2S2 is built solely by intermolecular Cl ^… H interactions (Figure 1h). All the hydrogen bonding parameters are tabulated in Table S2−S3.

FIGURE 1.

(a,g) ORTEP drawing of Cu‐L_N3S and Cu‐L_N2S2 with 50% ellipsoidal probability of thermal ellipsoids. Hydrogen atoms are omitted for clarity in figure (a,b,g,h). (b,h) Supramolecular framework of Cu‐L_N3S and Cu‐L_N2S2 involving O^…H, Cl^…H and O^…π interactions. (c,i) Comparison of PXRD plot of Cu‐L_N3S simulated (black) and Cu‐L_N2S2 experimental (red). (d,j) FESEM image of Cu‐L_N3S and Cu‐L_N2S2 revealing the uniform growth of the complexes. (e,k) Elemental mapping using EDX spectroscopy of Cu‐L_N3S and Cu‐L_N2S2. (f,l) Elemental mapping of individual elements present in Cu‐L_N3S and Cu‐L_N2S2 Cu, O, N, S, Cl, and C.

Cu‐L_N3S and Cu‐L_N2S2 were characterized through powder X‐ray diffraction (PXRD) analysis. The PXRD patterns of the two Cu complexes exhibited sharp and intense peaks across varying 2θ values, indicating their prominent crystalline nature. Moreover, the close agreement between the experimental and simulated PXRD patterns confirms the high phase purity of the bulk materials (Figure 1c,i). Further, the coordination of the ligands L_N2S2 and L_N3S to the metal center was justified by the IR spectral analysis (Figures S8 and S9). The characteristic peaks of (‐Cu–N‐) stretching were observed at 764 and 693 cm⁻¹, indicating the formation of Cu‐L_N3S and Cu‐L_N2S2, respectively [54]. Other IR peaks for Cu‐L_N3S were observed at 1118, 1445, and 1637 cm⁻¹ due to stretching vibration. Similarly, Cu‐L_N2S2 showed IR peaks at 1635, 2353 cm⁻¹.

Nonetheless, the solution phase characteristics of the ligands (L_N2S2 and L_N3S) and Cu‐L_N2S2 and Cu‐L_N3S were studied through the Ultraviolet‐Visible (UV‐Vis) spectroscopy (Figure S10) in water medium. The ligands L_N2S2 and L_N3S exhibited high intensity electronic bands at 259 and 260 nm, respectively, assignable to the π→π*/n→π* electronic transitions of the pyridine moiety of ligand origin. Moreover, Cu‐L_N2S2 and Cu‐L_N3S render bands at 327 nm (151 M⁻¹ cm⁻¹) and 320 nm (82 M⁻¹cm⁻¹), respectively, along with the broad transition centered at 682 nm (5 M⁻¹ cm⁻¹) and 661 nm (3 M⁻¹ cm⁻¹), respectively, correlating with the d–d transition. These observed electronic bands align well with those reported for structurally analogous square‐pyramidal Cu(II) complexes [52, 53, 55]. TD‐DFT analysis reveals the emergence of an absorption band at 320 nm for Cu‐L_N3S, attributed to a ligand‐to‐metal charge transfer (LMCT) involving a β‐HOMO‐5 to β‐LUMO excitation (Figure S11a). In contrast, the corresponding electronic transition in Cu‐L_N2S2 appears at 327 nm and is characterized as a metal‐to‐ligand charge transfer (MLCT), also arising from a β ‐HOMO‐2 → β ‐LUMO transition with an oscillator strength of f = 0.06789 (Figure S11b).

Further, the molecular integrity of the Cu(II)‐complexes was deduced through electron spray ionization mass spectrometry (ESI‐MS) recorded in MeOH. ESI‐MS of the Cu‐L_N2S2 displays a molecular ion peak at m/z 374 with an isotopic distribution correlating to [Cu−L_N2S2]⁺ species (Figure S6). Conversely, ESI‐MS of Cu‐L_N3S reveals a molecular ion peak at m/z 357 with an isotopic distribution attributed to [Cu−L_N3S]⁺ species (Figure S5).

Moreover, the geometrical distortions of Cu‐L_N2S2 and Cu‐L_N3S were elucidated from X‐band EPR spectra recorded in methanol at room temperature. The EPR spectra of both Cu(II)‐complexes exhibit two distinct g‐values (g‖ and g⊥), each exceeding the free electron g‐value, as shown in Figure S12. The fact that g‖ > g⊥ > g_e indicates distorted square‐pyramidal geometry [56].

2.2. Water Phase Stability of Cu‐L_N2S2 and Cu‐L_N3S

Prior to evaluating their HER activity, the aqueous phase stability of the Cu(II) complexes was assessed via time‐dependent UV‐Vis spectrophotometric studies conducted over 72 h. For this purpose, the Cu(II) complexes were dissolved in water, and their UV‐Vis spectra were initially recorded. The same solutions were then left undisturbed for 72 h. Upon re‐recording the spectra after this period, the electronic absorption bands remained essentially unchanged, indicating excellent solution‐phase stability of the complexes (Figure S13).

However, in our earlier study involving a structurally analogous Cu(II) complex, we observed that the coordinated chloride ligand dissociated under electrochemical conditions and was subsequently replaced by a water molecule [52]. In light of this, we performed semimicro analysis to further investigate and validate the potential chloride ion dissociation phenomenon. For semimicro analysis, multiple CV cycles were conducted in 0.1 M KNO₃ electrolyte at a scan rate of 0.1 V/s within the potential window of 0 to –1.2 V (vs. Ag/AgCl) (Figure S14a,d). During successive CV sweeps followed by a few drops of AgNO₃, the formation of a white, curdy precipitate of AgCl was observed, which readily dissolved upon the addition of ammonium hydroxide and reappeared upon treatment with dilute HNO₃. This behavior confirms the presence of free chloride ions in the aqueous medium, thereby supporting the chloride ligand dissociation under electrochemical conditions (Figure S14b,e).

Furthermore, the rapid coordination of water molecules to the metal center is convincingly demonstrated through a combination of experimental observations and detailed theoretical studies. Experimentally, when the complex is dissolved, the electronic absorption spectrum of the resulting bluish filtrate exhibits a clear blue shift, manifested by the appearance of a distinct d–d transition band at 628 nm (Figure S14c,f). This shift indicates a change in the ligand field surrounding the copper center and is characteristic of an altered coordination environment. Such a modification is fully consistent with the substitution of the chloride ligand by water molecules, which would generate a new set of ligand‐field interactions and consequently shift the d–d transition energy.

