# Proton Relays in Molecular Catalysis for Hydrogen Evolution and Oxidation: Lessons From the Mimicry of Hydrogenases and Electrochemical Kinetic Analyses

Matthieu Haake, Bertrand Reuillard, Murielle Chavarot-Kerlidou, Cyrille

Costentin, Vincent Artero

► To cite this version:

Matthieu Haake, Bertrand Reuillard, Murielle Chavarot-Kerlidou, Cyrille Costentin, Vincent Artero. Proton Relays in Molecular Catalysis for Hydrogen Evolution and Oxidation: Lessons From the Mimicry of Hydrogenases and Electrochemical Kinetic Analyses. Angewandte Chemie International Edition, 2024, 63 (52), pp.1-22. 10.1002/anie.202413910. hal-04805512

HAL Id: hal-04805512

https://hal.science/hal-04805512v1

Submitted on 19 Nov 2025

HAL is a multi-disciplinary open access
archive for the deposit and dissemination of sci-
entific research documents, whether they are pub-
lished or not. The documents may come from
teaching and research institutions in France or
abroad, or from public or private research centers.

L'archive ouverte pluridisciplinaire HAL, est
destinée au dépôt et à la diffusion de documents
scientifiques de niveau recherche, publiés ou non,
émanant des établissements d'enseignement et de
recherche français ou étrangers, des laboratoires
publics ou privés.

<figure><img src="image_1.png" /></figure>

Distributed under a Creative Commons Attribution 4.0 International License
Bioelectrocatalysis

# Proton Relays in Molecular Catalysis for Hydrogen Evolution and Oxidation: Lessons From the Mimicry of Hydrogenases and Electrochemical Kinetic Analyses

Matthieu Haake, Bertrand Reuillard, Murielle Chavarot-Kerlidou, Cyrille Costentin,* and Vincent Artero*

How to cite: Angew. Chem. Int. Ed. **2024**, 63, e202413910
doi.org/10.1002/anie.202413910

<figure><img src="image_2.png" /></figure>
**Abstract:** The active sites of metalloenzymes involved in small molecules activation often contain pendant bases that act as proton relay promoting proton-coupled electron-transfer processes. Here we focus on hydrogenases and on the reactions they catalyze, i. e. the hydrogen evolution and oxidation reactions. After a short description of these enzymes, we review some of the various biomimetic and bioinspired molecular systems that contain proton relays. We then provide the formal electrochemical framework required to decipher the key role of such proton relay to enhance catalysis in a single direction and discuss the few systems active for H₂ evolution for which quantitative kinetic data are available. We finally highlight key parameters required to reach bidirectional catalysis (both hydrogen evolution and hydrogen oxidation catalyzed) and then transition to reversible catalysis (both reactions catalyzed in a narrow potential range) as well as illustrate these features on few systems from the literature.

# 1. Introduction

Electrocatalysis is now established as a key technology for the energy transition as it underlies power-to-chemical/fuel (or chemical/fuel-to-power) and, indirectly, light-to-chemical/fuel processes[1] encompassing the production and use of hydrogen, a key enabler of a net-zero emission economy, but also the CO₂ valorization,[2] the still eluding N₂ fixation to produce ammonia[3] as well as the electrification of synthetic chemistry.[4] Yet, efficient, stable and cheap catalysts are still missing for most of these inherently kinetically demanding multiproton-multielectron processes. Interestingly, similar processes are also key for the metabolism of microorganisms, expressing specific metalloenzymes to drive them at high rate and often close to the thermodynamic equilibrium, which is an asset for energy conversion at high efficiency.[5] Since these enzymes have evolved under strong bioavailability constrain, they only contain abundant metal ions from the first transition metal series to which Mo and W should be added. Therefore, lessons learned from these enzymes should allow to derive both efficient and sustainable catalytic systems. In this review, we will focus on the hydrogen evolution (HER) and oxidation (HOR) reactions that are key for the production of H₂ in electrolyzers and its use in fuel cells where platinum currently stands as the reference catalyst.[6] Pt catalyzes the interconversion of protons and electrons into H₂ in water reversibly, meaning that catalysis operates in both directions at the thermodynamic equilibrium. This property has long been restricted to platinum and other electrode materials are only active for H₂ evolution. Interestingly, hydrogenase

enzymes[7] also catalyze this reaction reversibly in most cases after immobilization onto a carbon-based electrode.[8] This unique behavior is related to the peculiar structure of their active sites and in particular with the fact that they contain proton relays installed in the immediate vicinity of the metallic centers, leading to a very active area of research towards the mimicry of their activity. In this review, we will elaborate on few biomimetic and bioinspired catalysts to show how such proton relays effectively impact the efficiency of catalysis in one direction (or monodirectional catalysis), namely towards hydrogen evolution, measured via its overpotential requirement and turnover frequency—TOF, but also allow for bidirectional (both hydrogen evolution and hydrogen oxidation catalyzed) and/or reversible (both reactions catalyzed in a narrow potential range) operation (measured via metrics that will be defined later in the review).

# 2. Hydrogenases

Hydrogenase (H₂ase) enzymes are expressed in a variety of microorganisms (archaea, bacteria, green algae and microalgae...), where they partake the bioenergetics metabolism.[7] They belong to two classes characterized by the nature of their organometallic active sites, a diiron cluster for [FeFe]-H₂ases and a nickel-iron cluster for [NiFe]-H₂ases (Figure 1). Both catalyze the interconversion of protons and electrons with H₂ in both directions.

The development of protein film voltammetry[8a,b] allowed to gain more insights into the catalytic activity of the various hydrogenase enzymes. The diversity of electrochemical responses (Figure 2) not only depends on the nature of the active site but also reflects the influence of the polypeptidic envelope and of the other cofactors, namely iron-sulfur clusters that mediate electron transfer between

<figure><img src="image_3.png" /></figure>

<figure><img src="image_4.png" /><figcaption>**Figure 1.** Structure of the active sites of the [FeFe]-hydrogenases (left) and [NiFe]-hydrogenases (right) in their reduced state.</figcaption></figure>

<footnote>[*] M. Haake, Dr. B. Reuillard, Dr. M. Chavarot-Kerlidou, Dr. V. Artero
Univ. Grenoble. Alpes, CNRS, CEA, IRIG, Laboratoire de Chimie et
Biologie des Métaux
38000 Grenoble (France)
E-mail: vincent.artero@cea.fr

Prof. C. Costentin
Département de Chimie Moléculaire
Univ. Grenoble. Alpes, CNRS
38000 Grenoble (France)
E-mail: cyrille costentin@univ-grenoble-alpes.fr</footnote>

<footnote>© 2024 The Authors. Angewandte Chemie International Edition published by Wiley-VCH GmbH. This is an open access article under the terms of the Creative Commons Attribution Non-Commercial NoDerivs License, which permits use and distribution in any medium, provided the original work is properly cited, the use is non-commercial and no modifications or adaptations are made.</footnote>
<figure><img src="image_5.png" /><figcaption>**Figure 2.** Examples of reversible, bidirectional (irreversible) and mono-directional (irreversible) catalysis by *C. acetobutylicum* HydA (black), *T. mathranii* HydS (red) and *E. coli* Hyd-1 (blue) hydrogenases, respectively. Protein film voltammograms obtained with the enzymes adsorbed onto rotating carbon electrodes. Black: *C. acetobutylicum* [FeFe]-H₂ase (*Ca* HydA, replotted from ref. [30]), pH 7, 1 atm of H₂, 30 °C, scan rate 20 mV s⁻¹, electrode rotation rate 3000 rpm; red: *T. mathranii* [FeFe]-H₂ase (*Tam* HydS, replotted from ref. [11]), pH 7, 1 atm of H₂, 30 °C, scan rate 5 mV s⁻¹, electrode rotation rate 3000 rpm; red: *E. coli* [NiFe]-H₂ase (*Ec* Hyd-1 replotted from ref. [31]), pH 6, 1 atm with 10%H₂, 30 °C, scan rate 1 mV s⁻¹, electrode rotation rate 8500 rpm.</figcaption></figure>

the metal sites, buried within the polypeptide chain, and the surface of the enzyme. In this respect, turnover frequency values are rarely reported from protein film voltammetry and it is difficult to relate these numbers with activity values obtained from solution assays using soluble electron mediators.[9] We will only comment here the electrochemical data that allow to quantify the reversible vs non-reversible bidirectional or even monodirectional behavior of the enzymes.

In short, [FeFe]-hydrogenases immobilized at a carbon electrode generally display a reversible behavior.[10] One notable exception is the [FeFe]-hydrogenase from *Thermoanaerobacter mathranii*, characterized with an irreversible behavior (Figure 2).[11] [FeFe]-hydrogenases appear to be catalytically biased towards HER at pH < 7. Typical TOF are 5–20×10³ s⁻¹ for HER and 2–5×10³ s⁻¹ for HOR. HOR is favored under basic conditions (pH 8–9) where the catalytic

bias (i.e. the ratio between HER and HOR catalytic currents) was found close to unity for the *A. woodii* hydrogen-dependent CO₂ reductase (HDCR) enzyme.[10f] An exception is found with the [FeFe]-hydrogenase II from *Clostridium pasteurianum* (CpII) that is biased towards HOR over a large pH range including acidic values.[12]

On the contrary, immobilized [NiFe]-hydrogenases are rather biased towards H₂ oxidation under near neutral conditions with TOF for HOR generally below 2×10³ s⁻¹.[8c,13] Because [NiFe]-hydrogenases have an affinity (K<sub>M</sub>(H₂) is 10² atm⁻¹[13h]) for H₂ one order of magnitude higher than [FeFe]-hydrogenases,[10f] their HER activity is more strongly inhibited in the presence of H₂[14] resulting in low HER TOF values. While most [NiFe] hydrogenases display reversible behavior, the membrane bound O₂-tolerant [NiFe]-hydrogenase (Hyd-1) from *E. Coli* only catalyzes HOR at pH 7 (Figure 2); still a reversible behavior could be restored when the pH is turned acidic with a catalytic bias resuming to unity at pH 3.[13c] The electron transfer chain has been shown not being the cause for the observed bias in [NiFe]-hydrogenases.[15] This demonstrates that simple models[16] based on Butler-Volmer relations are insufficient to explain the catalytic bias in reversible catalysts and motivated the development of more refined models including a variety of active states.[17] Of note the observed inhibition of hydrogenase activity by H₂ is not controlled at the active site but relates to H₂ transport within the [NiFe]- or [FeFe]-H₂ases.[18]

The crystallographic structures of the two classes of hydrogenases revealing the details of their metal active sites are only available since the mid 90's with seminal contributions from the groups of Fontecilla-Camps,[19] Peters[20] and Higuchi.[21]

[NiFe]-hydrogenases contain a heterobinuclear Ni–Fe center bridged by two cysteinate ligands, while the Ni atom completes its strongly distorted coordination sphere with two additional cysteинates and the Fe atom with two carbon monoxide and one cyanide ligand (Figure 1).

The active site of [FeFe]-hydrogenases is a diiron cluster with carbon monoxide and cyanides as terminal ligands at each Fe atom and a bidentate bridging dithiolate ligand connecting the two Fe atoms (Figure 1). A cysteine, as the unique protein ligand, connects the diiron center to a [4Fe-

<figure><img src="image_6.png" /></figure>

Vincent Artero (1973) studied at the Ecole Normale Supérieure (Paris) and graduated in inorganic chemistry at Sorbonne Universities (PhD 2000). After a post-doctoral stay in RWTH Aachen with Prof. U. Kölle, he moved in 2001 to the Fondamental Research Division of CEA in Grenoble to develop bioinspired catalysis and artificial photosynthesis related to H₂ technologies. He currently leads the Laboratoire de Chimie et Biologie des Métaux and chairs the Board of Management of the ARCANE Excellence Laboratory Network (LABEX) for bio-driven chemistry in Grenoble.

<figure><img src="image_7.png" /></figure>

Cyrille Costentin studied at the Ecole Normale Supérieure Paris-Saclay and graduated at Université Paris Diderot (Ph.D. 2000) under the guidance of Prof. Savéant. After a postdoctoral fellowship at University of Rochester, he joined the faculty at Université Paris Diderot as Associate Professor in 2001. He was promoted Professor in 2007. He is an expert in molecular electrochemistry and was Visiting Scholar at Harvard University (2016–19) working with Prof. No-cera. In 2019 he joined the Département de Chimie Moléculaire at Université Grenoble-Alpes.
4S] cluster. The presence of a bridgehead nitrogen atom, playing a central role as proton relay or base during catalysis, in this bridging ligand was first intuited by Fontecilla-Camps in 2001[19d] and confirmed by Fan and Hall using quantum methods.[22] Still the very nature of this bridgehead atom has been long debated. A first experimental evidence for a nitrogen bridgehead atom was provided using ¹⁴N HYSCORE (Hyperfine sublevel correlation spectroscopy, a 2D pulsed EPR technique) by the Lubitz group,[23] still with the possibility that the identified residue is another amine moiety within the active site pocket. An elegant and definitive demonstration was provided by Berggren et al. via the artificial maturation of apo-hydrogenase with synthetic dinuclear active sites containing various bridgehead atoms, the so-called Grenoble insertion method, and the demonstration that the dithiomethylamine derivative is the only one displaying hydrogenase activity at a level similar to that of native enzymes.[24]

Conversely, the proton relay in [NiFe]-hydrogenases has been rapidly identified as one of the terminal cysteinate ligands by comparing the activity of [NiFe]-hydrogenase with that of [NiFeSe]-hydrogenases having a similar active site but a terminal selenocysteinate ligand.[19c,25] This hypothesis was also verified by quantum calculations[26] and recently confirmed by protein film voltammetry.[27] An arginine residue located on top of the active site could also play a key role in the mechanism.[28] For both [FeFe]- and [NiFe]-hydrogenases, the key transition state corresponding to the formation or cleavage of the H–H bond implies a dihydrogen bond[29] between the metal(s) and the proton relay as depicted in Figure 3.

