Reaction Pathways of Hydrogen-Evolving Electrocatalysts

Perspective pubs.acs.org/acscatalysis

Reaction Pathways of Hydrogen-Evolving Electrocatalysts: Electrochemical and Spectroscopic Studies of Proton-Coupled Electron Transfer Processes Noémie Elgrishi,‡ Brian D. McCarthy,‡ Eric S. Rountree, and Jillian L. Dempsey* Department of Chemistry, University of North Carolina, Chapel Hill, North Carolina 27599-3290, United States ABSTRACT: Proton-coupled electron transfer (PCET) reactions are at the heart of the catalytic processes involved in hydrogen evolution. In this Perspective, the state-of-the-art spectroscopic and electrochemical methods available to elucidate the mechanisms of PCET reactions of fuel-forming catalysts are presented. Through examples of our recent work, the applications of these methods are discussed with a focus on the type of information and the accuracy that can be obtained from each. Three case studies are presented to illustrate different possible origins for peak shifts observed in cyclic voltammograms. KEYWORDS: PCET, electrocatalysis, cyclic voltammetry, stopped-flow, solar fuels

1. INTRODUCTION Development of cost-effective solar energy storage technology is paramount for realizing the transition away from fossil fuel energy.1−3 In terms of chemical fuels, chemical bonds in energypoor molecules must be rearranged to yield energy-rich fuels. Frequently, these rearrangements involve the movement of multiple protons and electrons such as in H2 evolution, CO2 reduction, and water oxidation.4−6 Minimizing energy loss during these rearrangements is imperative for practical largescale fuel production. As energy losses are intricately related to the kinetics of elementary chemical steps, careful modulation of the pathway taken from energy-poor reactants to energy-rich products is critical. It is particularly attractive to mediate these complex transformations using molecular transition metal catalysts. As such, opportunities exist to modulate reaction mechanism and perturb reaction kinetics through systematic changes to catalyst structure and electronics, as well as reaction conditions. A common motivation for understanding reaction mechanisms of catalytic transformations is the hope that this insight can lead to modification and improvement of a catalyst. For example, if the hydrogen evolution mechanism for a specific catalyst is understood, a second generation catalyst can be rationally designed. An elegant example is the evolution and improvement of the [Ni(PR2 NR2 ′)2]2+ family of hydrogen evolution catalysts.7 Notably, within the context of hydrogen evolution, the catalyst orchestrates the union of two electrons and two protons, and elementary steps are frequently proton-coupled electron transfer (PCET) reactions. PCET is traditionally viewed in terms of a square scheme (Scheme 1) where concerted and stepwise routes are possible. Reactions that proceed via the stepwise routes © XXXX American Chemical Society

Scheme 1. Square Scheme Representation of a Homogeneous PCET Reaction Where Species M Receives One Electron and One Proton To Form MHa

a

H−A represents an acid molecule, and ET, PT, and CPET indicate electron transfer, proton transfer, and concerted proton−electron transfer, respectively.

necessarily proceed through high-energy intermediates. Consequently, mechanistic study of PCET reactions involving transition-metal complexes is an important step toward designing more efficient catalysts for hydrogen evolution reactions. Further, as PCET reactions are additionally relevant outside the subfield of fuel formation reactions, the quest to understand PCET in hydrogen evolution catalysts simultaneously contributes more broadly useful fundamental knowledge. Our group has focused on understanding PCET processes in matters relevant to energy conversion and storage processes; specifically, in the context of homogeneous electrochemical catalysts for the hydrogen evolution reaction (HER). As noted above, it has become increasingly clear that an understanding of Received: March 16, 2016 Revised: April 21, 2016

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DOI: 10.1021/acscatal.6b00778 ACS Catal. 2016, 6, 3644−3659