To complement these findings, molecular structure optimizations were carried out in both the gas phase and in solution, with water (ε = 78.3553) incorporated as the solvent using the IEFPCM solvation model. The optimized geometries obtained for each phase were then compared with the solid‐state structure derived from X‐ray crystallography. The computational results reveal a notable elongation of the Cu—Cl bond in the aqueous‐phase optimized structure relative to the experimentally determined solid‐state geometry (Tables S4 and S5) and gas phase. This elongation provides a clear indication that the Cu—Cl interaction weakens significantly in a polar solvent environment, reflecting the increased lability of the chloride ligand under aqueous conditions. Such weakening is entirely consistent with the thermodynamic favorability of water coordination, as a more labile chloride ligand is readily displaced when water molecules interact with the copper center.

Together, these lines of evidence present a coherent and detailed mechanistic picture. The spectral blue shift signals a reorganization of the ligand field caused by chloride substitution, while the computationally observed bond elongation confirms the reduced stability of the Cu—Cl bond in water. As a result, water molecules can readily bind to the copper center, inducing a restructuring of the coordination sphere. Collectively, the experimental and theoretical insights firmly establishes water coordination as an early and significant structural event upon dissolution.

Upon examining the aqueous‐phase compatibility and the molecular integrity of the complexes, the electrochemical behavior of the complexes was studied in 0.1 M aqueous KCl. Initially, CVs were recorded between 0 to −1.2 V (V vs. Ag/AgCl) at the scan rate of 0.1 V/s. The CVs of Cu‐L_N3S displayed the two reduction events at −0.25 and −0.90 V (V vs. Ag/AgCl) (Figure 2a). The first reduction is ascribed to the Cu(II)/Cu(I) redox transition, while the second arises from a ligand‐centered reduction involving the Cu^I(L_N3S)/Cu^I(L_N3S)⁻. This interpretation is well‐aligned with established literature reports [13]. Furthermore, the DPV analysis exhibits a strong correlation with the observed reduction potentials of the complex (Figure S15a). To better understand the nature of redox‐active species at highly cathodic potentials, cyclic voltammetry (CV) of the free ligand L_N3S was recorded in aqueous medium over the potential range of 0 to –1.6 V under conditions identical to those used for the Cu complex (Figure 2a). Notably, the free L_N3S did not exhibit any discernible reduction wave within this potential window, suggesting its limited electrochemical activity in the cathodic region. However, to confirm the ligand‐centric second reduction event observed at –0.90 V in the copper complex, a structurally analogous zinc (II) complex, Zn–L_N3S was synthesized and fully characterized by spectroscopic techniques, including UV‐Vis and IR spectroscopy (Figure S16). The generation of a catalytic peak at 1.3 V from DPV studies of Zn–L_N3S in 0.1(M) KCl aqueous solution corroborated the reduction wave (Figure S17), firmly establishing the ligand‐centric redox wave.

FIGURE 2.

(a) CVs of 1 mM copper complex Cu‐L_N3S, L_N3S and GCE in 0.1 M aqueous KCl at 0.1 V/s. (b) CVs of the complex Cu‐L_N3S at various scan rates. (c) Variation of peak current (i _p) vs. square root of the scan rate; 9 CVs of 1 mM Cu‐L_N3S in 0.1 M aqueous KCl at 0.1 V/s scan rate without acid addition. (d) CVs of 1 mM copper complex Cu‐L_N2S2, L_N2S2, and GCE in 0.1 M aqueous KCl at 0.1 V/s. (e) CVs of the complex Cu‐L_N2S2 at various scan rates. (f) variation of peak current (i _p) vs. square root of the scan rate; 9 CVs of 1 mM Cu‐L_N2S2 in 0.1 M aqueous KCl at 0.1 V/s scan rate without acid addition.

In contrast, Cu‐L_N2S2 exhibited the Cu(II)/Cu(I) reduction peak at a more cathodic potential of –0.45 V (V vs. Ag/AgCl, indicating that the Cu(II) center within the N₂S₂ coordination environment is more resistant to reduction compared to its N₃S counterpart (Figure 2d). An anodic shift of approximately 10 mV was also observed in the ligand‐centered reduction wave, appearing at –0.89 V (V vs. Ag/AgCl) for Cu‐L_N2S2 relative to Cu‐L_N3S. Moreover, DPV studies under identical electrochemical conditions consistently revealed two distinct reduction events, which correlate well with the CV cycle (Figure S15b). In contrast, the anodic peak at −0.08 V (V vs. Ag/AgCl) for Cu‐L_N3S and −0.02 V (V vs. Ag/AgCl) for Cu‐L_N2S2 is assigned to Cu(I)/Cu(II) redox couple (Figure S41) [57].

Furthermore, the diffusion coefficient (D₀) is a fundamental parameter in homogeneous catalysis, especially in electrochemical studies, and plays a crucial role in understanding the kinetics and transport behavior of catalytic species in solution. Typically, the diffusion coefficient signifies the extent of molecular movement through the solvent. In homogeneous catalysis, the catalytic species must diffuse to and from the electrode surface. A higher D₀ suggests faster transport, which can lead to higher catalytic turnover, especially when electron transfer is coupled with mass transport. D₀ was calculated using Equation (S2), between the peak current and the square root of scan rate over the range of 100–800 mV/s. The calculated D₀ values were 4.1475 × 10⁻⁶ cm² s⁻¹ for Cu‐L_N3S and 5.0706 × 10⁻⁶ cm² s⁻¹ for Cu‐L_N2S2. The significantly higher D₀ of the later complex indicates a more rapid diffusion of the electroactive species toward the GCE surface, reflecting enhanced mass transport efficiency in the electrochemical medium.

2.3. Homogenous Electrocatalytic HER Studies of Cu‐L_N2S2 and Cu‐L_N3S

Homogenous HER studies of both Cu‐L_N2S2 and Cu‐L_N3S complexes were carried out with a three‐electrode setup consisting of a glassy carbon electrode as the working electrode (GCE), a platinum wire counter electrode, and an Ag/AgCl reference electrode in water using 0.1 M KCl as the supporting electrolyte in the presence of acetic acid (AcOH) proton source. Initially, CV was performed using only the electrolyte solution (0.1 M KCl), which showed no noticeable redox features. Separate addition of acetic acid, to the KCl solution in the absence of acetic acid/proton source failed to produce any progressive increment of the current (Figure S18). However, upon introducing 1 mM of Cu‐L_N3S and Cu‐L_N2S2 into acetic acid‐water media (Figure S18a,b), a significant rise in the catalytic currents for both complexes was observed. This clearly highlights the electrocatalytic role of the molecular complexes in facilitating proton reduction [16, 44, 58].