<figure><img src="image_8.png" /><figcaption>**Figure 3.** Heterolytic H₂ formation and cleavage via the formation of a dihydrogen bond involving proton relay at the active sites of [FeFe] (top) and [NiFe]-hydrogenases. The structures shown between square brackets are the proposed transition states or the non-isolated intermediates that feature a dihydrogen bond as discussed in the text.</figcaption></figure>

### 3. Biomimetic and Bioinspired Approaches Towards HER Catalysis

The unusual and unique nature of the active sites of hydrogenases opened a very active field of research to mimics these structures with the aim of discovering novel catalysts for HER and HOR. This field has already been reviewed by us and others[32] and we will only summarize here the efforts at mimicking the proton relays installed in close proximity of the metallic sites. Very few mimics of the [NiFe]-H₂ase active sites have been designed with installed thiolate function able to act as a proton relay, mainly because nickel complexes with terminal thiolate ligands tend to oligomerize. Lubitz and co-workers reported the Ni–Fe complex 1 with such a terminal thiolate ligand (Figure 4) and could demonstrate that this ligand could be protonated, although under strongly acidic conditions that were not relevant to the observed catalytic HER activity.[33]

More interestingly, the hemilabile nature[34] of nickel-dithiolate fragments, viewed here as a ligand of the iron center allows a terminal thiolate function to be generated concomitantly to the installation of a metal-bound hydride ligand upon reduction of dithiolate-bridged nickel-iron models such as 2 or 3. DFT studies on several such complexes revealed that the generated nickel thiolate ligand can then be protonated under catalytic conditions to form a thiol that interacts with the metal hydride to generate H₂ (Figure 4).[35]

Much more work has been devoted to install azadithiolate or aryl derivatives in diiron mimics. In CO-rich Fe¹Fe¹ mimics however, and although the installed amine function could be protonated, this did not prove functional for catalysis, mostly because the complexes did not contain a site for hydride binding directly facing the installed proton

<figure><img src="image_9.png" /><figcaption>**Figure 4.** [NiFe]-H₂ase mimics featuring a thiolate function able to act as a proton relay.</figcaption></figure>
relay.[36] This is due to the fact that the most stable conformation contains a Fe–Fe bond as the site for protonation, while the “rotated” conformation that would contain a bridging CO ligand and display an open coordination site for hydride binding facing the bridgehead amine group is higher in energy (see discussion regarding complex 5 in part 6). Thus, the observed gain in catalytic performance compared to propyldithiolate analogues might be rationalized in terms of charge effect and/or lower reorganization energy for hydride formation upon protonation of the relay.

Rauchfuss and co-workers used 1,1-bis(diphenylphosphino) ethylene (dppv = cis-C₂H₂(PPh₂)₂) as ligand in such Fe¹Fe¹ mimics and were able to isolate compound 4 containing a terminal hydride and bridgehead ammonium moiety with a NH–HFe distance of 1.88(7) Å indicating dihydrogen-bonding (Figure 5).[37] This complex proved active for electrochemical H₂ evolution and the role of the azadithiolate ligand as proton relay to install the hydride ligand and produce hydrogen via proton-hydride coupling under strongly acidic conditions could be demonstrated.

Dey and co-workers designed the hexacarbonyl-diiron mimic 5 with an ortho-substituted dithiomethylaniline bridge (Figure 5).[38] Ortho substitution at the bridgehead aniline induces a conformation change due to steric reasons, thus preventing conjugation of the nitrogen doublet with the phenyl ring and directing this more electron-rich doublet towards one Fe site, where it was shown to act as a proton relay during HER after chemisorption onto graphite electrodes. This positioned proton relay also provides the complex with O₂ tolerance during catalysis. More recently, this complex has been shown to display the first bidirectional catalytic behavior among mimics of the active site of [FeFe]-H₂ase, as discussed in part 6.[39]

A breakthrough in the mimicry of the proton relays found at the active site of [FeFe]-H₂ase was reached when DuBois designed the series of “P₂N₂” (1,5-diaza-3,7-diphosphacyclooctane) and related ligands.[40] DuBois' catalysts, i.e. Ni(II) complexes with two P₂N₂ ligands (Figure 6), elegantly combine features of both active sites of [NiFe]-

<figure><img src="image_10.png" /><figcaption>**Figure 5.** [FeFe]-H₂ase mimics featuring an azadithiolate bridge able to act as a proton relay.</figcaption></figure>

<figure><img src="image_11.png" /><figcaption>**Figure 6.** Bioinspired Ni and Co complexes based on «P₂N₂» and related phosphine ligands.</figcaption></figure>

and [FeFe]-H₂ase, i.e. the Ni ion from [NiFe]-H₂ases and the pendant amine group from the [FeFe]-H₂ases but as part of a cyclic diphosphine ligand instead of the dithiolate ligand found in the enzyme. These catalysts were initially[41] studied in CH₃CN and display a remarkable activity for HER with low overpotential requirements.[42] Depending on the nature of the substituents at the N and P atoms and the acido-basic conditions of the medium, they preferentially catalyze the HER (in the presence of triflic acid or protonated DMF) or the HOR (see below) in the presence of trimethylamine. Correct positioning of the basic amine centers close to the nickel atom facilitates intra- and inter-molecular proton transfers and the coupling of proton and electron transfer processes as discussed in detail below. This series has been later on extended to mononuclear cobalt complexes with P₂N₂ ligands,[43] to nickel and cobalt complexes of related tetraphosphine ligand featuring the same P₂N₂ moiety (Figure 6),[44] as well as to nickel complexes with P₂N ligands.[45] All these catalysts have been intensively investigated to establish structure-function relationships regarding their catalytic properties for HER in link with the positioning and the pKₐ of the proton relays relative to the hydricity of the metal site.[40–41,42]

More akin to the proton relay topology found at the active site of [NiFe]-H₂ases, mononuclear aryldithiolate, aryldiselenoate and dithiolene Ni[46] Fe[47] and Co[48] complexes, such as 6 and 7 (Figure 7), have been described as HER catalysts and DFT studies have shown that, in the general case, the metal-coordinated sulfur atoms are the primary sites for protonation and serve as relay, first to install a hydride on the metal site and then to generate H₂ through the coupling of the hydride with another proton located on a metal-coordinated sulfur ligand.[46a,49] Marinescu then showed that such catalytic motif can also be incorporated in 2D metal-organic catalytic materials deposited on electrode surfaces.[50]

Metal coordinated N and O ligands have also been shown to act as proton relays. Representative examples are
<figure><img src="image_12.png" /><figcaption>**Figure 7.** Co-, Ni- and W-dithiolene complexes.</figcaption></figure>

provided by the bisdithiolene W⁴⁻=O complex **8** (Figure 7) that enjoys two subsequent PCET steps to generate first a W¹¹⁻OH complex and then a W⁴⁻H-OH intermediate, from which H₂ is evolved.[⁵¹]

The “NH” amine function within the tetraazamacrocyclic CoN₄H HER catalyst **9** (Figure 8)⁵² has also been proposed to act as a proton relay during the rate determining intramolecular H–H bond forming step (see below) from a CoIIH–NH/NH₂⁺ intermediate;⁵³ nonetheless, no consensus has yet been reached regarding the protonation (and therefore coordination) state of this amine ligand in the different states of the catalytic cycle.[⁵³⁻⁵⁴] DFT studies indicate a relatively high barrier for both intramolecular and intermolecular protonation of the CoII–H species with the coordinated amine compared to non-coordinated ammonium.[⁵³a]

Cobaloximes and cobalt diimine-dioxime complexes form a class of (pseudo)macrocyclic HER catalysts in which proton relays have been proposed to play a key role in catalysis.[⁵⁵] Complexes **10–12** of this family (Figure 8) contain a proton bridging two oximato functions to close the (pseudo)macrocyclic ligand. Protonation of this bridge leads to two protonated oxime functions. The possibility for these oxime functions to act as proton relays has been formulated

in 2009 for cobalt diimine-dioxime catalysts, including **12**.[⁵⁶] after the observation that protonation of the bridge was responsible for a strong anodic shift of the H₂ evolution catalytic wave.[⁵⁷] This system is detailed in part 5 and it has been shown that the oxime protons serve to pin the catalytic potential to more positive values but otherwise act as spectators during catalysis, while another, yet to be identified, site acts as proton relay.[⁵⁸]

Another example is the tetradentate equatorial ligand 6,6'-bis-(2-aminopyridyl)-2,2'-bipyridine (or bapbpy), featuring a redox-active bipyridine moiety connected to two pyridine units via non-coordinated NH linkers that can act as proton relays. The [CoII(bapbpy)Cl]⁺ complex **13** catalyzes HER at a potential 100 mV more positive when Et₃NH⁺ (pKₐ=9.3 in DMF) is used instead of acetic acid (pKₐ=13.5 in DMF) and this shift has been assigned to the protonation of one of the amino groups in the reduced bapbpy ligand.[⁵⁹] Inspired from the active site of the mononuclear [Fe]-hydrogenase,[⁶⁰] the installation on a Cp*Co platform of a pyridinone ligand that gets protonated in its reduced pyridinol state allowed Choudhury and coworkers to prepare catalyst **14** able to promote H₂ formation from the coupling of a Co-bound hydride with the pyridinol proton.[⁶¹]

Another strategy consists in incorporating a dangling proton relay to a catalytic platform, the challenge here being that steric constraints should be introduced to avoid coordination of the basic form of the relay to the metal.[⁶²] This was achieved on the cobalt diimine-dioxime complex **15** (Figure 9) through the use of a methyl substituted pyridine that cannot coordinate as in other diimine-dioxime Co complexes with pendant pyridine moieties.[⁶²] In that case, protonation of this relay does not modify the potential for

<figure><img src="image_13.png" /><figcaption>**Figure 8.** Cobaloxime and related cobalt-based catalysts.</figcaption></figure>

<caption>**Figure 9.** Bioinspired catalysts featuring dangling proton relays (other than Dubois' P₂N₂ ones).</caption>
Co reduction and theoretical study showed that the gen-
eration of the hydride is facilitated thanks to an intra-
molecular protonation step.[63]

Similarly, dangling pyridine moieties have been identi-
fied to act as proton relays to ease the formation of the
hydride or the formation of the H–H bond during HER
catalysis mediated by various Mn,[64] Ni[65] or Co[66] com-
plexes. In most cases such as for complex **16** (Figure 9), this
proton relay results from the decoordination of the pyridine
ligand in the course of the catalytic cycle under the protic
conditions used. Dangling hydroxymethyl groups have been
introduced in Ni, Fe and Co systems thanks to di or
tetraphosphine ligands, with very low overpotential
requirements.[67] Masuda and co-workers designed a bi-
phosphinopyridyl nickel catalyst **17** with two positioned
aniline groups that can act as proton relays during H₂
evolution catalysis.[68] Very recently, Dempsey and co-work-
ers studied the HER properties of the cobalt polypyridyl
complex **18** with a pendant pyridine base and could
demonstrate the implication of this proton relays both for
the intramolecular H₂-evolution step when a weak acid is
used and for the installation of the hydride ligand onto the
metal center if an acid strong enough to protonate the
pendant pyridine is used.[69]

Porphyrins have been extensively used as platforms to install proton relays. Nocera and co-workers developed Hangman Co porphyrins, such as **19** (Figure 9) with a positioned carboxylic moiety onto the Co macrocycle; they demonstrated that the presence of this proton relay allowed to lower the overpotential requirement for HER[70] via the acceleration of the protonation of the reduced Co⁰ center to yield a Co(II)-hydride intermediate.[71] The first order rate of 8.5×10⁶ s⁻¹ for this intramolecular process was determined as a result of a local acid concentration of > 3000 M (see below for an explanation of such high local concentration enabled by proton relays). While intramolecular proton transfer occurs after generation of the metal(0) state in the cobalt system, the Ni analogues experience proton transfer from the Ni(I) state to yield a Ni(III)-hydride which is further reduced before H₂ is formed. Similarly Hangman Fe porphyrins were shown able to mediate HER from the Fe¹ state in the presence of PPh₃ as an axial ligand.[72] Changing the proton-donor ability of the hanging group allowed to modulate the rate of catalysis by 3 orders of magnitude. Using different designs, Dey and co-workers also demonstrated that second sphere proton relays allowed to lower the overpotential requirement for HER by iron porphyrins such as **20** (Figure 9). Computational and experimental data indicated the formation of a dihydrogen bond in the Fe<sup>III</sup>H<sup>δ</sup>···H<sup>δ</sup>+N state<sup>[73]</sup> and linear free-energy correlations of the observed rate with the pK<sub>a</sub> of the acid source used suggested intramolecular proton transfer in the rate-determining step.[74]

In most of the examples listed above, the involvement of
proton relays, either synthetically installed on purpose or
generated through ligand decoordination during the catalytic
cycle, is inferred based on a body of experimental and/or
theoretical investigations. Yet, very few studies have made a
formal demonstration of the proton shuttling role played by

such proton relays through a quantification of the kinetic
advantage introduced. Before describing the cases for which
this has been achieved, or on the contrary refuted, through
detailed electrochemical investigations, we will first intro-
duce the kinetic and thermodynamic bases for these studies.

4. How Does It Work?

Protonation reactions are generally fast reactions, especially
in aqueous media when oxygen and nitrogen-based proton
donors or acceptors are concerned. This is however not the
case when the protonation or deprotonation step involves
making or cleaving a bond with significant covalency, such
as the formation or deprotonation of metal-hydride as well
as the protonation of metal hydride or the deprotonation of
a metal-bound dihydrogen molecule.[75] These key steps are
exemplified in Scheme 1 that also shows how a proton relay
can be involved in their kinetic enhancement.

That proton relays are required in enzymes to install protons on the spot ready for a reaction at the active site (for HER) or for removing it from the active site (for HOR) seems to be a reasonable idea because the active site is buried in the interior of the polypeptide chain, separated from the reaction medium. When it comes to molecular catalysts, either in solution or in some cases immobilized at an electrode surface, accessibility of the active site (the metal center in metal-complex catalysts) is much less of a problem. Therefore, questioning the origin of a boosting effect of proton relays and its limitations is legitimate. This problem was formally analyzed in a seminal contribution by Savéant using a simple reaction sequence so as to focus on the essential traits of the phenomenon.[76] In this section we will recall the main findings of this analysis and further consider slightly different hypotheses regarding the reaction sequence that are more adequate to the context of HER with molecular catalysts than the original ones.

The basic questions related to proton relays are twofold.
The first one is: why would it be beneficial that a
protonation step proceeds via a relay rather than directly
from a proton donor in solution? (Transposition to a
deprotonation step and a proton acceptor in solution is
straightforward and, for the sake of clarity, we will focus

<figure><img src="image_14.png" /><figcaption>**Scheme 1.** Key proton relay-assisted protonation/deprotonation steps involving the formation or cleavage of covalent bonds during the HER or HOR processes. Top: formation (forward) and deprotonation (backward) of metal-hydride; Bottom: protonation of metal hydride (forward) and deprotonation of a metal-bound dihydrogen molecule (backward).</figcaption></figure>
here on the case of a protonation, therefore associated to a reduction). The answer to that question is clearly of kinetic nature because the overall process is the same whether the proton goes through a relay or not.

Although the problem is of kinetic nature, thermodynamic considerations have also to be taken into account, as we will detail below. Indeed, the ability of the solution proton (AH) to protonate the relay moiety on the catalyst in its oxidized and/or reduced form will depend on the relative values of the corresponding pKₐ values as detailed below.

<figure><img src="image_15.png" /><figcaption>**Scheme 2.** Reaction schemes. (a) PCET square-scheme of a catalyst with a proton relay. (b) Mechanism in the strong acid regime (c) Mechanism in the weak acid regime (d) Mechanism with proton transfers to the relay at unconditional equilibrium.</figcaption></figure>

<figure><img src="image_16.png" /><figcaption>**Figure 10.** Pourbaix-type diagram for the PCET square scheme shown in Scheme 2a.</figcaption></figure>

The second question is: if metal protonation from the relay is kinetically more efficient than direct protonation from a solution proton, to which extent re-protonation of the relay from a solution proton may be a limitation of the relay boosting effect in a catalytic process?