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reaction for a CPET process (Figure 1A). Although the formation of high-energy, charged intermediates are circumvented, this termolecular reaction can give rise to a high kinetic barrier. As such, stepwise ET-PT or PT-ET mechanisms are often favored kinetically.37,38 Conversely, for electrochemical PCET, the reaction between the PCET substrate and the proton donor/acceptor occurs at an essentially infinite electrode surface (Figure 1B). As such, the kinetic barriers for an electrochemical CPET process will be different than the termolecular homogeneous example. Second, the electron transfer driving force changes during a potential-sweep electrochemical experiment. It has been demonstrated that via careful selection of chemical oxidants, the PCET oxidation of a tungsten-hydride bond can be induced to undergo a CPET mechanism.10 In an electrochemical experiment such as cyclic voltammetry, the driving force for electron transfer changes as a function of time. This suggests that the PCET mechanism may change during a single experiment as the potential is swept (Figure 2). It has already been

electrochemical PCET processes are crucial for advancing HER and other catalytic processes. Studies of PCET reaction mechanisms to date have primarily involved organic substrates relevant to biological charge transport pathways, with only a few examples of transition metal-based systems.8−11 Further, while a rich literature exists for the study of homogeneous-based PCET, where electron and proton transfers occur between discrete molecules (Figure 1A),9,12−16 there are fewer reports which

Figure 1. Cartoon depictions of (A) homogeneous termolecular PCET (case for two donors and one acceptor shown; more partners are possible) and (B) electrochemical PCET.

detail individual electrochemical PCET steps (specifically where the electrode is a partner in the electron transfer reaction, Figure 1B).11,17−24 This is not due to a lack of methodology or theory extensive information and experimental examples exist for the study of electrochemical mechanism.21,22,25−32 In addition, detailed theoretical analyses for mechanism-dependent electrochemical PCET rate expressions have been reported.33−35 However, only recently have these methods received more use in the study of PCET processes relevant to energy storage.21,23,24,27,31,36 In the study of electrochemical mechanisms, elementary steps are generally divided into two categories: electrochemical (E) steps and chemical (C) steps. An E step involves the movement of an electron to/from the electrode, whereas any other homogeneous chemical transformation corresponds to a C step. Additional subscripts may designate reversibility (subscript r) or irreversibility (subscript i) of individual steps, while the term (EC) indicates a concerted E and C step. This classic electrochemical nomenclature is readily paired with traditional homogeneous PCET terminology. For example, an electron transfer followed by a proton transfer, referred to as ET-PT in PCET literature, is designated more generally as an EC electrochemical mechanism. Meanwhile, the concerted-proton electron transfer (i.e., CPET) pathway would be designated (EC). Proton transfer generally occurs between discrete molecular proton donors and proton acceptors; however, electron transfer may occur between solution species or, as considered more frequently, to and from the electrode directly. As electron transfer occurs between solution and an electrode, two interesting and, as of yet, incompletely explored influences on the PCET mechanism arise. First, the kinetic barrier of the CPET pathway is impacted. In purely homogeneous systems, PCET may involve three discrete molecules which must encounter one another in a ternary

Figure 2. As applied potential changes throughout a variable-potential experiment, the PCET mechanism accessed may change as the driving force changes.

demonstrated that mechanisms for hydrogen evolution can in some cases switch from ECEC to EECC as the applied potential is scanned cathodically.24,39 This strongly suggests that a change of potential may promote a change in the mechanism of an individual PCET event. In this Perspective, the methods available to chemists for the study of PCET reactions in organic solvents will be introduced, with a focus on recent developments in electrochemical methods and their applications in acetonitrile. Our efforts in applying these methods to the elucidation of PCET mechanisms of hydrogen-evolving catalysts will then be discussed, highlighting the role of each method in building a broad mechanistic picture. Finally, three case studies of model systems relevant to hydrogen production will be presented to reflect upon different origins of peak shifts observed in cyclic voltammograms and their mechanistic implications.