The enhancement in catalytic current upon acid addition is a common feature of molecular HER systems and is primarily attributed to increased proton availability, which accelerates the rate‐determining proton‐coupled electron transfer (PCET) steps. In many molecular catalysts, protonation of a reduced metal center or metal–ligand intermediate governs the HER kinetics. Thus, the introduction of acid facilitates faster formation of the key metal–hydride (M–H) intermediate, leading to higher catalytic turnover and increased current density. Moreover, weak acids can stabilize critical intermediates such as M–H species or protonated ligand states, thereby lowering the thermodynamic and kinetic barriers for hydrogen evolution. Increased proton concentration can also shift the catalytic pathway from a slower electron‐limited sequence (EEC) to more efficient proton‐responsive mechanisms (ECE or direct hydride‐proton coupling), resulting in a more pronounced catalytic wave. In addition, acid addition enhances local proton concentration near the electrode surface, improving proton transport and minimizing diffusion limitations. For Cu‐based complexes, protonation can further induce favorable electronic and structural changes, including enhanced ligand‐assisted proton shuttling and modulation of metal electron density, which collectively promote the formation of catalytically active Cu–H intermediates. Together, these effects account for the observed acid‐induced increase in HER current.

In Cu‐L_N3S, the emergence of a new catalytic wave at –0.95 V marked the onset of proton reduction, which progressively intensified (Figure 3a), clearly indicating the onset of electrocatalytic activity in aqueous media. The shifting of the ligand centric‐wave (−0.90 V) towards more negative potentials, reaching –1.2 V (vs. Ag/AgCl) as the AcOH concentration increased from 4 to 92 mM, confirmed the electrocatalytic activity of Cu‐L_N3S. However, beyond 96 mM AcOH, no further increase in catalytic current was observed, implying that all catalytic sites were saturated and excess protons could no longer be utilized [59, 60]. In contrast, Cu‐L_N2S2 exhibited the emergence of a catalytic wave at –0.89 V (V vs. Ag/AgCl), and upon successive addition of AcOH from 8 to 144 mM, the catalytic wave for proton reduction shifted to more negative potentials (Figure 3d), towards –1.4 V (V vs. Ag/AgCl). This pronounced cathodic shift, in comparison to Cu‐L_N3S, suggests that in the case of Cu‐L_N2S2, the electron transfer step plays a more prominent role as the rate‐determining step in the catalytic proton reduction process.

FIGURE 3.

(a) CVs of 1 mM complex Cu‐L_N3S in 0.1 M aqueous KCl at 1 V/s scan rate without acid addition (red) and upon the addition of different concentrations of AcOH. (b) i _C vs. [AcOH] fitted with a linear correlation exhibiting an R ² value of 0.998. (c) i _Cat/i _p vs. equivalent of AcOH fitted with a linear correlation. (d) CVs of 1 mM complex Cu‐L_N2S2 in 0.1 M aqueous KCl at 1 V/s scan rate without acid addition (red) and upon the addition of different concentrations of AcOH. (e) i _C vs. [AcOH] fitted with a linear correlation exhibiting an R ² value of 0.999. (f) i _Cat/i _p vs. equivalent of AcOH fitted with a linear correlation.

Another key electrochemical parameter used to assess catalytic performance is the i _cat/i _p value. Remarkably high values were observed for both complexes, with Cu‐L_N2S2 exhibiting an i _cat/i _p value of 122.1 (Table S6) at 144 mM AcOH and Cu‐L_N3S showing an i _cat/i _p value of 51.34 (Table S7) at 92 mM AcOH. These correspond to impressive turnover frequencies (TOFs) of 2.90*10³ s⁻¹ and 5.11*10² s⁻¹, for Cu‐L_N2S2 and Cu‐L_N3S, respectively. The detailed equations used for these calculations are provided in the Supplementary Information and summarized in Table S9.

Moreover, the i _cat/i _p ratio increases linearly with rising acid concentration (Figure 3c,f), suggesting a first‐order dependence [44, 61] of the catalytic rate on AcOH concentration (Figure S19). To probe the dependence on electrocatalyst concentration, AcOH concentration was fixed at 96 mM for Cu‐L_N3S and 144 mM for Cu‐L_N2S2, while the electrocatalyst concentration was systematically varied from 1.0 to 2.5 mM (Figure S20a,c). The resulting catalytic current i _cat/i _p exhibited a linear correlation with electrocatalyst concentration (Figure S20b,d), indicating a first‐order dependence on catalyst concentration as well.

To further judge the potential of the electrocatalysts for HER, overpotential (η) was deduced. Overpotential is a crucial metric in electrocatalysis that reflects how much additional energy (beyond the thermodynamic requirement) must be applied to drive HER. The overpotentials of Cu‐L_N2S2 and Cu‐L_N3S for HER were determined to be –0.657 and –0.882 V (V vs. Ag/AgCl, respectively, highlighting their commendable catalytic performance in comparison to other reported homogeneous molecular electrocatalysts (Table S8).

Building on these promising performance metrics, and in pursuit of greener and safer HER conditions, the alternative, environmentally benign electrolyte environment was also scrutinized. While KCl is commonly employed as an electrolyte in HER studies due to its widespread availability and established use in the literature, its corrosive nature and potential for chlorine gas evolution pose notable sustainability and safety challenges. To evaluate more environmentally benign alternatives, the electrocatalytic efficiencies of both electrocatalysts in the presence of acetic acid were examined in 0.1 M K₂SO₄ under similar electrochemical conditions (Figure S21). The scope for this alternative electrolyte offered good catalytic current responses, indicating that the catalytic performance is largely independent of the electrolyte medium. Thus, the HER in water for these newly designed bioinspired complexes offers safer and greener operational conditions without compromising efficiency.