We will answer both questions via a formal analysis considering the minimal reaction scheme involving the monoelectronic reduction of a substrate S<sub>ox</sub> associated to its protonation, hence leading to the formal product S<sub>redH</sub>. This overall irreversible reaction (S<sub>ox</sub>+e<sup>-</sup> + AH=S<sub>redH</sub>+A<sup>-</sup>) is catalyzed by a molecular catalyst (rP), typically a transition metal complex (P), bearing a basic moiety (r) possibly serving as a proton relay (noted Hr<sup>+</sup> when protonated). A PCET (Proton-Coupled Electron Transfer) square scheme can be considered for such a catalyst (Scheme 2a), thermodynamically characterized by a Pourbaix-type diagram (Figure 10). The Pourbaix-type diagram has been drawn here considering that the relay moiety is more acidic in the oxidized form of the catalyst than in the reduced form of the

catalyst, i. e. $\frac{K_{\mathrm{H}}^{\mathrm{O}}}{K_{\mathrm{H}}^{\mathrm{P}}} < 1$, (and thus

$E_{\mathrm{HrP}^{+}/\mathrm{HrQ}}^{0} - E_{\mathrm{rP}/\mathrm{rQ}_{-}}^{0} = \frac{RT}{F} \ln \frac{K_{\mathrm{H}}^{\mathrm{P}}}{K_{\mathrm{H}}^{\mathrm{O}}} > 0$) a likely situation in practice from simple electrostatic considerations. This situation is nonetheless at variance with the case discussed in reference[76] in which it was considered that $E_{\mathrm{HrP}^{+}/\mathrm{HrQ}}^{0} \approx E_{\mathrm{rP}/\mathrm{rQ}_{-}}^{0}$, i. e., that the redox state of the catalytic center had no influence on the acid/base properties of the relay and vice-versa.

According to the presently considered Pourbaix diagram (Figure 10), we can delineate two situations depending on the choice of the external source of proton AH. If a strong acid is used in solution with $pK_{a}^{\mathrm{AH}} < pK_{\mathrm{H}}^{\mathrm{P}}$, the predominant form of the oxidized catalyst in solution is HrP<sup>+</sup> and the catalyst redox couple to be considered is HrP<sup>+</sup>/HrQ leading to the mechanism depicted in Scheme 2b. Alternatively, if a weak acid is used in solution with $pK_{\mathrm{H}}^{\mathrm{P}} < pK_{a}^{\mathrm{AH}}$, the predominant form of the oxidized catalyst in solution is rP and the catalyst redox couple to be considered is rP/rQ<sup>-</sup> leading to the mechanism depicted in Scheme 2c. Both situations will be examined successively. In both cases we will consider irreversible proton transfers. Indeed, molecular catalysts are often studied in organic solvents in which acids are used in non-buffered conditions. Very weak acids are not considered because situations corresponding to $pK_{\mathrm{H}}^{\mathrm{Q}} < pK_{a}^{\mathrm{AH}}$ would lead to thermodynamically unfavorable protonation of the relay and are not of interest. Note, however, that it is worth examining the case of equilibrated proton transfers to describe systems operating in water, this will be done in the third sub-section considering the reaction scheme depicted in Scheme 2d.

In all cases, examining the effect of the proton relay relies on the analysis of the kinetic competition between two pathways: a pathway in which the relay is bypassed leading to a direct protonation of the substrate S<sub>ox</sub> via an external acid molecule (AH) while this substrate S<sub>ox</sub> is reduced by the catalyst Q ($k_e$ or $k_e^H$ steps in Schemes 2b, 2c, 2d), and a pathway in which the relay is involved via a step ($k_i$) where
the substrate is reduced by Q while it is protonated by the internal proton sitting on the relay r.

The analysis of the boosting effect of proton relays and its limitations is conveniently performed via description of the corresponding responses in cyclic voltammetry (CV). In that context, and for the sake of clarity, we make two additional simplifying assumptions: (i) the acid (AH) concentration in the bulk $C_{\text{AH}}^b$ is large enough to remain constant throughout the diffusion-reaction layer and (ii) all pseudo-first order rate constants ($k$) are large enough compared to the characteristic time-scale of the CV ($kRT/Fv \gg 1$) so that an S-shaped steady-state CV is obtained. The efficiency of the catalytic process is then evaluated considering the plateau current ($i_{pl}$). The mono-electronic PCET process considered here (see Scheme 2) is formally one-half of a HER process and therefore can be considered as the first PCET leading to the hydride (in such a case $S_{ox}$ is not to be considered) or a second PCET corresponding to reaction of the hydride ($S_{ox}$ being an acid molecule AH and $S_{red}$H being $H_2$), with, in both cases, rapid regeneration of the initial catalyst species.

## 4.1. Strong Acid Regime

According to Scheme 2b, in the presence of an exogeneous acid AH which $pK_a$ is smaller than the relay's $pK_a$, the catalyst couple to be considered is $HrP^+/HrQ$ and the competition between the relayed and non-relayed pathways occurs at the level of HrQ.

Considering first that reprotonation of the relay is fast, i.e. $\frac{k_i^P C_{\text{AH}}^b}{k_i} \to \infty$, the boosting effect of the relay is conveniently evaluated from an amplification factor defined as the ratio of the plateau current in the presence of the relay ($i_{Hr,pl}^{r+nr}$) and without the relay ($i_{Hr,pl}^{nr}$):

$$ \frac{i_{Hr,pl}^{r+nr}}{i_{Hr,pl}^{nr}} = \sqrt{1 + \frac{k_i}{k_e^H C_{\text{AH}}^b}} $$

(all rate constants are defined in Scheme 2; see the Supporting Information for details on the formula's derivation).

It immediately appears that the boosting effect is due to a local concentration effect resulting from the proximity of the proton to the catalytic center once installed on the relay.

The boosting term $\frac{k_i}{k_e^H C_{\text{AH}}^b}$ can indeed be rewritten as $\frac{k_i^P C_{\text{AH}}^{ap}}{k_i^P C_{\text{AH}}^b}$ where $k_i^P$ is an apparent second order rate constant and $C_{\text{H}}^{r,ap}$ the apparent local concentration. We thus emphasize that a key point of the boosting effect lies in the fact that the proton relay installed near a metal site generates a high apparent local concentration that increases the rate of these slow steps. As indicated in reference [76], a very rough estimate of this apparent local concentration can be made considering that the proton relay ensures that one proton ($1.66 \times 10^{-24}$ mol) is located at any time within a sphere of 2 Å diameter (thus with a volume of $4.2 \times 10^{-27}$ L) leading to

400 M (Figure 11). Therefore, $\frac{k_i^{ap} C_{\text{AH}}^{ap}}{k_i^P C_{\text{AH}}^b}$ is likely larger than unity even if $k_i^{ap} < k_e^H$ because the relay is less acidic than the exogeneous acid AH ($pK_a^{AH} < pK_e^Q$).

Considering now that there is indeed a boosting effect of the relay, i.e. $k_i \gg k_e^H C_{\text{AH}}^b$ but that the reprotonation of the relay may interfere kinetically, we can evaluate the limitation of the boosting effect of the relay (see the Supporting Information) knowing that

$$ \frac{\left( \frac{i_{Hr,pl}^{r+nr}}{i_{Hr,pl}^{nr}} \right)}{\left( \frac{i_{Hr,pl}^{r+nr}}{i_{Hr,pl}^{nr}} \right)_{\text{max}}} \approx \frac{1}{1 + \frac{1}{\sqrt{\frac{k_i^P C_{\text{AH}}^b}{k_i} \left( \sqrt{\frac{k_i^P C_{\text{AH}}^b}{k_i}} + 1 \right)}}} $$

Decreasing values of $\frac{k_i^P C_{\text{AH}}^b}{k_i}$ lead to limitation of the boosting effect, down to:

$$ \frac{i_{Hr,pl}^{r+nr}}{i_{Hr,pl}^{nr}} \approx \sqrt{\frac{k_i^P}{k_e^H}} $$

when the relayed process is fully controlled by reprotonation of the relay. In such a case, there is still a beneficial effect of the relay because it can be anticipated that $k_i^P > k_e^H$. Indeed, protonation of the catalytic center (whether it is formation of a hydride or protonation of an hydride to make dihydrogen), via $k_e^H$ or $k_i$, is kinetically difficult[75] whereas reprotonation of the relay is intrinsically fast because it involves proton transfer between nitrogen or oxygen acid/base couples.[77]

From this analysis, it clearly appears that the installation of a proton relay may be valuable because it allows to decorrelate the process of providing the proton close to the catalytic center and the actual protonation at the catalytic center which is intrinsically slow. First, thanks to the relay, and at the essential condition that it involves nitrogen or oxygen acid-base, bringing the proton close to the catalytic center is fast ($k_i^P$) whereas it is slow in the direct process ($k_e^H$). Second, thanks to the local concentration effect generated by the proximity of the proton to the catalytic

<figure><img src="image_17.png" /><figcaption>**Figure 11.** Illustration of the high proton concentration provided by a pendant ammonium group positioned in close proximity of a metal center.</figcaption></figure>
center, the intrinsically slow protonation step ($k_i^{ap}$ and $k_e^H$) is boosted ($C_H^{r,ap} \gg C_{AH}^b$).

## 4.2. Weak acid Regime

Considering now weak acids ($pK_H^P < pK_a^{AH} < pK_H^Q$), the catalyst couple to be considered is rP/rQ⁻ and the competition between the relayed and non-relayed pathways occurs at the level of rQ⁻ (Scheme 2c). When the local concentration effect of the relay makes $\frac{k_e C_{AH}^b}{k_i} \ll 1$ but also $\frac{k_e^{Q} C_{AH}^b}{k_i} \ll 1$, a maximal amplification factor is reached (see the Supporting Information):

$$ \left( \frac{i_{r,pl}^{r+nr}}{i_{r,pl}^{r+nr}} \right)_{\max} = \sqrt{1 + \frac{k_e^Q}{k_e}} $$

The boosting effect of the relay then originates from the faster proton transfer from the solution acid AH to the relay than to the catalyst active site $\frac{k_e^Q}{k_e} > 1$. If, alternatively, despite the local concentration effect, protonation of the relay is not the rate determining step of the relayed pathway, i.e., $\frac{k_e^{Q} C_{AH}^b}{k_i} \gg 1$, then:

$$ \frac{i_{r,pl}^{r+nr}}{i_{r,pl}^{r+nr}} = \sqrt{\frac{k_i}{k_e C_{AH}^b}} $$

A boosting effect is still possible provided $\frac{k_i}{k_e C_{AH}^b} > 1$.

Thus, as in the case of a strong acid, the value of having a proton relay relies on (i) a local concentration effect and (ii) fast protonation of the relay.

## 4.3. Fast Equilibrated Proton Transfers

So far, we have only considered irreversible protonation of the relay, either in its oxidized form (strong acid regime) or in its reduced form (weak acid regime). However, in aqueous media in the presence of a buffer, proton transfers between the buffer components and the relay may be fast enough to be considered at unconditional equilibrium in both oxidized and reduced forms (Scheme 2c), meaning that, whatever the timescale considered in CVs, the equilibrium is maintained without kinetic limitation. The amplification factor is then (see the Supporting Information):

$$ \frac{i_{pl}^{nr+r}}{i_{pl}^{nr}} = \sqrt{1 + \frac{k_i}{\left( k_e \frac{K_H^Q}{[H^+]} + k_e^H \right) C_{AH}^b}} $$

The boosting effect relies on the local concentration effect: $k_i = k_i^{ap} C_H^{r,ap} > k_e^H C_{AH}^b$ and/or

$k_i = k_i^{ap} C_H^{r,ap} > k_e \frac{K_H^Q}{[H^+]} C_{AH}^b$. If we assume a Brønsted type activation/driving force relationship for proton transfer from the relay or from AH (acid form of the buffer) to the active site,[78] we can write: $k_i^{ap} = A \sqrt{K_H^Q}$ where A is a constant and similarly $k_e \approx k_e^H = A \sqrt{K_a^{AH}}$. Moreover, we have $C_{AH}^b = \frac{C_{buffer}}{1 + K_a^{AH} / [H^+]}$ where $C_{buffer}$ is the total concentration of the buffer. Finally, if we consider that at a given pH, a specific buffer is used which $pK_a^{AH}$ is equal to the pH, then we have:

$$ \frac{i_{pl}^{nr+r}}{i_{pl}^{nr}} \approx \sqrt{1 + \frac{K_H^Q}{[H^+]} \frac{2}{(1 + \frac{K_H^Q}{[H^+]} )} \frac{C_H^{r,ap}}{C_{buffer}}} $$

thus, showing that the maximal amplification factor, is obtained for pH = $pK_H^Q$ and is given by the local concentration effect:

$$ \left( \frac{i_{pl}^{nr+r}}{i_{pl}^{nr}} \right)_{\max} = \sqrt{1 + \frac{C_H^{r,ap}}{C_{buffer}}} $$

This optimal boosting effect results from the fact that the relayed pathway is in competition with two direct reactions, one being favored in acidic media via the intrinsic acceleration of $k_e^H$, the other being favored in basic media via the increase of the equilibrium concentration of the deprotonated relay.

# 5. Deciphering the Role of Proton Relays in H₂ Evolution Catalysis

Surprisingly, despite intensive efforts aiming at understanding the molecular mechanisms of catalysis mediated by the biomimetic and bioinspired systems containing proton relays described in section 3, very few reliable kinetic data are available that allows to unambiguously decipher the role of the proton relay on the key steps in catalysis. In this section, we will show how the use of the most recent methodologies of molecular electrochemistry can provide quantitative insights into the role of the pendant relays on catalytic hydrogen evolution. It is a fact that only few studies contain reliable data from which it is possible to extract rate constants for the successive protonation steps,[42] namely the key steps corresponding to the formation of the hydride and of the H-H bond (Scheme 1). Of note, the examples below are studied in organic media and we will show that reprotonation of the relay are not always fast processes which complicates the analysis.
## 5.1. Effect of a Proton Relay on the Formation of a Hydride Intermediate

Although the rate of the first protonation step, directly or indirectly leading to the formation of a hydride intermediate, is often reported in the literature for HER catalysts there is only rare definitive evidence of the boosting role of the proton relay for this specific step. As mentioned above, Nocera and co-workers could quantity both intra and intermolecular rate constants for the protonation of a Co⁰ center leading to a Co¹⁻H species in the case of the hangman Co porphyrin 19 with a pendant benzoic acid and a similar compound that lacks the proton relay, respectively. The intramolecular protonation rate constant of 8.5×10⁶ s⁻¹ equates to an intermolecular process (2nd order rate constant of 2500 M⁻¹s⁻¹, as observed when the proton relay is lacking) with an apparent local concentration of >3000 M.[79] In a similar metal-free hangman chlorin system,[80] the kinetic ratio between the unimolecular rate constant for the protonation of the macrocycle from a pendant benzoic moiety and the bimolecular rate constant for protonation from an external acid (benzoic acid) is only of the order of 2 mM. Large concentration of external acid can thus bypass the relay. Larger local concentration may thus be reached if the positioning of the relay is optimized.