2. METHODS FOR ELUCIDATING REACTION MECHANISMS AND EXPERIMENTAL CONSIDERATIONS In this section, methodology for the elucidation of PCET mechanisms, particularly PCET reactions in catalysis, will be discussed. The primary methods utilized in our laboratory and others will be described, along with the information that can be gained from their application. Care will be taken to highlight factors influencing the quality of the data and analysis, as well as potential pitfalls. 2.1. Time-Resolved Spectroscopy. Two additional spectroscopic techniques have found utility in the study of PCET reactionstopped-flow spectroscopy and transient absorption 3645


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py, which has a time resolution of microseconds due to the necessary mixing time, transient absorption spectroscopy allows investigation of kinetics in the pico- to nanosecond regime. 2.2. Electrochemistry. Electrochemical methods have vast precedent in the study of reaction mechanisms,25,26 and consequently, electrochemistry has been increasingly utilized in the study of PCET reactions. One key benefit of electrochemistry is the dynamic control over the energy of electrons. Of the multitude of electrochemical techniques, cyclic voltammetry has emerged as the favorite tool for uncovering reaction mechanisms and determining rate constants due to the rich information available from current−potential responses and the extensive analytical methods established for interpreting resulting data. Additionally, cyclic voltammetry is a nondestructive technique; only the minute volume of reactants immediately next to the working electrode (the reaction layer) is involved in a measurement. Finally, over the last 5 years, great strides have been made in the development of mathematical models to extract quantitative kinetic information from cyclic voltammograms of multielectron, multisubstrate reactions. Presented here is a brief summary of the electrochemical tools most commonly used in our laboratory; more extensive discussions are available.26,32 2.2.1. Kinetic Analysis Using Catalytic Plateau Current and Half-Wave Potential. In the case of electrochemical catalysis under ideal conditions with no limiting side phenomena or substrate depletion, an S-shaped response with superimposable forward and reverse traces is obtained by cyclic voltammetry.32 An observed rate constant kobs can be extracted using the experimental catalytic plateau current (ipl, which is independent of scan rate) and the potential at which half the catalytic current is reached (Ecat/2). For a two-electron catalytic process and assuming that both electron transfers occur at the electrode with the second reduction being easier than the first, the catalytic plateau current is given by eq 1 where [Cat]0 is the bulk solution catalyst concentration, D is the catalyst diffusion coefficient, F the Faraday constant, and A is the electrochemically accessible electrode surface area.

spectroscopy. Both techniques take advantage of changes in the absorption profiles during a reaction to gain mechanistic and kinetic information; specific details are discussed below. 2.1.1. Stopped-Flow Spectroscopy. Stopped-flow spectroscopy is a powerful tool for extracting kinetic information and has been used to examine PCET in various systems.10,40−43 Stoppedflow has been applied to the study of both catalytic cycles and elementary reaction steps where the reaction is accompanied by clear spectral changes. In a typical stopped-flow experiment, the contents of two syringes are rapidly mixed and the absorbance of the system recorded over time, either as full spectra at designated time delays or as single-wavelength kinetics traces (absorbance vs time). A general-purpose UV−vis absorption spectrometer may be used in cases where reaction kinetics are slow.44 Stopped-flow methods are amendable to monitoring the reaction between many reagent types or even combinations of reagents. For PCET studies, these may be chemical reductants, oxidants, hydride donors, acid, or base, for example. Once the two solutions have been rapidly mixed, the absorbance is recorded as a function of time. When monitoring a single elementary step, pseudo-first-order conditions are often employed, enabling the reaction kinetics to be fit to a single exponential to extract an observed rate constant kobs. For complex reactions, the kinetics traces can be fit using a kinetics model based on a series of differential equations describing the concentrations of each reactant, intermediate, and product as a function of time. Obtaining rate constants through stopped-flow measurements has been shown to corroborate electrochemically derived rate constants.24,40,41 2.1.2. Transient Absorption Spectroscopy. Time-resolved optical monitoring techniques can also be coupled to phototriggering methods to initiate PCET reactions. In traditional laser flash-photolysis methods, excited-state reductants (including ruthenium, rhenium, iridium, and copper polypyridyl complexes) are employed to initiate the one-electron reduction of a redox-active molecule upon excitation by a pulsed laser.8,45−53 In the presence of a proton source, reduction will often be accompanied by a protonation reaction.12,54,55 As in stoppedflow spectroscopy, the distinct optical signatures of the different redox forms of the catalysts, as well as the protonated species, allow the reaction intermediates and products to be spectroscopically detected and monitored. However, in these phototriggered transient absorption experiments, it is the change in the sample absorbance (as opposed to the absolute absorbance) that is monitored as a function of time. Data can be recorded as difference spectra at designated time delays or as singlewavelength kinetics traces, depending upon experimental setup. While phototriggered reactions typically involve the initiation of electron transfer, proton transfer reactions may be photoinitiated as well using photoacids.12 Photoacids are species that become powerful proton donors in their electronic excited states, allowing in situ generation of a proton donor upon pulsed laser excitation.56,57 Beginning with the reduced metal species, photoexcitation of the photoacid will promote proton transfer to yield a PCET reaction intermediate. Like photoinduced electron transfer reactions, proton transfer reactions can also be monitored optically if protonation is accompanied by a distinct change in absorption. While identifying the conditions to phototrigger an electron transfer or proton transfer reaction for a laser flash-photolysis experiment may be challenging, especially when other reactive substrates are present in solution, this time-resolved spectroscopy allows much faster reactions to be optically monitored. Compared to stopped-flow spectrosco-