Additionally, we have also prepared and isolated the water coordinated copper complex. The details of preparation, characterization, and analysis are summarized in SI. Although repeated attempts were made to isolate single crystals suitable for X‐ray diffraction, crystals of sufficient quality could not be obtained. Nevertheless, the successful formation of the desired aqua‐coordinated species was convincingly confirmed through complementary spectroscopic techniques, including UV–visible absorption spectroscopy (Figure S40a) and mass spectrometry (Figure S40b), which together validate the proposed molecular composition and coordination environment. Cu‐L_N3S‐H₂O exhibited high intensity electronic band at 298 nm (€ = 678 M⁻¹ cm⁻¹) which can be assigned as π to π*/ n to π* electronic transitions of the pyridine moiety of ligand origin. On the other hand, band at 628 nm (€=8 M⁻¹ cm⁻¹) aligns well with d–d transition. ESI‐MS shows the presence of molecular ion peak for [Cu‐L_N3S‐H₂O]²⁺ at m/z 340.0512 with an isotropic distribution that correlated with [Cu‐L_N3S‐H₂O]²⁺ species (Figure S40). In this context, the use of chloride‐coordinated copper complexes emerges as a pragmatic and scientifically justified choice. As discussed above, chloride ligands impart enhanced structural integrity and synthetic control, facilitating the formation of discrete, well‐defined coordination environments that are amenable to isolation as crystalline solids. This intrinsic stability not only simplifies handling and storage but also enables comprehensive and unambiguous characterization, particularly through crystallographic techniques. In contrast, the high lability of aqua‐coordinated species often leads to poorly crystalline or amorphous products, limiting detailed structural insight. Therefore, employing chloride‐coordinated copper complexes provides a reliable and experimentally tractable platform, while still allowing access to the corresponding aqua species through controlled ligand exchange, thereby effectively bridging synthetic feasibility with mechanistic and functional relevance, providing valuable mechanistic insights into catalyst activation under aqueous HER conditions.

2.4. Control Experiments

Understanding the mechanistic nature of electrocatalysis is crucial for the rational design of efficient and durable molecular catalysts. One of the key challenges lies in discerning whether the observed catalytic activity arises from truly homogeneous species in solution or from electrode‐adsorbed intermediates that can subtly alter the electrode surface. To interrogate the homogeneity and adsorption characteristics of isostructural Cu(II) complexes, Cu‐L_N3S and Cu‐L_N2S2, a series of meticulously designed electrochemical control experiments were performed.

To confirm whether the catalytic activity arose from freely diffusing molecular species or from electrode‐adsorbed intermediates, rinse tests were conducted. The glassy carbon electrode (GCE) was first cycled in the catalyst‐containing acid‐saturated electrolyte and then thoroughly rinsed with deionized water and transferred into a fresh solution of saturated acetic acid devoid of any catalyst .The absence of catalytic features in the resulting voltammograms ruled out any significant adsorption or deposition of active species on the electrode surface, thus confirming that the proton reduction proceeds via a homogeneous catalytic pathway under the applied conditions.

To reinforce this conclusion, soak tests were carried out in accordance with the method reported by Dempsey and coworkers [62]. In this experiment, the GCE was immersed overnight in a catalyst‐containing electrolyte‐acid solution and then rinsed and tested in a fresh electrolyte devoid of catalyst. The electrochemical response was then compared with that of a freshly polished electrode. In both cases, no detectable redox activity was observed, confirming the absence of any persistent, redox‐active surface species (Figure S18a,b).

Collectively, these control studies provide strong evidence that the catalytic activity arises exclusively from molecularly dissolved Cu(II) complexes, operating through a homogeneous pathway. Furthermore, these highlight the excellent stability and nonadsorptive behavior of these complexes in acidic electrochemical environments, underscoring their potential as reliable and efficient molecular electrocatalysts.

2.5. Molecular Stability and Durability Assessment

The long‐term operational stability and structural durability of molecular electrocatalysts are key parameters in determining their viability for practical applications in hydrogen evolution. Control potential electrolysis (CPE), conducted under sustained catalytic conditions, serves as a robust tool to evaluate not only the electrocatalytic performance but also the molecular integrity of catalysts over time.

To assess these factors, CPE was carried out for two Cu(II)‐based complexes, Cu‐L_N3S and Cu‐L_N2S2, at a fixed potential of –1.2 V vs Ag/AgCl for 2 h in acid‐saturated aqueous electrolyte. Both Cu‐L_N3S and Cu‐L_N2S2 displayed a steady and linear current profile, with characteristic serrated fluctuations caused by the periodic evolution and release of hydrogen bubbles, indicative of continuous and stable catalysis (Figure S22a,c).

Quantitative analysis revealed total charges of 25.82 C and 29.26 C (Figure S22b,d) for Cu‐L_N3S and Cu‐L_N2S2, respectively, corresponding to hydrogen evolution of 13.0 × 10⁻⁵ and 14.8 × 10⁻⁵ moles, respectively. Gas chromatography with a thermal conductivity detector (GC‐TCD) confirmed molecular hydrogen as the sole gaseous product, with Faradaic efficiencies (FE) of 96% for Cu‐L_N3S and 88% for Cu‐L_N2S2.The details of the calculation are provided in the SI.

The slightly lower FE observed for the copper‐based catalysts arises from minor non‐Faradaic processes occurring during prolonged electrocatalysis. Under reductive HER conditions, the coordinated organic ligands are not completely electrochemically inert and may undergo partial defragmentation or degradation, leading to side reactions that consume a small fraction of the applied current without contributing to hydrogen evolution. Such processes can be accompanied by surface reconstruction of the copper centers into Cu⁰/Cu nanoparticles or copper acetate species in the presence of AcOH. Minor contributions from capacitive charging, surface redox processes, and transient adsorption–desorption phenomena may also contribute to the non‐Faradaic current. Experimental evidence supports this interpretation. After controlled potential electrolysis, visible deposits formed on the electrode surface, which dissolved in dichloromethane to give a pale solution. UV–Vis analysis of this solution showed absorption features consistent with fragmented ligand species (2‐picolyl chloride), providing direct evidence of partial ligand defragmentation during catalysis (Figure S42). Furthermore, UV–Vis spectra of the post‐electrolysis solutions showed a reduced intensity of the characteristic d–d transition compared to fresh catalyst solutions, indicating partial structural transformation of the molecular copper complexes (Figure S43). Despite these secondary processes, the FE values remain high, confirming that hydrogen evolution is the dominant electrochemical pathway. The modest deviation from 100% FE is therefore attributed to catalyst activation and structural evolution under operating conditions rather than any loss of catalytic selectivity.