While the proton relay in the hangman systems is protonated in the oxidized state (Scheme 2b), the first series of DuBois Ni(P₂N₂)₂ catalysts (Figure 6) are only protonated after reduction of the Ni center in organic media and their mechanism follows Scheme 2c in the case of hydride formation. The preferential protonation of the ligand (vs the metal), under thermodynamical control, has been demonstrated for DuBois's Ni(P₂N₂)₂ catalysts following the establishment of Pourbaix-type diagrams.[81] However, the first protonation constants determined for these systems[42] vary from 8×10² M⁻¹s⁻¹ and 1.3×10³ M⁻¹s⁻¹ for [Ni(P₂PhN₂Ph)₂]²⁺[41b] and [Ni(P₂PhN₂C₆H₄CH₂P(OEt)₂)₂]²⁺[41c,42] to 7.5×10⁴ M⁻¹s⁻¹ for [Ni(P₂PhNPh)₂]²⁺,[45a] and rather reflect a slow intermolecular proton transfer process to the ligand in organic media depending on pKₐ values of the proton source and the ligands as well as steric factors. This process is likely followed by a fast intramolecular proton transfer leading to the hydride species. Similarly, protonation of the pendant pyridine base of the cobalt polypyridine complex 18 studied by Dempsey and coll. is diffusion-controlled when the acid is strong. Reduction triggers fast tautomerisation to form a hydride derivative.[69]

A similar conclusion can be made for cobaloximes. In a comprehensive study involving the [Co(dmgBF₂)₂(CH₃CN)₂] complex,[82] Dempsey and co-workers showed that the first protonation rate constants range from 2×10⁴ to 2×10⁸ M⁻¹s⁻¹ depending on the pKₐ of the acid used with a negative linear free energy relationship that only deviates near the diffusion limit. Such high protonation constant does not seem to match with the direct protonation of a metal center[75a] and suggests the intervention of a proton relay, yet to be identified (see below). The same consideration may hold for the cobalt diimine-dioxime catalyst [Co((DO)₂BF₂)pn]Br₂]

for which the rate constant of the first protonation event was found to be as high as 1.2 10⁶ M⁻¹s⁻¹.[58]

## 5.2. Intramolecular Proton-Hydride Coupling: Can It Be Diagnosed?

Among the molecular catalysts reported in section 2 and for which kinetic data are available, there are few examples for which the first protonation yielding the metal-hydride intermediate is not the rate determining step. Cobaloximes 10–11,[42] cobalt diimine-dioxime 12,[58] and the tetraazamacrocyclic cobalt complex 9[53c] all exhibit a 2nd order rate law for the second protonation step at low acid concentration, with a rate constant lower than that of the first protonation step. If this second step involves a proton relay, the observation of a 2nd order rate law indicates that the rate determining step is the (re)protonation of the relay but it could as well reflect intermolecular protonation of a hydride formed during the first protonation step. Still, for these complexes and others, interestingly, a distinct regime of catalysis has been observed at high acid concentration. Under such conditions, the catalytic current becomes independent of the acid concentration and a first order rate constant kᵩ can be determined.

The first example has been reported by Dempsey et al. with [Co(dmgBF₂)₂(CH₃CN)₂] that exhibits a kᵩ value of 125 s⁻¹.[82] A similar behavior was then reported for the related [Co((DO)₂BF₂)pn]Br₂] catalyst with a much higher kᵩ rate constant (1.5×10⁴ s⁻¹).[58] We also present in Figure 12 the cyclic voltammograms measured in CH₃CN with complex 9 in the presence of *p*-cyanoanilinium as the proton source in which we extend the acid concentration to higher values compared to our precedent study.[53c,83] Here again we observe that the catalytic current saturates above 80 mM acid, allowing to estimate that the value of kᵩ is of the same order of magnitude as that measured for [Co(dmgBF₂)₂-(CH₃CN)₂]. The molybdenum-sulfur dimer [(CpMoS)₂(µ-S₂CH₂)] complex reported by Appel et al.[84] also display catalytic current saturation with a kᵩ rate constant of 1.5 ± 0.3 s⁻¹ after a first regime in which the catalytic current linearly varies with the acid concentration. A similar behavior was observed with the bisphosphinopyridyl nickel complex 17 with positioned aniline ligands reported by Masuda and co-workers.[68] Very recently, Dempsey and coll. could also evidence that protonation of a pendant pyridine occurs after formation of the Co(II)–H intermediate of complex 18, also leading to saturation of the catalytic current that becomes independent of both the acid concentration and the acid strength. A kᵩ rate constant of 4.5 10² s⁻¹ could be determined.[69]

This behavior could be rationalized as follows: at low acid concentration, the observed rate determining step is the (re)protonation of the relay, in line with catalytic rate being first order of the acid concentration. Raising the acid concentration then allows to reach a catalytic regime where the rate determining step is *intramolecular* proton-hydride coupling resulting in H₂ formation. A kinetic model has been developed that takes the later hypothesis into account
<figure><img src="image_18.png" /><figcaption>**Figure 12.** a) Cyclic voltammograms (CV) of 9 (0.5 mM) in CH<sub>3</sub>CN with 0.1 M *n*Bu<sub>4</sub>NBF<sub>4</sub> as supporting electrolyte recorded at a glassy carbon electrode (1.6 mm diameter) in the absence and in the presence of 5, 10, 20, 40, 80, 120 and 160 mM of *para*-cyanoanilinium tetrafluoroborate. CV recorded at a scan rate of 100 mV s<sup>-1</sup>. b) Plot of the peak current *i*<sub>p</sub> vs the square root of the acid concentration.</figcaption></figure>

in the context of a 2 e⁻ + 2 H⁺ mechanism[58] to determine the k<sub>Ω</sub> rate constants. Still it is not easy to identify the proton relay in the cobaloxime and cobalt diimine-dioxime complexes **10–12** and the possibility that the amine function in the tetraazamacrocyclic cobalt complex **9** acts as a proton relay is debated.[53a,c,85] This intramolecular step can also be the decoordination of the H₂ molecule: Then, the mechanism would not involve proton relay but consist first in an irreversible protonation of the hydride with an external acid leading to a coordinated H₂ that subsequently goes off at a rate corresponding to k<sub>Ω</sub> which becomes the rate determining step of the catalytic process at high acid concentration. Finally, saturation at high acid concentration leading to a k<sub>Ω</sub> rate constant to 20 s⁻¹ for DuBois' Ni(P<sup>Ph</sup><sub>2</sub>N<sup>Ph</sup><sub>2</sub>)₂ catalyst (Figure 6)[55b] was assigned to the presence of off-cycle intermediates with protonation of the ligand in exo position where the protonated relay is unable to interact with the metal center.[86]

**5.3. The Proton Relay That Never Was: Installing Is Not Enough!**

As explained in section 4, it is of upmost importance for a proton relay to operate during HER that the reprotonation of the relay occurs much faster than the rate-determining step consuming the relayed proton, so as to maintain its high apparent local concentration. In aqueous solution and with oxygen and nitrogen-based acids, proton transfers are usually greater than 10⁶ M⁻¹s⁻¹[75a] but such reaction may be slower in non-aqueous solvents.

An informative example is shown by the cobalt diimine-dioxime complex **12**,[58] that features a proton in the bridge between two equatorial oximato functions. In earlier HER studies using this complex as catalyst, it has been observed that the catalytic potential is shifting anodically depending on the strength of the acid used, which was assigned to a protonation of the bridge to generate to equatorial oxime functions.[56,87] DFT calculations confirmed such a shift in potential upon bridge protonation[57,88] and also suggested that the additional proton could interact with a Co<sup>II</sup>–H intermediate to generate H₂ in an intramolecular step, though with a quite high calculated activation energy.[57] Finally, the establishment of the right acid concentration range to reach canonical conditions to observe catalysis[89] confirmed that catalysis occurs from a bridge protonated intermediate[58] and allowed to observe a regime independent of acid concentration with k<sub>Ω</sub> = 190 ± 30 s⁻¹ for an intramolecular step that could be H₂ formation from a Co<sup>II</sup>H-(LH) intermediate. This last step would generate the initial Co(II) complex with a singly protonated bridge. For the HER to cycle again, the bridge must be reprotonated with a rate constant greater than 190 s⁻¹. However, the variation of the Co<sup>III/II</sup> response with acid allowed to determine an upper value of 200 M⁻¹s⁻¹ for the second order rate reprotonation rate constant, and thus a maximum reprotonation rate of 24 s⁻¹ as the acid concentration used is below 120 mM. This clearly indicated that the additional proton positioned on the dioxime bridge cannot act as a proton relay during catalysis but rather acts as a spectator, simply influencing the catalytic potential via electronic and electrostatic effects.[90] Therefore we were left with having to postulate the presence of another proton relay to explain the observed behavior but H₂ formation could also results from intermolecular protonation with the intramolecular step being H₂ dissociation from the complex.

As a take-home message, the installation of a proton relay through bio-inspired synthetic design is not a master key for successful catalytic enhancement. In the latter case, electronic and electrostatic effects resulting from protonation indeed shifts the potential at which catalysis occurs but has no effect on the actual rate of catalysis.

It is also frustrating to conclude that only in the case of the nickel phosphinopyridyl complex **17** and the cobalt polypyridine complex **18**, it is possible to link the presence of a pendant base installed by design and the observation of a first-order rate-determining step that does not depend on the proton concentration that could be related to intramolecular proton-hydride coupling. It is thus worth to
devote efforts in the future to understand the mechanism that underlies behaviors similar to that shown in Figure 12 and decipher if it really involves proton relays, or not.

**6. Frustrated Lewis Pairs: a Key Structural Feature for Catalytic H₂ Activation**

The activation of H₂ and, beyond, catalytic HOR introduces additional constrains: it requires both H₂ binding and the cleavage of the covalent H–H bond. Here again, the mechanism at play at the active sites of hydrogenases shows how a proton relay can polarize this H–H bond and ultimately trigger in its heterolytic cleavage yielding a protonated relay and a metal hydride. A key structural feature is therefore the possibility for the unprotonated form of the relay to establish a frustrated Lewis pair with the metal site so as to activate a coordinated H₂ molecule.

This feature could be reproduced in several Fe,[91] Ni,[45b] Mn[92] and Mo[93] complexes of the P₂N₂ ligands that spontaneously react with H₂ to either form metal-H₂ adducts[91] or proceed to heterolytic H₂ cleavage.[45b,92-93] These complexes generally catalyze rapid H/D exchange when placed in a H₂ atmosphere in the presence of a deuterated acid (or a D₂ atmosphere with protons in solution) which reveals the existence of an equilibrium between bound H₂ molecule and the metal hydride species with a protonated relay. The structure of the H₂ adduct of the iron complex 21, solved via single-crystal neutron diffraction, displays a Fe–H···H–N dihydrogen bond with a remarkably short H···H distance of 1.49 Å between the protonated relay N–H⁶⁺ and the iron bound hydride in Fe–H⁶⁻ as hypothesized at the active site of [FeFe]-hydrogenases (Figure 13).[94]

A full picture (Figure 14) of the heterolytic H₂ splitting has been obtained by metadynamic simulation for the

<figure><img src="image_19.png" /><figcaption>**Figure 14.** Classical-nuclei free-energy landscape determined using hybrid DFT(PBE)/MM metadynamics simulations for the H₂ splitting by the all-boat [Ni(P<sub>Cy</sub><sub>2</sub>N<sub>Me<sub>2</sub></sub>)<sub>2</sub>]<sup>2+</sup> in acetonitrile projected onto the nitrogen/hydrogen coordination number of the two ligands and the H–H distance. Local minima and transition states between them are marked with blue and red dots, respectively, with free energy value in kcal mol<sup>-1</sup> indicated by a number placed nearby. The dotted lines indicate lowest-free-energy pathways: lowest-energy path (heterolytic H₂ splitting) as blue line, highest-energy path (homolytic H₂ splitting) as green line. All of the Ni complexes shown here have a +2 charge that is not shown in the figure. Reproduced with permission from ref.[95].</figcaption></figure>

<figure><img src="image_20.png" /><figcaption>**Figure 13.** Mono- and dinuclear iron complexes catalyzing the hydrogen oxidation reaction. Top right: single-crystal neutron diffraction structure of the hydrogen adduct of 21. Reproduced with permission from ref.[94].</figcaption></figure>

[Ni(P<sub>Cy</sub><sub>2</sub>N<sub>Me<sub>2</sub></sub>)<sub>2</sub>]<sup>2+</sup> complex (Figure 6).[95] A side-on H₂ adduct (bottom left in Figure 14) is first formed, then the H₂ molecule is polarized to yield a proton-hydride form. A last step is then observed yielding a Ni⁰ complex with a bis-protonated ligand set that has been characterized in solution[96] but may not be involved in the catalytic cycle in the presence of proton acceptors.

While the nickel complexes have early been shown to catalyze the HOR in CH₃CN,[41a,b] their Fe (Figure 13),[97] Ru[98] and Mn[99] analogues require more specific conditions, namely non-coordinating solvents such as fluorobenzene, and/or bulky substituents preventing the binding of the amines introduced as external bases to capture protons generated by the HOR to be able to cycle. Interestingly, the addition of an outer coordination sphere proton relay allowed to boost the catalytic rate for HOR by one such iron complex.[97a]

Similar strategies have been implemented in diiron mimics of the active site of [FeFe]-H₂ases. Namely, a dppv diiron bisdithiolate derivative **22** (Figure 13) incorporating
an N-benzyl group that reportedly functions as a proton relay unit and complemented with a functional electron transfer relay, in the form of a phosphine functionalized ferrocene group (FcP*) to mimic the whole H-cluster was shown by Rauchfuss and co-workers to catalyze HOR in CH₂Cl₂ solution in the presence of 6 equivalents of base (P(o-tolyl)₃) and excess of oxidant (4 equiv. of FcBAR₄, BAR₄⁻ = tetrakis[3,5-bis(trifluorométhyl)phenyl]borate) with 2 turnovers achieved over a 5 h period at atm H₂ pressure and 25°C (Figure 13).[100] Another diiron bisdithiolate complex **23** with a P₂N ligand acting as proton relay positioned so as to be able to interact with the bridging position between the two Fe centers has also been reported by Sun and co-workers (Figure 13) to be competent for HOR in CH₂Cl₂ in the presence of 20 equivalents of both base (P(o-tolyl)₃) and oxidant (FcBAR₄) with 6 turnovers achieved within 10 h at atm H₂ pressure and 25°C.[101]

A breakthrough was recently achieved by Ahmed et al. via the introduction of an ortho-substituted dithiomethylaniline bridge (Figure 5) within the hexacarbonyl diiron mimic **5** of the active site of [FeFe]-hydrogenases.[39] This complex exhibited bidirectional H₂ evolution and oxidation after chemisorbed onto various carbon electrode materials. A key intermediate (**5H** in Figure 15) could be characterized that corresponds to the monoreduced complex with a protonated anilinium ligand that is the resting state stable in a 500 mV range of potential between the HER and HOR processes that both display PCET behavior with a catalytic potential dependence on the pH of the solution with a 59 mV/unit slope. DFT calculations allowed to show that the ortho substitution at the bridgehead aniline induces a conformation change due to steric reasons, thus preventing conjugation of the nitrogen doublet with the phenyl ring and directing this more electron-rich doublet towards one Fe

<figure><img src="image_21.png" /><figcaption>**Figure 15.** Proposed mechanism for bidirectional H₂ evolution and oxidation mediated by complex **5**.[39]</figcaption></figure>

site. At the same time, this conformational change releases steric constrain so as to allow the binding of H₂ at a Fe site in inverted conformation, setting the scene for its heterolytic cleavage thanks to the properly positioned amine moiety. Of note the positioned proton relay may also play a key role during the oxidative deprotonation of the hydride to regenerate the starting complex as demonstrated by Hammarström and co-workers on another model system.[102]

## **7. From Bidirectional to Reversible Catalysis**

As indicated in the introduction, hydrogenase enzymes have the peculiar ability to catalyze the interconversion of protons and electrons into H₂ in water reversibly, meaning that besides bidirectionality, i. e., catalysis in both directions (HER and HOR), catalysis occurs at a substantial rate at low overpotential. Such a definition of reversibility is qualitative and efforts are currently devoted not only to clarify it but also to decipher the reasons that can make a non-platinum group metal catalyst, either an enzyme or a molecular catalyst, working reversibly. In that endeavor, it has been suggested that the presence of proton relays may be a useful characteristic, extending beyond their capability to promote H₂ activation as detailed in section 6.