i pl = 2FA[Cat]0 D kobs

(1)

eq 1 is often divided by the one-electron peak current in the absence of substrate, i0p, given by the Randles−Sevcik equation (eq 2, where F is the Faraday constant, T is temperature, and υ is scan rate) to yield eq 3, which eliminates the need for the diffusion coefficient and electrode surface area to be independently determined: i p0 = 0.446FA[Cat]0

i pl i p0

= 0.448

kobsRT υF

DυF RT

(2)

(3)

Expressions for kobs as a function of the rate constants for the first and second chemical steps, k1 and k2 respectively, can be derived for specific mechanisms.28 For the mechanisms ECEC and EECC, assuming S-shaped catalytic responses, expressions for kobs as well as Ecat/2, the potential at which half of the plateau current is reached, are given in Table 1 as a function of k1, k2, and E1/2, the potential of the electrochemical event triggering catalysis. In many cases, either k1 or k2 is rate limiting, resulting in a collapse of the equations in Table 1 simpler expressions. 3646


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figures of merits of different homogeneous catalysts by illustrating the trade-off between activity and overpotential. The turnover frequency is defined as the moles of product generated per unit of time divided by the moles of the active form of the catalyst present in the reaction-diffusion layer, and the expression of the maximal turnover frequency TOFmax is dependent upon the mechanism considered.28,32

Table 1. Values of kobs and Ecat/2 as a Function of the Rate Constants of the Two Chemical Steps Considered28,39 ECEC

EECC k1

k1

kobs

Ecat/2

1+

1+

k1 k2

E1/2 +

RT ⎛ ln⎜1 F ⎝

⎞ ⎟ + k2 ⎠ k1

k1 ⎛ k2 ⎞ k2 ⎜⎜1 + ⎟⎟ k1 ⎠ ⎝

⎛

E1/2 +

RT ⎜ ln⎜1 F

⎜ ⎝

+

k1

⎛ k 2 ⎜1 + ⎝

⎞ ⎟ ⎟ k2 ⎞ ⎟⎟ k1 ⎠ ⎠

TOF =

(5)