2.6. Proposed Homogenous HER Mechanism

A comprehensive understanding of the mechanistic pathway is essential to elucidate the nature of catalytic hydrogen evolution and to guide the rational design of improved molecular electrocatalysts. In this section, we propose a detailed homogeneous mechanism for the HER catalyzed by the Cu‐L_N3S and Cu‐L_N2S2 complexes, based on an integrated approach combining experimental and computational techniques. Ultraviolet‐visible (UV‐Vis) spectroscopy, electron paramagnetic resonance (EPR), gas chromatography (GC), and density functional theory (DFT) calculations were employed to unravel the electronic structure, protonation behavior, and stepwise reaction intermediates of the catalyst under reductive conditions.

2.7. Initiation: Aquation of the Cu(II) Complexes

Preliminary electrochemical studies of both Cu(II) complexes revealed that the catalytic cycle begins with the dissociation of the coordinated chloride ligand, which is subsequently replaced by a water molecule (see section: Solution Phase Behavior and Figure S23). This indicates that both electrocatalysts initiate their first cycle through a common mechanistic step wherein the water‐coordinated species serve as the active catalysts for both Cu complexes.

2.7.1. Step I: Sequence of Protonation and Reduction

Furthermore, to determine the sequence of events, whether protonation occurs before reduction or vice versa, UV‐Vis titration was employed to gain deeper mechanistic insight.

2.7.2. Cu‐L_N3S: Reduction Before Protonation

The titration spectra of Cu‐L_N3S displayed consistent absorbance features, closely resembling those of the pristine complex, even upon successive additions of AcOH. Notably, no isosbestic point was observed throughout the titration (Figure S24a), eliminating the possibility of protonation, thereby suggesting reduction occurs first, followed by protonation in the next step. This reduction was further validated by EPR studies conducted in the presence of cobaltocene, a standard reductant (Figure S25a). The EPR‐silent response of Cu‐L_N3S upon cobaltocene addition is consistent with the formation of a Cu(I) species. This interpretation is corroborated by the spin‐density distribution plot (Figure S26), which shows that the unpaired electron density is primarily localized on the copper center rather than the ligand framework. Collectively, these results confirm that the complex undergoes a copper‐centered reduction as the initial activation event prior to protonation.

2.7.3. Cu‐L_N2S2: Protonation Before Reduction

In the case of Cu‐L_N2S2, a noticeable isosbestic point was observed at 302 nm (Figure S24b), suggesting the possibility of dynamic equilibrium between protonated and unprotonated species of the catalyst and signifying the protonation at Cu‐L_N2S2 in the primary stage. The peak observed at 327 nm upon acid addition suggests that the transition is hindered in the presence of acetic acid. The HOMO is predominantly localized on the two nitrogen atoms and the copper center, indicating that protonation preferentially occurs at the nitrogen sites rather than at the sulfur atom, which is consistent with the Mulliken charge values (N =  −0.641, −0.640; S = 0.090, 0.072). Mulliken charges is obtained from the Mulliken population analysis for the nitrogen and sulfur donor atoms in the optimized structures. They describe the distribution of electron density on these coordinating sites and help illustrate their relative electron‐donating abilities toward the metal center. Furthermore, the two Cu–N bond lengths are nearly identical, measuring 1.990 and 1.988 Å, respectively. Nonetheless, the protonation site in the Cu‐L_N2S2 was convincingly elucidated through pH‐dependent CV measurements in citrate‐phosphate buffer across a pH range of 3.2–6.4 (Figure S27a). A Pourbaix diagram revealed a distinct vertical shift in the redox potential with a slope of approximately 23 mV per pH unit for Cu‐L_N2S2, indicative of a proton‐coupled electron transfer process. This transition was centered around pH 5.2 for (Figure S27b), indexing well with the reported pK_a value of pyridinium ions in aqueous media (∼5.2) [63]. Such agreement strongly supports the hypothesis that protonation in the Cu(II) complex occurs specifically at the pyridine nitrogen atom. Further corroborating this assignment, literature precedents report significantly higher pKa values for aqua ligands coordinated to Cu(II) centers, typically falling within the range of 8–10 [64, 65, 66]. This substantial disparity in acidity effectively rules out coordinated water as the protonation site under the experimental conditions. Taken together, these findings validate that the pyridine nitrogen, owing to its congruent pKa and its strategic positioning within the ligand environment, is the most plausible site of proton uptake. This conclusion not only reinforces the proton‐responsive nature of the pyridine moiety but also underscores its critical role in modulating the redox behavior of the Cu(II) complex through proton‐coupled electron transfer mechanisms.

2.7.4. Step II: Protonation and Reduction of the Activated Complex

In the second step, for the Cu‐L_N3S, protonation may occur at four plausible donor sites: the thioether sulfur (S), the amine nitrogen (N), and the two pyridine‐arm nitrogen atoms (N‐py‐N and S‐py‐N). Mulliken charge analysis indicates that the pyridine‐arm nitrogen atoms (N‐py‐N and S‐py‐N) are the most favorable protonation sites, exhibiting significant negative partial charges of –0.587 and –0.589, respectively (Figure S28). Between the two pydidine‐Ns, aliphatic‐N‐adjacent py‐N is slightly more favorable than aliphatic‐S adjacent py‐N, as suggested by a marginally longer Cu–N bond distance of 1.97 Å compared to S‐py‐N (1.95 Å) (Figure S29a). Furthermore, the optimized structure of the N‐py‐N protonated species is energetically more stable by 5.72 kcal/mol, reinforcing its preference as the favored protonation site (Figure S30). The protonation site in the Cu‐L_N2S2 complexes was clearly identified through pH‐dependent CV studies (pH 3.2–6.4), which showed a shift in redox potential (∼49 mV/pH) indicative of proton‐coupled electron transfer centered at pH 5.2 (Figure S31), closely matching the pKa of pyridinium (∼5.2) [63]. On the other hand, the second step for Cu‐L_N2S2 involves reduction, as supported by theoretical calculations. The reduction was verified by recording the EPR spectrum in the presence of cobaltocene, which acted as a standard reducing agent (Figure S25b). In the monoprotonated Cu‐L_N2S2, the spin density is primarily localized on the copper center rather than the ligand framework (Figure S32). Likewise, the SOMO is centered on the copper atom, confirming that the reduction takes place at the copper site and originates from a singlet Cu(I) species (Figure S33).