De facto, besides enzymes themselves, the only reversible catalysts for HER/HOR are based on the Ni(P₂N₂)₂ platform designed by DuBois (Figure 6). The first evidence for a reversible behavior has been described for classical [Ni(P<sub>2</sub>N<sub>2</sub>)₂]<sup>2+</sup> complexes (R=Cy, Ph) grafted on multiwalled carbon nanotubes (MWNT) via the R' residue thanks to covalent[103] (Figure 16) or π-π[104] interactions. It is remarkable that the same complexes in organic solution only display monodirectional catalysis in one or the other direction depending on the operating conditions (acidic for HER and basic for HOR). It is noteworthy that this reversible behavior was observed under aqueous acidic conditions where deprotonation and reprotonation of the proton relay is expected to be fast. Later on, reversible catalysis was observed in solution for a similar complex with R'=CH₂CH₂OCH₃.[105] Finally, Shaw and co-workers designed a series of catalysts based on amino-acids inserted within the P₂N₂ ligands that show reversible behavior,[55a,106] sometimes in aqueous conditions,[17c,55a,106] as we will detail below in one representative example.[17c,55a,106] This activity was maintained after immobilization onto nanocarbon electrodes using covalent[107] or electrostatic[108] interactions, allowing to construct noble-metal free H₂ fuel cells[108–109] and to reach industrially relevant HOR current densities.[110]

Yet, future progress towards better catalytic systems depend on a better understanding on how the presence of proton relay(s) provides the catalysts with such a reversible catalytic activity. We will thus first briefly discuss the choice of a specific mechanistic reaction scheme that seems relevant for a catalyst bearing proton relays, operating in water, and catalyzing both HER and HOR. Then, in the framework of cyclic voltammetry studies we will recall the diagnostic criteria characterizing the bidirectionality and reversibility of a catalytic system following the chosen
<figure><img src="image_22.png" /><figcaption>**Figure 16.** Top: Structure of MWNT electrodes covalently functionalized with a [Ni(P<sup>Ph</sup>,N<sup>Ar</sup>)<sub>2</sub>]<sup>2+</sup> complex; bottom: Linear sweep voltammetry experiments (2 mV s<sup>-1</sup>) measured in 0.5 M H<sub>2</sub>SO<sub>4</sub> aqueous solution at membrane-electrode assemblies consisting of a gas diffusion layer (GDL) assembled with a Nafion membrane: black, unfunctionalized GDL; blue, GDL with [Ni(P<sup>Ph</sup><sub>2</sub>N<sup>Ar</sup>)<sub>2</sub>]<sup>2+</sup>-functionalized MWNT under an atmosphere of N<sub>2</sub>; red, GDL with [Ni(P<sup>Ph</sup><sub>2</sub>N<sup>Ar</sup>)<sub>2</sub>]<sup>2+</sup>-functionalized MWNT under an atmosphere of H<sub>2</sub> (10<sup>5</sup> Pa). The x-axis is positioned at null zero current density. Diamonds represent stabilized values of current at electrodes set for 10 minutes at various potentials. Reproduced with permission from ref.[103].</figcaption></figure>

pertinent mechanism. We will then show how intrinsic thermodynamic and kinetics parameters of the catalyst control reversibility with particular emphasis on the specific features associated with the presence of proton relays. Finally, application to an experimental example will be discussed.

## 7.1. Mechanistic Reaction Scheme

In section 4, we have discussed the conditions for a boosting effect of a proton relay in the framework of an irreversible system and considering only monoelectronic processes to get the nitty-gritty of proton relay effect. As far as reversibility is to be considered for the reversible catalysis of the 2 H<sup>+</sup> + 2 e<sup>-</sup> ⇌ H<sub>2</sub> reaction catalyzed by molecular systems, bielectronic mechanisms have to be considered. Several mechanisms can be proposed differing by the sequence of electron and proton transfers along the cycle. Because of the presence of a proton relay, and considering only aqueous media where proton transfers between the buffer components and the relay may be fast enough to be considered at

unconditional equilibrium, we postulate that only the two chemical steps previously discussed (Scheme 1) are kinetically relevant: (i) binding/release of molecular H<sub>2</sub> associated to its heterolytic cleavage/formation and (ii) formation/cleavage of a metal hydride as a key intermediate via proton transfer from/to the relay (red arrows in scheme 3). These two steps are the heart of the transition metal catalysis at work, where changes in coordination and oxidation state take place. That both steps can be operative at substantial rates is, at least partially, the result of the local concentration effect previously described. Therefore, in the present section, we do not consider competition with direct reaction of the metal center with the buffer, hence we will focus on the role of proton relay to facilitate reversibility, considering that the boosting role of the relay is actually achieved via the local concentration effect. We also consider that the two

<figure><img src="image_23.png" /><figcaption>**Scheme 3.** Catalytic reaction scheme involving N- or O-based proton relay under consideration. In the green notations, the relays are noted r when deprotonated and <sup>+</sup>Hr when protonated, M represents the central metal with corresponding oxidation state and hydride ligands (H<sup>-</sup>) are simply denoted H<sup>-</sup>. The total charge of the complex can be obtained by addition of the charge of the relay, the oxidation number of the metal, the charge of the hydride (if present) and the charge of other ligands. The pertinent electron transfer and proton transfer thermodynamic constants are noted alongside the corresponding reaction. The red equilibrium arrows correspond to the two kinetically relevant steps discussed in the text.</figcaption></figure>
pertinent chemical steps are separated by an electron transfer event implying oxidoreduction of the hydride derivative (green arrows in Scheme 3).

The four pertinent species (rMII, +HrM', rMIIIH- and + HrMIIH- in Scheme 3) are connected by proton-coupled electron transfers remaining at unconditional equilibrium because they involve fast electron transfers with the electrode and fast proton transfers between oxygen or nitrogen-based acid/base functionalities and the solution buffer (Scheme 3).

We note that alternative mechanisms could be considered, in particular regarding the separation of the two electron transfer events by the kinetically relevant chemical steps or not. Nonetheless, we focus here on the reaction scheme depicted in scheme 3 because of its relevance regarding the experimental example analyzed in section 7.4.

## 7.2. Diagnostics for Bidirectional and Reversible Catalysis

A convenient tool to diagnose bidirectionality and reversibility of a molecular catalytic system for HER/HOR is cyclic voltammetry (CV).[111] Provided the relevant rate constants are fast enough on the timescale of the CV (equal to RT/Fv with v being the scan rate), pure kinetics condition is obtained. Combined with negligible consumption of the substrates (H₂ and AH/A⁻) during the potential scan, pure kinetics condition leads to S-shaped current-potential responses corresponding to a steady-state regime. Considering a homogeneous catalyst, the steady state results from compensation between diffusion and the fast chemical steps parallel to electron transfer.

The additional crucial consideration is that, because of the proton relays, the proton transfers in both square schemes in Scheme 3 are at unconditional equilibrium and therefore the reaction scheme in Scheme 3 can be conveniently replaced by the simplified scheme shown in Scheme 4 corresponding to an apparent ECₜECₜ mechanism (E designates an electron transfer step and Cₜ designates a reversible chemical step) with apparent thermodynamic and kinetics constants (noted with stars in Scheme 4) which expressions are detailed in the SI.

<figure><img src="image_24.png" /><figcaption>**Scheme 4.** Equivalent reaction scheme simplified from scheme 3.</figcaption></figure>

Within this framework, a dimensionless equation of the CV, $\frac{i}{|i_{pt}|}$ vs. $\xi = \frac{F}{RT}(E - E_{eq})$, can be written as (see the Supporting Information):

$$ \frac{i}{|i_{pt}|} = \frac{e^{2\xi} - 1}{1 + \left(\frac{2}{p} - 1 - d\right)e^{\xi} + d e^{2\xi}} \quad (1) $$

where $|i_{pt}|$ is the absolute value of the cathodic plateau current, $d = |i_{pt}|/i_{pt}^a$ is a measure of the directionality (bias) of the catalysis as it is the ratio of the cathodic and anodic plateau currents. $p$ is the slope of the current at the equilibrium potential. Note also that both contributions to the current at the equilibrium potential are: $\frac{|i_{eq}^-|}{|i_{pt}|} = \frac{i_{eq}^-}{|i_{pt}|} = \frac{i_{eq}^-}{p}$, thus indicating that parameter $p$ describes the slope of the CV at the equilibrium potential ($E_{eq}$) but also the rate of catalysis at the equilibrium potential. The shape of the CV is hence fully described by two parameters $d$ and $p$ noting that $0 \le p < p_{max} = 2/(d+1)$. $d$ is an obvious metrics quantifying the directionality of catalysis. Regarding reversibility, for a given value of $d$, there is a transition from a behaviour corresponding to full reversibility when $p \to 2/(d+1)$ to full irreversibility when $p \to 0$. Few typical examples are shown in Figure 17. Intermediate situations can be qualitatively described as quasi-reversible. We note that a different parameter than $p$ can be used to evaluate the reversibility that is, for example, the separation between the anodic and cathodic waves using the catalytic potentials introduced in reference [17b] which can be more convenient for system with large reversibility while $p$ is more convenient for quasi-reversible systems. Note also that the dimensionless equation **1** is actually valid for any ordered mechanism, not only for the apparent ECₜECₜ mechanism considered here.[17b]

## 7.3. Relating Reversibility to Intrinsic Factors: Role of Proton Relays

In the framework of the considered reaction scheme depicted in Scheme 3 and replaced by the simplified scheme shown in Scheme 4 corresponding to an apparent ECₜECₜ mechanism, we just recalled that the CV shape is determined by two parameters $d$ characterizing directionality and $p/p_{max}$ characterizing reversibility. Equation (1) can be recast as:

$$ \frac{i}{|i_{pt}|} = \frac{e^{2\xi} - 1}{1 + (d + 1)\left(\frac{p_{\max}}{p} - 1\right)e^{\xi} + d e^{2\xi}} \quad (2) $$

One key question is thus to decipher the intrinsic characteristics of the system controlling these parameters and in particular the reversibility. Therefore, it is necessary to provide the relationship relating $d$ and $p$ to the apparent thermodynamic and kinetics constants $K_1^*, K_2^*, k_1^*, k_2^*, E_{sq,1}^0$,
<figure><img src="image_25.png" /><figcaption>**Figure 17.** Dimensionless CV corresponding to a EC,EC, mechanism (eq. 1) with (a) *d* = 1, (b) *d* = 10 and (c) *d* = 0.1 In each case, (left) *p* = *p*<sub>max</sub> = 2/(d + 1). The dotted lines indicate the cathodic and anodic contributions and the dashed line shows the slope at the equilibrium potential. (right) *p*/*p*<sub>max</sub> =1 (black), 0.5 (purple), 0.1 (pink), 0.01 (orange) and 0.001 (grey).</figcaption></figure>

$E_{\text{sq,2}}^0$ and hence to the true constants defined in Scheme 4. First and foremost, it is useful to take a look at both Pourbaix diagrams corresponding to the two square schemes in Scheme 4 and represented in Figure 18. In order to get as predominant species those involved in the kinetically relevant steps, it is required that the operative pH is in between the *pK*<sub>a</sub>'s of the square schemes. As a consequence, to get the largest range of pH fulfilling this condition, a first favorable condition for having a bidirectional and reversible system is to have a large difference between *pK*<sub>H</sub><sup>1</sup> and *pK*<sub>H</sub><sup>II</sup> on the one side and between *pK*<sub>H</sub><sup>IIIH</sup> and *pK*<sub>H</sub><sup>IIIH</sup> on the other side. Thus, for the remainder of this section we will consider that the operative pH is in between those values, leading to the mechanistic scheme in Scheme 5.

We can refer potentials to the equilibrium potential $E_{eq}$ and introduce the dimensionless potentials

<figure><img src="image_26.png" /><figcaption>**Figure 18.** Pourbaix diagrams corresponding to square schemes 1 and 2 in scheme 3. Reproduced with permission from ref. [17c].</figcaption></figure>

<figure><img src="image_27.png" /><figcaption>**Scheme 5.** Equivalent simplified reaction scheme.</figcaption></figure>

$\frac{F}{RT}(E_{sq1}^0 - E_{eq}) = \xi_{sq1}^0$ and $\frac{F}{RT}(E_{sq2}^0 - E_{eq}) = \xi_{sq2}^0$ with $\frac{\xi_{sq1}^0 + \xi_{sq2}^0}{2} = \ln \sqrt{\frac{K_2}{K_1}}$. Also noting that we are here discussing

relative currents because we are only discussing the shape of the CVs to find out conditions for reversibility, the system is described by four intrinsic parameters $\Delta\xi^0 = \frac{\xi_{sq1}^0 - \xi_{sq2}^0}{2}$, $K_1$, $K_2$, and $k_1/k_2$. The general expressions of *d* and *p*/*p*<sub>max</sub> as function of these four intrinsic parameters are quite cumbersome (see the Supporting Information). However, we can make an important comment: both are dependent on thermodynamics and kinetics parameters. That directionality may depend on the kinetic parameter $k_1/k_2$ is not surprising as it is related to the rate of catalysis of both HER and HOR. Regarding reversibility, it has been suggested that a flat energy landscape is required leading to the idea that $\Delta\xi^0$ should be close to zero and $K_1$, $K_2$ close to unity to achieve reversibility. As shown in reference [17b], it is actually not the case and this transpires by the fact that $p/p_{max}$ is dependent on both thermodynamics and kinetics parameters. To illustrate this state of affair, a specific case is described in detail in the SI, discussing the interplay of thermodynamic and kinetic parameter controlling reversibility.
### 7.4. Experimental Illustration

One of the unique and most salient example of molecular catalysts exhibiting bidirectional and reversible catalysis of HER and HOR is a nickel complex belonging to the DuBois-Shaw catalyst family with an arginine derivative $[Ni(P_2^{Cy}N_2^{Arg})_2]^{6+}$ (Figure 6 and 19a).