The overpotential η is defined as the difference between the applied potential at the electrode (E) and the apparent standard potential of the reduction of protons to hydrogen in the Solv conditions of acid and solvent studied ESolv HA/H2: η = EHA/H2 − E. The catalytic Tafel plots are generated using TOFmax, Ecat/2, and ESolv HA/H2 at a fixed stated acid concentration. Determination of the value of ESolv HA/H2 is discussed in section 2.3.1. Savéant and Artero suggest using the standard concentration of 1 M to generate the catalytic Tafel plots. Practically, researchers experimentally determine the rate constants for the transformations studied using the methods described above, and they generate catalytic Tafel plots by calculating at each overpotential the expected turnover frequency. The aim of this tool is to provide a simple visual representation of the activity of different catalysts at different overpotentials. Recent concerns have emerged over the usefulness of this tool as the activity reported is necessarily idealized. One such concern was over the actual acid concentration in the medium, as well as turning on catalyst-dependent side reactions and deactivation pathways at such high hypothetical acid concentrations.39 2.3. Considerations for Choosing an Acid Source. The nature and choice of acid source is of paramount importance in PCET processes. Several key parameters should be considered when choosing an acid, including acid pKa, structure, standard reduction potential for H2 production, potential of background reduction at the electrode, and solution thermal equilibrium processes. This section will briefly touch on the influence of each parameter, focusing on practical guidelines to be mindful of during PCET studies. A case study of the impact of these parameters on reactivity is presented in section 3.3. 2.3.1. Acid pKa, Thermodynamic Potential for Reduction, and Structure. Acid strength, quantified through pKa, is the parameter most commonly used in the choice of acid for the study of PCET reactions. Extensive tabulation of acid pKa values in acetonitrile is available; this existing data spans a wide range of acid types, including phenols, carboxylic acids, anilinium and pyridinium derivatives, among others.64−67 The pKa of the chosen acid directly relates to the overpotential for the specific electrochemical PCET reaction. The required overpotential for a catalytic reaction has implications for the comparison of catalysts as discussed in the context of catalytic Tafel plots and for analyzing catalytic thermochemical cycles. For the example of hydrogen evolution, evaluation of the overpotential η as a function of the applied potential E, necessitates knowledge of two parameters: the standard reduction potential of H+ to H2 in the solvent (E0,Solv H+/H2) and the pKa of the acid (HA) in the solvent (Solv):68

kobs collapses to k rate limiting for both cases, while Ecat/2 collapses to E1/2 only when k1 is rate limiting. 2.2.2. Foot-of-the-Wave Analysis (FOWA) for Kinetic Evaluation. Experimentally, S-shaped cyclic voltammograms and consequently accurate plateau currents may not be accessible due to substrate consumption, catalyst degradation, or other side phenomena.27,32 This prohibits the analysis of the plateau current to extract kinetic information. As the influence of these side reactions increase throughout a cyclic voltammetry experiment, the plateau current becomes inaccessible. Costentin and Savéant recently recognized that in these cases, the current (i) vs applied potential (E) plot of nonideal catalytic cyclic voltammograms overlays with the idealized cyclic voltammogram near the beginning of the catalytic scan.27 Consequently, kinetic information can still be gleaned using data recorded near the foot of the wave where the influence of such phenomena is minimized in a method coined foot-of-the-wave analysis (FOWA).27 The generic expression of the current near the foot of the wave, assuming fast electron transfers occurring only at the electrode, is i pl /i p0 i = F i p0 1 + exp⎡⎣ RT (E − Ecat/2)⎤⎦

TOFmax F Solv ⎡ 1 + exp⎣ RT (E HA/H − η − Ecat/2)⎤⎦ 2

(4)

where the specific definition of ipl and Ecat/2 varies with mechanism (see eq 1 and Table 1).28,32 FOWA has rapidly proven extremely valuable with successful application to the study of a number of electrocatalytic PCET reactions including H2 evolution, CO2 reduction, and water oxidation.24,27,39,58−63 Recently, the influence of the heterogeneous electron transfer rate and the percentage of the current response used in FOWA on the kinetic results were reported.39 It was found that utilizing larger percentages of the catalytic wave for FOWA resulted in larger errors; additionally, the linear portion of the wave decreased with smaller (and more realistic) values of the heterogeneous electron transfer rate constant. This study emphasizes the importance of only using the linear portion of the data for FOWA. Finally, obtaining accurate values from FOWA is only possible when an accurate value is known for the potential of the catalysisinitiating redox process (E1/2). The potential of an electrochemical event in the absence of substrate is often used, but care must be taken to ensure that this redox process is indeed the initiating event for catalysis. This can be difficult to obtain in some cases, especially when the solvent is the substrate. 2.2.3. Catalytic Tafel Plots and Standard States. To benchmark catalysts studied in different conditions, in the presence of different acids and at different overpotentials, Savéant and co-workers proposed the use of catalytic Tafel plots. These plots, of the log of the turnover frequency (TOF) as a function of overpotential (η) (the log plot of eq 5), showcase the

Solv η = E HA/H − E = E H0,Solv − 0.059 × pK a Solv − E + /H 2 2

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(6)


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depending on the acid structure, concentration, and the nature of the solvent.