2.7.5. Step III: Pyridine‐Arm Decoordination and Hydride Formation

The protonation event of Cu‐L_N3S is subsequently followed by (N)pyridine arm decoordination, accompanied by a reorganization of the metal's coordination environment. This mechanistic insight is further supported by the anodic shift observed in the Cu(II)/Cu(I) redox couple upon acid addition. Notably, this behavior closely parallels the findings of Smolentsev et al., who reported a similar pathway for the pentapyridyl cobalt complex [Co(aPPy)] under acidic conditions, wherein pyridine ligand decoordination precedes and facilitates protonation during HER [14]. Further insights from computational analysis of Cu‐L_N3S reveal that in the N‐pyridine‐protonated Cu(I) intermediate (D), the LUMO is primarily localized on the pyridine ring adjacent to the protonated nitrogen (Figure S34). This suggests that the electron‐deficient pyridine ring acts as the preferred site for accepting a second electron (E). The resulting increase in electron density promotes an intramolecular hydrogen transfer to the Cu(I) center, leading to the formation of a Cu(II)–hydride species (F). Following this step, the pyridine arm adjacent to the thioether sulfur undergoes protonation and subsequent decoordination of water, forming a new intermediate (G). This is supported by structural data showing Cu—N bond lengths of 1.92 Å for the N_py–S site and 1.88 Å for the N_py–N site (Figure S29b), indicating a weaker interaction and easier release of the N_py–S donor. In Cu‐L_N3S forms a Cu(II)–hydride species, with a Cu—H···N—H distance of 1.5 Å (Figure S35a). On the other hand, Cu(I) in Cu‐L_N2S2 forms a Cu(II)–hydride species, with a Cu—H···N—H distance of 1.65 Å. (Figure S35b).All the intermediated optimized structures are shown in the SI (Figure S36, S37).

Nonetheless, the formation of a Cu(II)–hydride intermediate during electrocatalytic HER by the Cu(II) complexes was validated through the hydrogenation reaction of styrene to ethylbenzene, following the reported methodology by Navarro [67]. Detailed experimental procedures are provided in the supplementary information (SI). The product was analyzed by gas chromatography–mass spectrometry (GC–MS), which revealed a prominent molecular ion peak at m/z 106, corresponding to 2‐phenylethan‐1‐ylium (C₆H₅CH₂CH₃ ⁺), thereby confirming the successful formation of ethylbenzene (Figures S38 and 39). This result strongly supports the presence of a copper‐hydride intermediate, as the hydrogenation of styrene necessitates a hydridic hydrogen source. Thus, the Cu—H species generated during catalysis is not only electronically feasible but also chemically competent in mediating hydrogen atom transfer, reinforcing its role as a key intermediate in the HER pathway.

2.7.6. Step IV: Hydrogen Evolution and Catalyst Regeneration

In both Cu(II) complexes, the catalytic sequence concludes with the formation of a Cu(II)–hydride species, a pivotal intermediate that facilitates the coupling of two adjacent hydrogen atoms. This concerted step results in the evolution of molecular hydrogen (H₂), marking the successful completion of the hydrogen evolution reaction. Importantly, this transformation simultaneously regenerates the active Cu(II) catalyst, thereby reinitiating the catalytic cycle (Figure 4).

FIGURE 4.

Proposed mechanistic cycle for CuL_N3S and CuL_N2S2.

3. Conclusion

This study presents a comprehensive investigation into the design, synthesis, and mechanistic elucidation of two structurally distinct copper(II) complexes, Cu‐L_N3S and Cu‐L_N2S2, bearing N,S‐donor ligands, for electrocatalytic hydrogen evolution in aqueous media. By combining rigorous experimental characterization with detailed electrochemical and theoretical studies, we establish the first comparative structure–activity relationship between these complexes and delineate their divergent catalytic pathways. Our findings underscore the pivotal role of ligand architecture in modulating the redox behavior, stability, and catalytic efficiency of copper‐based HER electrocatalysts. The Cu‐L_N3S complex, featuring an N₃S coordination sphere, exhibited superior diffusion kinetics, lower overpotential (–0.657 V), and greater long‐term catalytic stability, with a turnover frequency of 2.90*10³ s⁻¹ and a faradaic efficiency of 96%. In contrast, the Cu‐L_N2S2 complex, although demonstrating a higher catalytic current and turnover frequency (5.11*10² s⁻¹), showed signs of minimal structural degradation over prolonged electrolysis, reflected in a reduced faradaic efficiency of 88%. CPE of Cu‐L_N2S2 for 3h confirmed the direct evidence for partial ligand fragmentation, exhibiting a lower FE during the extended electrolysis. Mechanistically, both catalysts initiate their HER cycles via chloride dissociation and subsequent water coordination. However, their pathways diverge significantly. In Cu‐L_N3S, reduction precedes protonation, progressing through a CECE‐type mechanism, whereas Cu‐L_N2S2 likely follows an ECEC route. Computational studies further highlight the reactivity imparted by N,S‐type donors, particularly the preferential protonation at pyridine‐arm nitrogen atoms and the crucial role of intramolecular hydrogen transfer leading to the formation of Cu–hydride intermediates. Altogether, this work not only showcases the untapped potential of copper complexes with N,S‐donor ligands for sustainable hydrogen production but also offers valuable mechanistic insights that will inform the rational design of next‐generation molecular electrocatalysts.

It is indeed well established that homogeneous HER catalysts operating in aqueous electrolytes, such as KCl, typically exhibit higher overpotentials and lower current densities than state‐of‐the‐art heterogeneous electrocatalysts. However, the primary objective of homogeneous systems is not immediate industrial deployment, but rather the generation of mechanistic clarity under well‐defined molecular environments.

In the present study, the homogeneous catalyst operates in aqueous KCl, a neutral, noncorrosive, and industrially benign electrolyte that avoids the use of strongly acidic or alkaline media. This choice is particularly relevant for future decentralized and modular hydrogen‐generation technologies, where compatibility with mild electrolytes, corrosion resistance, and operational safety are critical considerations. The molecular nature of the catalyst enables precise control over the coordination sphere, ligand electronics, and proton‐coupled electron transfer pathways, allowing direct correlation between molecular structure and HER activity, an understanding that is largely inaccessible in conventional heterogeneous systems.