We recently analyze the kinetic behavior of this homogeneous molecular catalyst in buffered water toward HER and HOR using cyclic voltammetry. As shown in figure 19b, bidirectional and reversible catalytic CVs are obtained under H₂ atmosphere over a wide range of pH. In the framework of the mechanism depicted in scheme 3, with M=Ni, a thorough analysis of the CVs at different scan rates, pH, catalyst and H₂ concentrations, combined with a UV/

<figure><img src="image_28.png" /><figcaption>**Figure 19.** a) Structure of $[Ni(P_2^{Cy}N_2^{Arg})_2]^{6+}$, each R–N moieties stands for a protonated arginine residue similar to that highlighted in blue. b) Background corrected catalytic responses measured at a glassy carbon electrode at 40°C and 0.1 V/s under 1 atm. H₂ at various pH values: 1 (black), 2 (red), 3 (blue), 4 (green), 5 (magenta), 6 (orange), 7 (cyan), 8 (light green), 9 (navy). Catalyst concentration: 100 µM in 0.05 M MES buffer (+0.05 M Na₂SO₄) or 0.1 M sulfuric acid solutions. c) Apparent standard potential vs. pH (Pourbaix diagram) for the PCET couples corresponding to the Ni¹⁺/Ni¹⁻ couple ($E_{sq1}^0$) in black and the NiII¹⁺/NiII¹⁻ couple in red ($E_{sq2}^0$). The dashed straight line shows the apparent standard potential of the H⁺/H₂ couple ($E_{eq}$). Reproduced with permission from ref. [17c].</figcaption></figure>

<table><caption>**Table 1:** Values determined for the 14 kinetic and thermodynamic parameters listed in Scheme 3.</caption><tr><td>E<sup>0</sup><sub>1, Hr</sub> =</td><td>0.015 V vs SHE</td><td>E<sup>0</sup><sub>1, r</sub> =</td><td>-0.370 V vs SHE</td></tr><tr><td>E<sup>0</sup><sub>2, Hr</sub> =</td><td>-0.15 V vs SHE</td><td>E<sup>0</sup><sub>2, r</sub> =</td><td>0.24 V vs SHE</td></tr><tr><td>pK<sub>H</sub><sup>II</sup> =</td><td>1</td><td>pK<sub>H</sub><sup>I</sup> =</td><td>7</td></tr><tr><td>pK<sub>H</sub><sup>IIIH</sup> =</td><td>1</td><td>pK<sub>H</sub><sup>IIIH</sup> =</td><td>7</td></tr><tr><td>K<sub>1</sub> =</td><td>2×10<sup>5</sup></td><td>K<sub>2</sub> =</td><td>20</td></tr><tr><td>k<sub>1</sub> =</td><td>10<sup>5</sup> s<sup>-1</sup></td><td>k<sub>-1</sub> =</td><td>5×10<sup>9</sup> s<sup>-1</sup></td></tr><tr><td>k<sub>2</sub> =</td><td>10 s<sup>-1</sup></td><td>k<sub>-2</sub> =</td><td>0.5 s<sup>-1</sup></td></tr></table>

Vis study allowing to determine $K_2$, allowed to evaluate all the thermodynamic and kinetic constants characterizing the system. In particular, the Pourbaix diagrams corresponding to the PCET square schemes 1 and 2 have been obtained and are shown in figure 19c. Table 1 gathers the kinetic and thermodynamic parameters determined for this system.

According to this analysis, the proton relay is well-suited because a large difference exists between $pK_H^I$ (7.2) and $pK_H^{II}$ (1) and between $pK_H^{IIIH}$ and $pK_H^{IIIH}$, so that conditions are favorable for having a bidirectional and reversible system over a large pH range as indeed observed experimentally.

Whereas $E_{sq1}^0$ is closed to $E_{eq}$, it is not the case for $E_{sq2}^0$, so that in the PCET pH range (between $pK_H^I$ and $pK_H^{II}$), $\Delta\xi^0 \approx -3$. Similarly, $K_1 = 2 \times 10^{-5} \ll 1$ and $K_2 = 20 \gg 1$ (at 1 atm H₂) are far from unity. The energy landscape is thus far from being flat. Nonetheless, reversibility is observed and, although this experimental system is not exactly similar to the study case $K_1 = K_2 = K$ discussed in the SI, similar trends are expected. In other words, proton relays allow reversibility because step 1 is intrinsically much faster than step 2, i.e., $k_1^0/k_2^0 \gg 1$. It is indeed the case because it was evaluated that $k_1 \approx 10^5 s^{-1}$ and $k_2 = 10 s^{-1}$ (considering the pseudo first order rate constant at $p_{H_2} = 1$ atm) although step 1 is uphill and step 2 downhill. If we assume a Brønsted type activation/ driving force relationship for both chemical steps 1 and 2, then $k_1^0/k_2^0 \approx 10^7$. This large value ensures that reversibility is indeed possible despite $\Delta\xi^0 \approx -3$. Taking the evaluated kinetic and thermodynamic parameters, we have $p/p_{max} = 0.6$, in line with reversible CVs (Figure 17).

This experimental example thus shows that a molecular catalyst with a transition metal center which redox couple M¹⁺/M¹ is more negative than the corresponding hydride M¹⁺H/M¹⁺H couple can be a bidirectional reversible catalyst if a suitable proton relay is installed. The first role of the proton relay functionality is to ensure fast exchange with external proton donor and acceptor and to create a high local concentration of acid/base functionality. The second role of the proton relay is to ensure that the intramolecular proton transfer is intrinsically much faster than the hydrogen formation/splitting step to balance the thermodynamical mismatch between the two redox potentials.

The observed bias toward HOR ($d \approx 0.05$) is a consequence of $K_1 = 2 \times 10^{-5} \ll 1$, $K_2 = 20 \gg 1$ and $k_1^0/k_2^0 \gg 1$ leading to $d \approx 1/K_2$. The intrinsically faster intramolecular proton transfer compared to the hydrogen formation/ splitting step also make this last step rate determining in both directions so, for pH values between $pK_H^I$ and $pK_H^{II}$, the turnover frequencies are TOF(HER) $\approx k_{-2}$ and TOF(HOR) $\approx k_2$. We finally note that slower catalysis is observed at pH >7 due to the partial breaking of the assumptions that proton transfers between the relay and the buffer are at unconditional equilibrium.

Interestingly, this analysis indicated that the step that governs both the rate of catalysis in both directions and thus the catalytic bias is the capture and release of H₂. This may be due to the fact that the outer sphere arginine residues induce too many folding constrains and prevent rapid movement of the amine pendant bases to allow for fast H₂ insertion and desorption. At the same time, this tight
packing probably prevents the formation of protonated endo isomers that have been shown to be detrimental to catalysis[55b,81,112] and ensures fast proton delivery to the pendant bases allowing to reach pure kinetic conditions.

## 8. Conclusions

While the active sites of hydrogenases have inspired the design of hundreds of biomimetic and bioinspired catalysts for H₂ evolution and, to a lesser extent, oxidation, there are very few examples for which quantitative studies highlighting the role of pendant bases as protons relays are available. In this review, we provide theoretical basis showing that the boosting effect of proton relays on catalysis is of kinetic nature, thus calling for detailed and in-depth kinetic studies, while there is no simple method allowing to diagnose the implication of a proton relay in catalysis. Importantly, for a proton relay to function under optimal conditions, its protonation, reprotonation or deprotonation steps should occur faster than the catalyzed reaction, a rule more likely obeyed in aqueous solutions. Pendant bases are also key to promote bidirectional/reversible catalysis since the frustrated Lewis pair they form with the nearby metal site is able to capture and activate H₂. We hope that the collection of systems and their analysis in this perspective article will help the community to progress towards the design of more efficient catalytic systems, not only for hydrogen technologies but also to activate other small molecules such as CO₂, N₂O, N₂, Liu, O₂ or water through proton-coupled electron transfer processes.

## Acknowledgements

This work received funding from the French National Research Agency (Labex ARCANE, CBH-EUR-GS, ANR-17-EURE-0003; SelectO₂ project; ANR-22-CE07-0045-02).

## Conflict of Interest

The authors declare no conflict of interest.

## Data Availability Statement

Data sharing is not applicable to this article as no new data were created or analyzed in this study.

**Keywords:** hydrogen • electrocatalysis • bioinspiration • hydrogenases • proton-coupled electron transfer processes

[1] G. Segev, J. Kibsgaard, C. Hahn, Z. J. Xu, W. H. Cheng, T. G. Deutsch, C. X. Xiang, J. Z. Zhang, L. Hammarstrom, D. G. Nocera, A. Z. Weber, P. Agbo, T. Hisatomi, F. E. Osterloh, K. Domen, F. F. Abdi, S. Haussener, D. J. Miller, S. Ardo, P. C. McIntyre, T. Hannappel, S. Hu, H. Atwater, J. M. Gregoire, M. Z. Ertem, I. D. Sharp, K. S. Choi, J. S. Lee, O. Ishitani, J. W. Ager, R. R. Prabhakar, A. T. Bell, S. W. Boettcher, K. Vincent, K. Takanabe, V. Artero, R. Napier, B. Roldan Cuenya, M. T. M. Koper, R. Van de Krol, F. Houle, *Journal of Physics D-Applied Physics* **2022**, 55.

[2] I. E. L. Stephens, K. Chan, A. Bagger, S. W. Boettcher, J. Bonin, E. Boutin, A. K. Buckley, R. Buonsanti, E. R. Cave, X. Chang, S. W. Chee, A. H. M. da Silva, P. de Luna, O. Einsle, B. Endrödi, M. Escudero-Escribano, J. V. Ferreira de Araujo, M. C. Figueiredo, C. Hahn, K. U. Hansen, S. Haussener, S. Hunegnaw, Z. Huo, Y. J. Hwang, C. Janáky, B. S. Jayathilake, F. Jiao, Z. P. Jovanov, P. Karimi, M. T. M. Koper, K. P. Kuhl, W. H. Lee, Z. Liang, X. Liu, S. Ma, M. Ma, H.-S. Oh, M. Robert, B. R. Cuenya, J. Rossmeisl, C. Roy, M. P. Ryan, E. H. Sargent, P. Sebastián-Pascual, B. Seger, L. Steier, P. Strasser, A. S. Varela, R. E. Vos, X. Wang, B. Xu, H. Yadegari, Y. Zhou, *Journal of Physics: Energy* **2022**, 4, 042003.

[3] S. Z. Andersen, V. Čolić, S. Yang, J. A. Schwalbe, A. C. Nielander, J. M. McEnaney, K. Enemark-Rasmussen, J. G. Baker, A. R. Singh, B. A. Rohr, M. J. Statt, S. J. Blair, S. Mezzavilla, J. Kibsgaard, P. C. K. Vesborg, M. Cargnello, S. F. Bent, T. F. Jaramillo, I. E. L. Stephens, J. K. Nørskov, I. Chorkendorff, *Nature* **2019**, 570, 504–508.

[4] a) T.-Y. Chen, Y. W. Hsiao, M. Baker-Fales, F. Cameli, P. Dimitrakellis, D. G. Vlachos, *Chem. Sci.* **2022**, 13, 10644–10685; b) D. S. Mallapragada, Y. Dvorkin, M. A. Modestino, D. V. Esposito, W. A. Smith, B.-M. Hodge, M. P. Harold, V. M. Donnelly, A. Nuz, C. Bloomquist, K. Baker, L. C. Grabow, Y. Yan, N. N. Rajput, R. L. Hartman, E. J. Biddinger, E. S. Aydil, A. D. Taylor, *Joule* **2023**, 7, 23–41.

[5] V. Artero, *Nat. Energy* **2017**, 2, 17131.

[6] B. C. H. Steele, A. Heinzel, *Nature* **2001**, 414, 345–352.

[7] W. Lubitz, H. Ogata, O. Rüdiger, E. Reijerse, *Chem. Rev.* **2014**, 114, 4081–4148.

[8] a) F. A. Armstrong, R. M. Evans, S. V. Hexter, B. J. Murphy, M. M. Roessler, P. Wulff, *Acc. Chem. Res.* **2016**, 49, 884–892;

b) M. del Barrio, M. Sensi, C. Orain, C. Baffert, S. Dementin, V. Fourmond, C. Léger, *Acc. Chem. Res.* **2018**, 51, 769–777;

c) R. M. Evans, B. Siritanaratkul, C. F. Megarity, K. Pandey, T. F. Esterle, S. Badiani, F. A. Armstrong, *Chem. Soc. Rev.* **2019**, 48, 2039–2052.

[9] D. W. Mulder, J. W. Peters, S. Raugei, *Chem. Commun.* **2021**, 57, 713–720.

[10] a) J. N. Butt, M. Filipiak, W. R. Hagen, *Eur. J. Biochem.* **1997**, 245, 116–122; b) M. Winkler, J. F. Duan, A. Rutz, C. Felbek, L. Scholtysek, O. Lampret, J. Jaenecke, U. P. Apfel, G. Gilardi, F. Valetti, V. Fourmond, E. Hofmann, C. Leger, T. Happe, *Nat. Commun.* **2021**, 12; c) A. Parkin, C. Cavazza, J. C. Fontecilla-Camps, F. A. Armstrong, *J. Am. Chem. Soc.* **2006**, 128, 16808–16815; d) C. Madden, M. D. Vaughn, I. Diez-Perez, K. A. Brown, P. W. King, D. Gust, A. L. Moore, T. A. Moore, *J. Am. Chem. Soc.* **2012**, 134, 1577–1582; e) K. Pandey, S. T. A. Islam, T. Happe, F. A. Armstrong, *Proc. Natl. Acad. Sci. USA* **2017**, 114, 3843–3848; f) P. Ceccaldi, K. Schuchmann, V. Muller, S. J. Elliott, *Energy Environ. Sci.* **2017**, 10, 503–508.

[11] A. Fasano, H. Land, V. Fourmond, G. Berggren, C. Leger, *J. Am. Chem. Soc.* **2021**, 143, 20320–20325.

[12] J. H. Artz, O. A. Zadvornyy, D. W. Mulder, S. M. Keable, A. E. Cohen, M. W. Ratzloff, S. G. Williams, B. Ginovska, N. Kumar, J. Song, S. E. McPhillips, C. M. Davidson, A. Y. Lyubimov, N. Pence, G. J. Schut, A. K. Jones, S. M. Soltis, M. W. W. Adams, S. Raugei, P. W. King, J. W. Peters, *J. Am. Chem. Soc.* **2020**, 142, 1227–1235.