Recently, a methodology was outlined for the direct electrochemical determination of ESolv HA/H2 via measurement of the open circuit potential of an electrochemical cell containing the respective acid and 1 atm of H2 gas at equilibrium in the solvent studied.69 To date, the experimental complexity of the measurement has limited its use. A recent report investigating the kinetics of the CPET reaction between acids and gold electrodes in acetonitrile shed light on the influence of acid structure in terms of steric hindrance.70 Faster kinetics of the CPET process to the electrode were measured for triethylammonium compared to diisopropylethylammonium. This effect was attributed to the increased Tolman cone angle of 200° for the diisopropylethylammonium cation compared to 150° for triethylammonium. The possibility of interactions between catalysts and conjugate bases is another parameter to be considered in the choice of acid/ base pairs for mechanistic studies. The influence of base structure has recently been highlighted in our study of catalytic proton reduction by cobaloxime complexes with substituted aniliniums as an acid source (see below). It has been shown that during the course of catalysis, as conjugate base aniline is generated, the aniline molecules can bind to the cobalt center, influencing what the active catalyst is and so altering the cyclic voltammetry responses.71 This effect is readily apparent in the cyclic voltammetry of Co(dmgBF2)2(CH3CN)2 (dmgBF2 = difluoroboryldimethylglyoxime) with and without added 4-MeOaniline (Figure 3). Upon base addition, the CoIII/II wave

homoconjugation: BH+ + B → BHB+

(7)

dimerization: 2BH → (BH)2

(8)

heteroconjugation: BH+ + H 2O → BH3O+

(9)

Homoconjugation is particularly strong for phenols and carboxylic acids. For example, phenolate is reported to have a homoconjugation constant close to 16 000 M−1, while benzoate has a constant of almost 4000 M−1.64 By contrast, aniline has a much lower homoconjugation constants of 4 M−1;64 it is often assumed that most aniline derivatives have low homoconjugation constants.72 As was recently demonstrated in the study of [Ni(PR2 NR2 ′)2]2+ hydrogen evolution catalysts, failure to account for homoconjugation when determining rate constants from cyclic voltammetry data can lead to misinterpretations of the influence of added base as well as errors in the determination of rate constants. However, careful consideration of homoconjugation can permit accurate determination of kobs and mechanistic assignment.39 Acid dimerization is another process to consider, especially when the acid dimer has a lower pKa than the parent acid. The presence of an acidic species with a different pKa than the parent acid may result in multiple PCET mechanisms, complicating kinetic and mechanistic analysis. Dimerization can occur in aprotic organic solvents when interaction between two acid molecules yields a stabilization greater than solvation by the aprotic organic solvent. Dimerization has been found to influence the reactivity of carboxylic acids in acetonitrile.73 Heteroconjugation arises between acids and bases of a different species; for example, the interaction of a carboxylic acid with water. In the context of electrochemical studies in acetonitrile, the effect of heteroconjugation has only been considered when water is added to the organic solvent.73 2.3.3. Reduction of Acids at Electrode Surfaces. For solutions containing protic species, an important control experiment to keep in mind is at what potential these protic species are directly reduced at the electrode. In other words, control experiments for each individual component of an electrochemical system are crucial. Glassy carbon electrodes have long been used in the study of fuel-forming catalysts, especially hydrogen evolution, due to their low background activity for direct, heterogeneous acid reduction. While the background activity is less significant than on other electrode materials such as platinum, direct acid reduction at glassy carbon electrodes still occurs. Our group recently published a systematic study of the direct reduction of 20 acids in acetonitrile on glassy carbon electrodes.74 These acids spanned over 35 pKa units in acetonitrile, with the experimental direct acid reduction potential, reported as the current inflection potential (Einf), covering a range of over 2 V. These results are visually summarized in Figure 4, where the color intensity reflects the current intensity observed as a percentage of the current observed at Einf. This study underscored four factors influencing the observed voltammograms: electrode fouling, irreproducibility from scan to scan, curve crossing, and acids with erratic backgrounds. The influence of adding water to each acid was also studied. While no significant peak shifts were observed with water addition, the current increased for a number of acids known to have a high homoconjugation constants. The current increase was proposed