From a technological perspective, homogeneous HER catalysts serve as ideal prototypes for identifying key catalytic motifs, proton relay mechanisms, and metal‐hydride intermediates that govern hydrogen evolution kinetics. The insights obtained from such systems can be directly translated into the rational design of heterogeneous or hybrid catalysts, including MOF‐derived, polymer‐immobilized, or electrode‐anchored molecular architectures that are more amenable to industrial implementation. In this regard, the present work provides a valuable molecular blueprint for catalyst optimization beyond the homogeneous phase. With respect to economic and sustainability considerations, the reported catalyst is based on earth‐abundant, non‐noble metals and is synthesized using solution‐based, low‐energy methods without precious metals or hazardous reagents. The use of aqueous KCl further minimizes environmental footprint, reduces electrolyte cost, and aligns with green‐chemistry and sustainability principles. Moreover, homogeneous systems enable complete utilization of metal centers without the mass‐transport limitations often encountered in bulk heterogeneous materials, offering high intrinsic metal efficiency.

Although homogeneous HER catalysts in aqueous media are currently limited by challenges such as catalyst recovery, long‐term stability, and scalability, ongoing advances in catalyst immobilization, membrane separation, and flow‐electrolyzer design are rapidly narrowing this gap. Therefore, while the present catalyst is not intended as a direct industrial competitor, it establishes fundamental design principles and mechanistic insights that are essential for the development of cost‐effective, durable, and scalable hydrogen‐evolution technologies. Overall, this work positions homogeneous molecular catalysis in aqueous KCl as a critical stepping stone toward next‐generation electrocatalysts, bridging fundamental molecular electrochemistry and future industrial hydrogen production.