[13] a) J. Riethausen, O. Rudiger, W. Gartner, W. Lubitz, H. S. Shafaat, *ChemBioChem* **2013**, 14, 1714–1719; b) P. Ceccaldi, J. M. Gregoire, M. Z. Ertem, I. D. Sharp, K. S. Choi, J. S. Lee, O. Ishitani, J. W. Ager, R. R. Prabhakar, A. T. Bell, S. W. Boettcher, K. Vincent, K. Takanabe, V. Artero, R. Napier, B. Roldan Cuenya, M. T. M. Koper, R. Van de Krol, F. Houle, *Journal of Physics D-Applied Physics* **2022**, 55.
M. C. Marques, V. Fourmond, I. C. Pereira, C. Leger, *Chem. Commun.* **2015**, *51*, 14223–14226; c) B. J. Murphy, F. Sargent, F. A. Armstrong, *Energy Environ. Sci.* **2014**, *7*, 1426–1433; d) S. W. Ragsdale, *J. Biol. Chem.* **2009**, *284*, 18571–18575; e) A. Abou Hamdan, S. Dementin, P. P. Liebgott, O. Gutierrez-Sanz, P. Richaud, A. L. De Lacey, M. Roussett, P. Bertrand, L. Cournac, C. Leger, *J. Am. Chem. Soc.* **2012**, *134*, 8368–8371; f) A. A. Karyakin, S. V. Morozov, O. G. Voronin, N. A. Zorin, E. E. Karyakina, V. N. Fateyev, S. Cosnier, *Angew. Chem. Int. Ed.* **2007**, *46*, 7244–7246; g) G. Goldet, A. F. Wait, J. A. Cracknell, K. A. Vincent, M. Ludwig, O. Lenz, B. Friedrich, F. A. Armstrong, *J. Am. Chem. Soc.* **2008**, *130*, 11106–11113; h) H. R. Pershad, J. L. C. Duff, H. A. Heering, E. C. Duin, S. P. J. Albracht, F. A. Armstrong, *Biochemistry* **1999**, *38*, 8992–8999.

[14] a) A. A. Karyakin, S. V. Morozov, E. E. Karyakina, S. D. Varfolomeyev, N. A. Zorin, S. Cosnier, *Electrochem. Commun.* **2002**, *4*, 417–420; b) S. V. Morozov, E. E. Karyakina, O. A. Zadvornyi, N. A. Zorin, S. D. Varfolomeev, A. A. Karyakin, *Russ. J. Electrochem.* **2002**, *38*, 97–102.

[15] A. Fasano, C. Guendon, A. Jacq-Bailley, A. Kpebe, J. Wozniak, C. Baffert, M. d Barrio, V. Fourmond, M. Brugna, C. Léger, *J. Am. Chem. Soc.* **2023**.

[16] a) S. V. Hexter, F. Grey, T. Happe, V. Climent, F. A. Armstrong, *Proc. Natl. Acad. Sci. USA* **2012**, *109*, 11516–11521; b) J. M. Saveant, *ACS Catal.* **2018**, *8*, 7608–7611.

[17] a) S. V. Hexter, T. F. Esterle, F. A. Armstrong, *Phys. Chem. Chem. Phys.* **2014**, *16*, 11822–11833; b) V. Fourmond, E. S. Wiedner, W. J. Shaw, C. Leger, *J. Am. Chem. Soc.* **2019**, *141*, 11269–11285; c) B. Reuillard, C. Costentin, V. Artero, *Angew. Chem. Int. Ed.* **2023**, *62*, d) A. Fasano, H. Land, V. Fourmond, G. Berggren, C. Léger, *J. Am. Chem. Soc.* **2021**, *143*, 20320–20325; e) A. Fasano, C. Baffert, C. Schumann, G. Berggren, J. A. Birrell, V. Fourmond, C. Leger, *J. Am. Chem. Soc.* **2024**, *146*, 1455–1466; f) A. Fasano, V. Fourmond, C. Léger, *Bioelectrochemistry* **2024**, *155*.

[18] V. Fourmond, C. Baffert, K. Sybirna, S. Dementin, A. Abou-Hamdan, I. Meynial-Salles, P. Soucaille, H. Bottin, C. Leger, *Chem. Commun.* **2013**, *49*, 6840–6842.

[19] a) A. Volbeda, M. H. Charon, C. Piras, E. C. Hatchikian, M. Frey, J. C. Fontecilla-Camps, *Nature* **1995**, *373*, 580–587; b) A. Volbeda, E. Garcin, C. Piras, A. L. de Lacey, V. M. Fernandez, E. C. Hatchikian, M. Frey, J. C. Fontecilla-Camps, *J. Am. Chem. Soc.* **1996**, *118*, 12989–12996; c) E. Garcin, X. Vernede, E. C. Hatchikian, A. Volbeda, M. Frey, J. C. Fontecilla-Camps, *Structure* **1999**, *7*, 557–566; d) Y. Nicolet, A. L. de Lacey, X. Vernede, V. M. Fernandez, E. C. Hatchikian, J. C. Fontecilla-Camps, *J. Am. Chem. Soc.* **2001**, *123*, 1596–1601.

[20] J. W. Peters, W. N. Lanzilotta, B. J. Lemon, L. C. Seefeldt, *Science* **1998**, *282*, 1853–1858.

[21] Y. Higuchi, T. Yagi, N. Yasuoka, *Structure* **1997**, *5*, 1671–1680.

[22] H. J. Fan, M. B. Hall, *J. Am. Chem. Soc.* **2001**, *123*, 3828–3829.

[23] A. Silakov, B. Wenk, E. Reijerse, W. Lubitz, *Phys. Chem. Chem. Phys.* **2009**, *11*, 6592–6599.

[24] G. Berggren, A. Adamska, C. Lambertz, T. R. Simmons, J. Esselborn, M. Atta, S. Gambarelli, J. M. Mouesca, E. Reijerse, W. Lubitz, T. Happe, V. Artero, M. Fontecave, *Nature* **2013**, *499*, 66–69.

[25] M. Teixeira, G. Fauque, I. Moura, P. A. Lespinat, Y. Berlier, B. Prickril, H. D. Peck Jr, A. V. Xavier, J. Le Gall, J. J. G. Moura, *Eur. J. Biochem.* **1987**, *167*, 47–58.

[26] S. Niu, N. B. Hall, *Inorg. Chem.* **2001**, *40*, 6201–6203.

[27] R. M. Evans, N. Krahn, J. Weiss, K. A. Vincent, D. Söll, F. A. Armstrong, *J. Am. Chem. Soc.* **2024**, *146*, 16971–16976.

[28] a) R. M. Evans, S. E. Beaton, P. Rodriguez Macia, Y. Pang, K. L. Wong, L. Kertess, W. K. Myers, R. Björnsson, P. A. Ash, K. A. Vincent, S. B. Carr, F. A. Armstrong, *Chem. Sci.* **2023**, *14*, 8531–8551; b) R. M. Evans, E. J. Brooke, S. A. Wehlin, E. Nomerotskaia, F. Sargent, S. B. Carr, S. E. Phillips, F. A. Armstrong, *Nat. Chem. Biol.* **2016**, *12*, 46–50.

[29] R. Custelcean, J. E. Jackson, *Chem. Rev.* **2001**, *101*, 1963–1980.

[30] C. Baffert, M. Demuez, L. Cournac, B. Burlat, B. Guigliarelli, P. Bertrand, L. Girbal, C. Leger, *Angew. Chem. Int. Ed.* **2008**, *47*, 2052–2054.

[31] M. J. Lukey, A. Parkin, M. M. Roessler, B. J. Murphy, J. Harmer, T. Palmer, F. Sargent, F. A. Armstrong, *J. Biol. Chem.* **2010**, *285*, 3928–3938.

[32] a) S. Canaguier, V. Artero, M. Fontecave, *Dalton Trans.* **2008**, *39*, 315–325; b) T. R. Simmons, G. Berggren, M. Bacchi, M. Fontecave, V. Artero, *Coord. Chem. Rev.* **2014**, 270–271, 127–150; c) G. Caserta, S. Roy, M. Atta, V. Artero, M. Fontecave, *Curr. Opin. Chem. Biol.* **2015**, *25*, 36–47; d) J. T. Kleinhaus, F. Wittkamp, S. Yadav, D. Siegmund, U.-P. Apfel, *Chem. Soc. Rev.* **2021**, *50*, 1668–1784.

[33] K. Weber, T. Krämer, H. S. Shafaat, T. Weyhermüller, E. Bill, M. van Gastel, F. Neese, W. Lubitz, *J. Am. Chem. Soc.* **2012**, *134*, 20745–20755.

[34] J. C. Jeffrey, T. B. Rauchfuss, *Inorg. Chem.* **1979**, *18*, 2658–2666.

[35] a) D. Brazzolotto, L. Wang, H. Tang, M. Gennari, N. Queyriaux, C. Philouze, S. Demeshko, F. Meyer, M. Orio, V. Artero, M. B. Hall, C. Duboc, *ACS Catal.* **2018**, *8*, 10658–10667; b) H. Tang, M. B. Hall, *J. Am. Chem. Soc.* **2017**, *139*, 18065–18070; c) S. Ding, P. Ghosh, M. Y. Darensbourg, M. B. Hall, *Proc. Natl. Acad. Sci. USA* **2017**, *114*, E9775–E9782; d) S. Ding, P. Ghosh, A. M. Lunsford, N. Wang, N. Bhuvanesh, M. B. Hall, M. Y. Darensbourg, *J. Am. Chem. Soc.* **2016**, *138*, 12920–12927; e) D. Brazzolotto, M. Gennari, N. Queyriaux, T. R. Simmons, J. Pécaut, S. Demeshko, F. Meyer, M. Orio, V. Artero, C. Duboc, *Nat. Chem.* **2016**, *8*, 1054–1060; f) P. Ghosh, S. Ding, R. B. Chupik, M. Quiroz, C.-H. Hsieh, N. Bhuvanesh, M. B. Hall, M. Y. Darensbourg, *Chem. Sci.* **2017**, *8*, 8291–8300; g) P. Ghosh, M. Quiroz, N. Wang, N. Bhuvanesh, M. Y. Darensbourg, *Dalton Trans.* **2017**, *46*, 5617–5624.

[36] A. Aster, S. H. Wang, M. Mirmohades, C. Esmieu, G. Berggren, L. Hammarstrom, R. Lomoth, *Chem. Sci.* **2019**, *10*, 5582–5588.

[37] M. E. Carroll, B. E. Barton, T. B. Rauchfuss, P. J. Carroll, *J. Am. Chem. Soc.* **2012**, *134*, 18843–18852.

[38] M. E. Ahmed, S. Dey, B. Mondal, A. Dey, *Chem. Commun.* **2017**, *53*, 8188–8191.

[39] M. E. Ahmed, A. Nayek, A. Krizan, N. Coutard, A. Morozan, S. G. Dey, R. Lomoth, L. Hammarstrom, V. Artero, A. Dey, *J. Am. Chem. Soc.* **2022**, *144*, 3614–3625.

[40] E. S. Wiedner, A. M. Appel, S. Raugei, W. J. Shaw, R. M. Bullock, *Chem. Rev.* **2022**, *122*, 12427–12474.

[41] a) C. J. Curtis, A. Miedaner, R. Ciancanelli, W. W. Ellis, B. C. Noll, M. R. DuBois, D. L. DuBois, *Inorg. Chem.* **2003**, *42*, 216–227; b) A. D. Wilson, R. H. Newell, M. J. McNevin, J. T. Muckerman, M. R. DuBois, D. L. DuBois, *J. Am. Chem. Soc.* **2006**, *128*, 358–366; c) U. J. Kilgore, J. A. S. Roberts, D. H. Pool, A. M. Appel, M. P. Stewart, M. R. DuBois, W. G. Dougherty, W. S. Kassel, R. M. Bullock, D. L. DuBois, *J. Am. Chem. Soc.* **2011**, *133*, 5861–5872.

[42] V. Artero, J.-M. Saveant, *Energy Environ. Sci.* **2014**, *7*, 3808–3814.

[43] a) E. S. Wiedner, J. Y. Yang, W. G. Dougherty, W. S. Kassel, R. M. Bullock, M. R. DuBois, D. L. DuBois, *Organometallics* **2010**, *29*, 5390–5401; b) M. Fang, E. S. Wiedner, W. G.
Dougherty, W. S. Kassel, T. Liu, D. L. DuBois, R. M. Bullock, *Organometallics* **2014**, *33*, 5820–5833; c) G. M. Jacobsen, J. Y. Yang, B. Twamley, A. D. Wilson, R. M. Bullock, M. R. DuBois, D. L. DuBois, *Energy Environ. Sci.* **2008**, *1*, 167–174.

[44] a) E. S. Wiedner, A. M. Appel, D. L. DuBois, R. M. Bullock, *Inorg. Chem.* **2013**, *52*, 14391–14403; b) C. M. Klug, W. G. Dougherty, W. S. Kassel, E. S. Wiedner, *Organometallics* **2019**, *38*, 1269–1279.

[45] a) M. L. Helm, M. P. Stewart, R. M. Bullock, M. R. DuBois, D. L. DuBois, *Science* **2011**, *333*, 863–866; b) R. M. Stolley, J. M. Darmon, P. Das, M. L. Helm, *Organometallics* **2016**, *35*, 2965–2974.

[46] a) A. Zarkadoulas, M. J. Field, V. Artero, C. A. Mitsopoulou, *ChemCatChem* **2017**, *9*, 2308–2317; b) A. Zarkadoulas, M. J. Field, C. Papatriantafyllopoulou, J. Fize, V. Artero, C. A. Mitsopoulou, *Inorg. Chem.* **2016**, *55*, 432–444; c) Y. Aimoto, K. Koshiba, K. Yamauchi, K. Sakai, *Chem. Commun.* **2018**, *54*, 12820–12823.

[47] C. P. Yap, K. Hou, A. A. Bengali, W. Y. Fan, *Inorg. Chem.* **2017**, *56*, 10926–10931.

[48] a) W. R. McNamara, Z. Han, P. J. Alperin, W. W. Brennessel, P. L. Holland, R. Eisenberg, *J. Am. Chem. Soc.* **2011**, *133*, 15368–15371; b) W. R. McNamara, Z. Han, C.-J. Yin, W. W. Brennessel, P. L. Holland, R. Eisenberg, *Proc. Natl. Acad. Sci. USA* **2012**, *109*, 15594–15599; c) C. A. Downes, J. W. Yoo, N. M. Orchanian, R. Haiges, S. C. Marinescu, *Chem. Commun.* **2017**, *53*, 7306–7309; d) T. Fogeron, J.-P. Porcher, M. Gomez-Mingot, T. K. Todorova, L.-M. Chamoreau, C. Mellot-Draznieks, Y. Li, M. Fontecave, *Dalton Trans.* **2016**, *45*, 14754–14763; e) C. S. Letko, J. A. Panetier, M. Head-Gordon, T. D. Tilley, *J. Am. Chem. Soc.* **2014**, *136*, 9364–9376.

[49] B. H. Solis, S. Hammes-Schiffer, *J. Am. Chem. Soc.* **2012**, *134*, 15253–15256.

[50] A. J. Clough, J. W. Yoo, M. H. Mecklenburg, S. C. Marinescu, *J. Am. Chem. Soc.* **2015**, *137*, 118–121.

[51] M. Gomez-Mingot, J.-P. Porcher, T. K. Todorova, T. Fogeron, C. Mellot-Draznieks, Y. Li, M. Fontecave, *J. Phys. Chem. B* **2015**, *119*, 13524–13533.

[52] C. F. Leung, Y. Z. Chen, H. Q. Yu, S. M. Yiu, C. C. Ko, T. C. Lau, *Int. J. Hydrogen Energy* **2011**, *36*, 11640–11645.

[53] a) S. Grau, M. Schilling, D. Moonshiram, J. Benet-Buchholz, S. Luber, A. Llobet, C. Gimbert-Surinach, *ChemSusChem* **2020**, *13*, 2745–2752; b) C. C. L. McCrory, C. Uyeda, J. C. Peters, *J. Am. Chem. Soc.* **2012**, *134*, 3164–3170; c) C.-B. Li, A. J. Bagnall, D. Sun, J. Rendon, M. Koepf, S. Gambarelli, J.-M. Mouesca, M. Chavarot-Kerlidou, V. Artero, *Sustain. Energy Fuels* **2022**, *6*, 143–149.

[54] M. Willery, P.-G. Julliard, F. Molton, F. Thomas, J. Fortage, C. Costentin, M.-N. Collomb, *ACS Catal.* **2024**, *11352–11365*.

[55] a) A. Dutta, S. Lense, J. Hou, M. H. Engelhard, J. A. S. Roberts, W. J. Shaw, *J. Am. Chem. Soc.* **2013**, *135*, 18490–18496; b) E. S. Rountree, J. L. Dempsey, *J. Am. Chem. Soc.* **2015**, *137*, 13371–13380; c) D. Dolui, S. Khandelwal, P. Majumder, A. Dutta, *Chem. Commun.* **2020**, *56*, 8166–8181; d) N. Kaeffer, M. Chavarot-Kerlidou, V. Artero, *Acc. Chem. Res.* **2015**, *48*, 1286–1295.

[56] P.-A. Jacques, V. Artero, J. Pécaut, M. Fontecave, *Proc. Natl. Acad. Sci. USA* **2009**, *106*, 20627–20632.

[57] A. Bhattacharjee, E. S. Andreiadis, M. Chavarot-Kerlidou, M. Fontecave, M. J. Field, V. Artero, *Chem. Eur. J.* **2013**, *19*, 15166–15174.

[58] D. Sun, A. K. Harshan, J. Pecaut, S. Hammes-Schiffer, C. Costentin, V. Artero, *ChemElectroChem* **2021**, *8*, 2671–2679.

[59] N. Queyriaux, D. Y. Sun, J. Fize, J. Pecaut, M. J. Field, M. Chavarot-Kerlidou, V. Artero, *J. Am. Chem. Soc.* **2020**, *142*, 274–282.

[60] F. J. Arriaza-Gallardo, Y.-C. Zheng, M. Gehl, S. Nomura, J. P. Fernandes-Queiroz, S. Shima, *ChemBioChem* **2023**, *24*, e202300330.

[61] S. K. Mandal, C. Sunil, J. Choudhury, *ACS Catal.* **2024**, *14*, 2058–2070.

[62] A. Gerli, M. Sabat, L. G. Marzilli, *J. Am. Chem. Soc.* **1992**, *114*, 6711–6718.

[63] P. Huo, C. Uyeda, J. D. Goodpaster, J. C. Peters, T. F. Miller, III, *ACS Catal.* **2016**, *6*, 6114–6123.

[64] G. K. Rao, M. P. Jamshidi, J. I. G. Dawkins, W. Pell, I. Korobkov, D. Richeson, *Dalton Trans.* **2017**, *46*, 6518–6522.

[65] a) Z. J. Han, L. X. Shen, W. W. Brennessel, P. L. Holland, R. Eisenberg, *J. Am. Chem. Soc.* **2013**, *135*, 14659–14669; b) Z. J. Han, W. R. McNamara, M. S. Eum, P. L. Holland, R. Eisenberg, *Angew. Chem. Int. Ed.* **2012**, *51*, 1667–1670.

[66] a) J. Xie, Q. Zhou, C. Li, W. Wang, Y. Hou, B. Zhang, X. Wang, *Chem. Commun.* **2014**, *50*, 6520–6522; b) E. Deponti, A. Luisa, M. Natali, E. Iengo, F. Scandola, *Dalton Trans.* **2014**, *43*, 16345–16353; c) N. Queyriaux, R. T. Jane, J. Massin, V. Artero, M. Chavarot-Kerlidou, *Coord. Chem. Rev.* **2015**, *304–305*, 3–19; d) C. Bachmann, M. Guttentag, B. Spingler, R. Alberto, *Inorg. Chem.* **2013**, *52*, 6055–6061; e) M. B. Brands, J. N. H. Reek, *Inorg. Chem.* **2024**, *63*, 8484–8492.

[67] a) L. Chen, M. Wang, K. Han, P. Zhang, F. Gloaguen, L. Sun, *Energy Environ. Sci.* **2014**, *7*, 329–334; b) Y.-Q. Zhang, R.-Z. Liao, *Phys. Chem. Chem. Phys.* **2017**, *19*, 32589–32596.

[68] R. Tatematsu, T. Inomata, T. Ozawa, H. Masuda, *Angew. Chem. Int. Ed. Engl.* **2016**, *55*, 5247–5250.

[69] J. Amtawong, C. L. Montgomery, G. P. Bein, A. L. Raithel, T. W. Hamann, C.-H. Chen, J. L. Dempsey, *J. Am. Chem. Soc.* **2024**, *146*, 3742–3750.

[70] C. H. Lee, D. K. Dogutan, D. G. Nocera, *J. Am. Chem. Soc.* **2011**, *133*, 8775–8777.

[71] B. H. Solis, A. G. Maher, T. Honda, D. C. Powers, D. G. Nocera, S. Hammes-Schiffer, *ACS Catal.* **2014**, *4*, 4516–4526.

[72] D. J. Graham, D. G. Nocera, *Organometallics* **2014**, *33*, 4994–5001.

[73] A. Rana, B. Mondal, P. Sen, S. Dey, A. Dey, *Inorg. Chem.* **2017**, *56*, 1783–1793.

[74] S. Bhunia, A. Rana, S. Hematian, K. D. Karlin, A. Dey, *Inorg. Chem.* **2021**, *60*, 13876–13887.

[75] a) K. W. Kramarz, J. R. Norton, *Progress in Inorganic Chemistry*, Vol 42 **1994**, *42*, 1–65; b) K. R. Brereton, N. E. Smith, N. Hazari, A. J. M. Miller, *Chem. Soc. Rev.* **2020**, *49*, 7929–7948.

[76] J. M. Saveant, *Angew. Chem. Int. Ed.* **2019**, *58*, 2125–2128.

[77] C. L. Montgomery, J. Amtawong, A. M. Jordan, D. A. Kurtz, J. L. Dempsey, *Chem. Soc. Rev.* **2023**, *52*, 7137–7169.

[78] D. A. Kurtz, D. Dhar, N. Elgrishi, B. Kandemir, S. F. McWilliams, W. C. Howland, C.-H. Chen, J. L. Dempsey, *J. Am. Chem. Soc.* **2021**, *143*, 3393–3406.

[79] M. M. Roubelakis, D. K. Bediako, D. K. Dogutan, D. G. Nocera, *Energy Environ. Sci.* **2012**, *5*, 7737–7740.

[80] R. Sun, M. Liu, S.-L. Zheng, D. K. Dogutan, C. Costentin, D. G. Nocera, *Proc. Natl. Acad. Sci. USA* **2022**, *119*, e2122063119.

[81] E. S. Rountree, B. D. McCarthy, J. L. Dempsey, *Inorg. Chem.* **2019**, *58*, 6647–6658.

[82] E. S. Rountree, D. J. Martin, B. D. McCarthy, J. L. Dempsey, *ACS Catal.* **2016**, *6*, 3326–3335.

[83] A similar behavior has been reported for the acetonitrile derivative in reference [54].

[84] A. M. Appel, D. L. DuBois, M. R. DuBois, *J. Am. Chem. Soc.* **2005**, *127*, 12717–12726.

[85] No consensus is reached yet regarding the protonation state of the amine ligand in the Co(II) state as chloride axial ligand
exchange at the Co(III)/Co(II) redox wave being a possible explanation for its shift in the presence of acid.[54]

[86] M.-H. Ho, R. Rousseau, J. A. S. Roberts, E. S. Wiedner, M. Dupuis, D. L. DuBois, R. M. Bullock, S. Raugei, ACS Catal. **2015**, *5*, 5436–5452.

[87] V. Fourmond, P. A. Jacques, M. Fontecave, V. Artero, Inorg. Chem. **2010**, *49*, 10338–10347.

[88] B. H. Solis, Y. Yu, S. Hammes-Schiffer, Inorg. Chem. **2013**, *52*, 6994–6999.

[89] J. M. Saveant, Chem. Rev. **2008**, *108*, 2348–2378.

[90] N. G. Léonard, R. Dhaoui, T. Chantarojsiri, J. Y. Yang, ACS Catal. **2021**, *11*, 10923–10932.

[91] a) T. B. Liu, S. T. Chen, M. J. O'Hagan, M. R. DuBois, R. M. Bullock, D. L. DuBois, *J. Am. Chem. Soc.* **2012**, *134*, 6257–6272; b) T. Liu, D. L. DuBois, R. M. Bullock, *Nat. Chem.* **2013**, *5*, 228–233.

[92] a) E. B. Hulley, M. L. Helm, R. M. Bullock, *Chem. Sci.* **2014**, *5*, 4729–4741; b) E. B. Hulley, K. D. Welch, A. M. Appel, D. L. DuBois, R. M. Bullock, *J. Am. Chem. Soc.* **2013**, *135*, 11736–11739.

[93] S. Zhang, A. M. Appel, R. M. Bullock, *J. Am. Chem. Soc.* **2017**.

[94] T. Liu, X. Wang, C. Hoffmann, D. L. DuBois, R. M. Bullock, *Angew. Chem. Int. Ed. Engl.* **2014**, *53*, 5300–5304.

[95] S. Raugei, S. T. Chen, M. H. Ho, B. Ginovska-Pangovska, R. J. Rousseau, M. Dupuis, D. L. DuBois, R. M. Bullock, *Chem. Eur. J.* **2012**, *18*, 6493–6506.

[96] A. D. Wilson, R. K. Shoemaker, A. Miedaner, J. T. Muckerman, D. L. DuBois, M. R. DuBois, *Proc. Natl. Acad. Sci. USA* **2007**, *104*, 6951–6956.

[97] a) N. Kumar, J. M. Darmon, C. J. Weiss, M. L. Helm, S. Raugei, R. M. Bullock, *Organometallics* **2019**, *38*, 1391–1396; b) T. B. Liu, Q. Liao, M. O'Hagan, E. B. Hulley, D. L. DuBois, R. M. Bullock, *Organometallics* **2015**, *34*, 2747–2764.

[98] T. Liu, M. R. DuBois, D. L. DuBois, R. M. Bullock, *Energy Environ. Sci.* **2014**, *7*, 3630–3639.

[99] E. B. Hulley, N. Kumar, S. Raugei, R. M. Bullock, ACS Catal. **2015**, *5*, 6838–6847.

[100] J. M. Camara, T. B. Rauchfuss, Nat. Chem. **2012**, *4*, 26–30.

[101] N. Wang, M. Wang, Y. Wang, D. Zheng, H. Han, M. S. G. Ahlquist, L. Sun, *J. Am. Chem. Soc.* **2013**, *135*, 13688–13691.

[102] T. F. Liu, M. Y. Guo, A. Orthaber, R. Lomoth, M. Lundberg, S. Ott, L. Hammarström, Nat. Chem. **2018**, *10*, 881–887.

[103] A. Le Goff, V. Artero, B. Jousselme, P. D. Tran, N. Guillet, R. Metaye, A. Fihri, S. Palacin, M. Fontecave, *Science* **2009**, *326*, 1384–1387.

[104] P. D. Tran, A. Le Goff, J. Heidkamp, B. Jousselme, N. Guillet, S. Palacin, H. Dau, M. Fontecave, V. Artero, *Angew. Chem. Int. Ed.* **2011**, *50*, 1371–1374.

[105] S. E. Smith, J. Y. Yang, D. L. DuBois, R. M. Bullock, *Angew. Chem. Int. Ed.* **2012**, *51*, 3152–3155.

[106] a) A. Dutta, D. L. DuBois, J. A. S. Roberts, W. J. Shaw, *Proc. Natl. Acad. Sci. USA* **2014**, *111*, 16286–16291; b) A. Dutta, J. A. S. Roberts, W. J. Shaw, *Angew. Chem. Int. Ed.* **2014**, *53*, 6487–6491.

[107] P. Rodriguez-Macia, A. Dutta, W. Lubitz, W. J. Shaw, O. Rudiger, *Angew. Chem. Int. Ed.* **2015**, *54*, 12303–12307.

[108] a) N. Coutard, B. Reuillard, T. N. Huan, F. Valentino, R. T. Jane, S. Gentil, E. S. Andreiadis, A. Le Goff, T. Asset, F. Maillard, B. Jousselme, A. Morozan, S. Lyonnard, V. Artero, P. Chenevier, *Chem Catalysis* **2021**, *1*, 88–105; b) S. Gentil, N. Lalaoui, A. Dutta, Y. Nedellec, S. Cosnier, W. J. Shaw, V. Artero, A. Le Goff, *Angew. Chem. Int. Ed.* **2017**, *56*, 1845–1849; c) S. Gentil, J. K. Molloy, M. Carriere, A. Hobballah, A. Dutta, S. Cosnier, W. J. Shaw, G. Gellon, C. Belle, V. Artero, F. Thomas, A. Le Goff, *Joule* **2019**, *3*, 2020–2029.

[109] P. D. Tran, A. Morozan, S. Archambault, J. Heidkamp, P. Chenevier, H. Dau, M. Fontecave, A. Martinent, B. Jousselme, V. Artero, *Chem. Sci.* **2015**, *6*, 2050–2053.

[110] J. Schild, B. Reuillard, A. Morozan, P. Chenevier, E. Gravel, E. Doris, V. Artero, *J. Am. Chem. Soc.* **2021**, *143*, 18150–18158.

[111] A. Fasano, V. Fourmond, C. Léger, *Curr. Opin. Electrochem.* **2024**, *101489*.

[112] E. S. Wiedner, H. J. S. Brown, M. L. Helm, *J. Am. Chem. Soc.* **2016**, *138*, 604–616.

Manuscript received: July 23, 2024

Version of record online: November 18, 2024