Supporting Information

Additional supporting information can be found online in the Supporting Information section. Supporting Fig. S1: ¹H NMR spectrum (400 MHz, CDCl₃, 298 K) of L_N3S in CDCl₃. Supporting Fig. S2: ¹³C NMR spectrum (400 MHz, CDCl₃, 298 K) of L_N3S in CDCl₃. Supporting Fig. S3: ¹H NMR spectrum (400 MHz, CDCl₃, 298 K) of L_N2S2 in CDCl₃. Supporting Fig. S4: ¹³C NMR spectrum (400 MHz, CDCl₃, 298 K) of L_N2S2 in CDCl₃. Supporting Fig. S5: ESI‐MS of Cu‐L _N3S in methanol: Calculated (m/z) for [Cl‐Cu‐L_N3S]⁺ = 357.0127, Found = 357.0121. Supporting Fig. S6: (1) ESI‐MS of Cu‐L _N2S2 in methanol: Calculated (m/z) for [Cl‐Cu‐L_N2S2]⁺ = 373.97, Found = 374. Isotropic distribution of m/z value of (a) [Cl‐Cu‐L_N2S2]⁺ and (b) [Cl‐Cu‐L_N3S]⁺ (black) along with their theoretical values (red). Supporting Fig. S7: Representation of dihedral angle between two planes (a) Cu1−S1‐C6 and Cu1‐N2A‐C9{65.68°} of Cu‐L _N3S (b) Cu1−S1‐C6 and Cu1‐S2‐C9{87.93°} of Cu‐L _N2S2 . Supporting Fig. S8: FT‐IR spectra of 10⁻³M (a) L_N3S (b) L_N2S2. Supporting Fig. S9: IR spectra of (a) Cu‐L _N3S (b) Cu‐L _N2S2 . Supporting Fig. S10: UV‐Vis spectra of (a) 10⁻⁵M L_N3S and Cu‐L _N3S (b) 10⁻⁵ M L_N2S2 and Cu‐L _N2S2 in water. Supporting Fig. S11: The molecular orbitals assignable for the (a) LMCT band in the [Cu‐L _N3S ], and the (b) MLCT band in the [Cu‐L _N2S2 ], computed using B3LYP/6‐311g(d, p) in water. Supporting Fig. S12: X‐band EPR spectrum of (a) Cu‐L_N3S (experimental vs. simulated) and (b) Cu‐L_N2S2 (experimental vs. simulated) in methanol at room temperature. Supporting Fig. S13: UV‐Vis spectra of (a) Cu‐L _N3S (b) Cu‐L _N2S2 recorded in water for 3 consecutive days. Supporting Fig. S14: a) CV cycle of Cu‐L _N3S in 0.1 M KNO₃ at a scan rate of 100 mV/s in the range of 0 to −1.2V (V vs. Ag/AgCl) b) UV‐Vis spectrum focusing d‐d band of Cu‐L_N3S in aqueous 0.1 M KNO₃ (black) and after performing chloride dissociation test followed multiple CVs cycle (blue) c) CV cycle of Cu‐L_N2S2 in 0.1 M KNO₃ at a scan rate of 100 mV/s in the range of 0 to −1.4V (V vs. Ag/AgCl) d) UV‐Vis spectrum focusing d‐d band of [Cu‐L _N2S2 ] in aqueous 0.1 M KNO₃ (black) and after performing chloride dissociation test followed multiple CVs cycle (blue). Supporting Fig. S15: Differential pulse voltammetry (DPV) plot of (a) Cu‐L _N3S (b) Cu‐L _N2S2 in 0.1 M aqueous KCl at 100 mV/s. Supporting Fig. S16: UV‐Vis and IR spectra of (a), (c) 10⁻⁵ M Zn‐L_N3S and (b), (d) 10⁻⁵ M Zn‐L_N2S2 in water. Supporting Fig. S17: (a) Differential pulse voltammetry (DPV) plot of Zn‐L_N3S in 0.1 M aqueous KCl at 100 mV/s (b) DPV plot of of Zn‐L_N3S upon addition of AcOH upto 44 mM. Supporting Fig. S18: Post‐CV dip test of (a) Cu−L _N3S and (b) Cu‐L _N2S2 for GCE in fresh 0.1 M aqueous KCl. Supporting Fig. S19: a) Plot of [H⁺] vs. i_cat and b) double‐logarithmic plot of i_cat vs. [H⁺] of Cu‐L _N3S corresponding to typical electrocatalytic AcOH titration from 4 mM to 96 mM c) Plot of [H⁺] vs. i_cat and d) double‐logarithmic plot of i_cat vs. [H⁺] of Cu‐L _N2S2 corresponding to typical electrocatalytic AcOH titration from 4 mM to 144 mM. Supporting Fig. S20: a) CVs of different concentration of Cu‐L _N3S , 1.00 mM (Red), 1.5 mM (green), 2.00 mM (black) and 2.5 mM (blue) in 0.1 M aqueous KCl containing 96 mM AcOH at a scan rate of 100 mV/s in the region of 0 to −1.3 V (V vs. Ag/AgCl) b) peak current (i_c) vs. Cu‐L _N3S fitted with a linear correlation exhibiting an R² value of 0.99 c) CVs of different concentration of Cu‐L_N2S2, 1.00 mM (Red), 1.5 mM (green), 2.00 mM (black) and 2.5 mM (blue) in 0.1 M aqueous KCl containing 144 mM AcOH at a scan rate of 100 mV/s in the region of 0 to −1.6 V (V vs. Ag/AgCl) d) peak current (i_c) vs. Cu‐L_N2S2 fitted with a linear correlation exhibiting an R² value of 0.99. Supporting Fig. S21: DPV plots of a) 1 mM Cu−L _N3S in 0.1 M K₂SO₄ electrolyte (b) 1 mM Cu−L _N2S2 in 0.1 M K₂SO₄ electrolyte c) CVs of 1 mM Cu−L _N3S using K₂SO₄ as electrolyte with subsequent AcOH addition, (d) CVs of 1 mM Cu−L _N2S2 using K₂SO₄ as electrolyte with subsequent AcOH addition. Suppporting Figure 22: a) CPE of 1 mM Cu−L _N3S with 92 mM AcOH (blue) for 2h; b), d) The corresponding Charge vs. time plots c) CPE of Cu−L _N2S2 (purple) with 144 mM AcOH for 2h in 0.1 (M) KCl aqueous solution. Supporting Fig. S23: Energy profile for the intermediate (A) and (B) for Cu‐L _N3S by using DFT calculation. Supporting Fig. S24: Change in UV‐vis spectra with sequential addition of AcOH from 4 mM to 60 mM in a aqueous solution of a) Cu−L _N3S b) Cu−L _N2S2 . Supporting Fig. S25: The EPR‐silent nature of the a) Cu−L _N3S and b) Cu−L _N2S2 upon 1:1 reaction with a standard one‐electon reducing agent, cobaltocene, in DMF under N₂ atmosphere. Supporting Fig. S26: The spin density distribution plot of Cu‐L _N3S at an isovalue of 0.004. Supporting Fig. S27: a) DPV and b) The Pourbaix diagram of [Cu−L _N2S2 ] in the citrate‐phosphate buffer within the pH range 3.2 to 6.4. Supporting Fig. S28: Mulliken Charge distribution of the optimized structure for the intermediate C. Supporting Fig. S29: The optimized geometry of the intermediate (a) C and (b) F (Figure 4) showing the bond distances of N(Py) and S(Py) atom from the Cu centre in Cu‐L _N3S . Supporting Fig. S30: Energy profile for the two possibility of intermediate D; pyridine nitrogen adjacent to sulfur (N_py‐S), and the pyridine nitrogen adjacent to the amine group (N_py‐N) by using DFT calculation. Supporting Fig. S31: a) DPV and b) The Pourbaix diagram of [Cu−L _N3S ] in the citrate‐phosphate buffer within the pH range 3.2 to 6.4. Supporting Fig. S32: The spin density distribution plot of intermediate 3 for Cu‐L _N2S2 at an isovalue of 0.004. Supporting Fig. S33: The SOMO of the protonated intermediate 3 in Cu‐L _N2S2 . Supporting Fig. S34: LUMO of the N‐pyridine‐protonated Cu(I) intermediate D. Supporting Fig. S35: Computed optimized geometry of the doubly protonated reactive intermediate (a) G and (b) 5 showing the closest proximity for H⁺…H^‐ coupling. Supporting Fig. S36: The optimized geometries with various important aspects of Cu−L _N3S and reactive intermediates. Supporting Fig. S37: The optimized geometries with various important aspects of Cu−L _N2S2 and reactive intermediates. Supporting Fig. S38: Mass spectra for hydrogenation of styrene in presence of Cu‐L _N3S using GC–MS analysis. Supporting Fig. S39: Mass spectra for hydrogenation of styrene in presence of Cu‐LN _2S2 using GC–MS analysis. Supporting Fig. S40: a) UV‐Vis spectra of 10⁻⁵ M [H₂O‐Cu‐L_N3S]²⁺ in water. b) ESI‐MS of [H₂O‐Cu‐L_N3S]²⁺in methanol at room temperature (RT). Supporting Fig. S41: Complete CV cycle of 1 mM (a) Cu‐L_N3S and (b) Cu‐L_N2S2 using GCE in 0.1 M aqueous KCl at 0.1 V/S scan rate. Supporting Fig. S42. UV‐Vis spectra of deposited species of (a) Cu‐L_N3S and (b) Cu‐L_N2S2 in water comparing with bulk solution after chronoamperometry. Supporting Fig. S43: UV–Vis spectra of the post‐CPE solutions (black) were compared with those of freshly prepared (red) (a) Cu‐L_N3S and (b) Cu‐L_N2S2 solutions in 0.1 M aqueous KCl containing AcOH at identical concentrations. Supporting Table S1: Crystallographic structural parameters of [Cu−L_N3S]. Supporting Table S2: Hydrogen bond lengths (Å) and angles (°) for [Cu−L_N3S]. Supporting Table S3: Hydrogen bond lengths (Å) and angles (°) for [Cu−L_N2S2]. Supporting Table S4: Selected bond distances (Å) and angles (°) of [Cu−L_N3S]. Supporting Table S5: Selected bond distances (Å) and angles (°) of [Cu−L_N2S2] . Supporting Table S6: Electrochemical experimental results from a typical AcOH titration of [Cu‐L_N2S2] obtained using Equation S3. Supporting Table S7: Electrochemical experimental results from a typical AcOH titration of [Cu‐L_N3S] obtained using Equation S3. Supporting Table S8: Performance of the reported catalysts for hydrogen evolution studies. Supporting Table S9: A comparison of the electrocatalytic performances of [Cu‐L_N2S2] and [Cu‐L_N3S].

Funding

This study was supported by the Anusandhan National Research Foundation (ANRF/IRG/2024/000343/ CS).

Conflicts of Interest

The authors declare no conflicts of interest.

Supporting information

Supplementary Material

Data Availability Statement

The data that support the findings of this study are available from the corresponding author upon reasonable request.