Figure 3. Cyclic voltammograms recorded at 100 mV/s in 0.25 M [Bu4N][PF6] of 0.38 mM of Co(dmgBF2)2(CH3CN)2 in the absence (a) and presence (b) of 2 mM 4-methoxyaniline. Reprinted with permission from reference 71. Copyright 2016 American Chemical Society.

reversibility and potential is modified, without apparent changes to the CoII/I wave. Separate UV−vis titration experiments demonstrated that one and two aniline molecules can reversibly coordinate to the catalyst.71 2.3.2. Acid Homoconjugation, Heteroconjugation, and Dimerization. Acid and base molecules are well-known to interact in solution via the processes of homoconjugation, heteroconjugation, and dimerization. These three processes, depicted in eqs 7 to 9, have a varying degree of importance 3648


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As the number of reports grows, some key structural motifs have emerged as susceptible to decomposition. Cobalt and nickel compounds containing CN and N−O bonds were reported by a number of groups80,81,83,84,89 to decompose in acidic and reducing conditions to Co- and Ni-containing nanoparticles, respectively. Firm experimental evidence that CN bonds can predispose a complex to decomposition was provided by our group via direct comparison of two complexes either with or without CN bonds.87 The presence of Ni−S bonds also appears to render compounds lacking CN or N−O bonds susceptible to decompositions, as discussed by Roberts85 and by our group.11 In these cases, it appears that cleavage of adjacent S−C bonds with concomitant ligand destructionis favored by Ni−S bonds.11,85 Finally, we recently demonstrated with an Fe complex that protic conditions are not required for decomposition.88 2.5. Digital Simulations. The modeling of electroanalytical experiments has advanced to a level of great accuracy over the last half century. Two general approaches have been used: (1) The semianalytical solutions to the diffusion equations have been obtained through the integral equation method in order to describe the voltammetric response of specific mechanisms (e.g., equations presented in section 2.2.1). These solutions are often applied directly or used to generate working curves based on dimensionless parameters.26 (2) Finite difference or finite element numerical methods have been used to solve the diffusion equations in conjunction with the kinetic equations describing the chemical reaction of interest (known as digital simulation). Both methods are very valuable, but digital simulation is generally found to be more accessible to most researchers. Digital simulation software can be used to generate simulated voltammograms that can be directly compared with experimental voltammograms. However, successful simulation of experimental data requires accurate assessment of a large number of parameters, and the sensitivity of electroanalytical techniques to environmental factors that cannot be emulated in simulation often prevents satisfactory fitting. In our experience, we have found digital simulation to be most valuable in addressing “what if” questions and to test if the already derived equations can be extended to a proposed mechanism. Both of these cases are highlighted below in our mechanistic studies of cobaloxime.71 2.6. Traversing the Zone Diagram To Access S-Shaped Voltammograms. Zone diagrams pictorially depict the voltammetric waveform shapes expected as a function of dimensionless parameters;91 the zone diagram for an EC′ mechanism is shown in Figure 5. For an EC′ reaction, the kinetic (λ) and excess factors (γ) describe how waveforms will evolve as a function of parameters such as ke (the rate constant for the reaction of the reduced catalyst and the substrate), υ (the scan rate), and [Cat]0 and [CA]0 (the bulk concentration of catalyst and substrate A, respectively):26,91

Figure 4. Acid reduction potentials in acetonitrile (25 mM acid, 100 mV/s) depicting the range between the direct reduction potential Einf. and the approximate thermodynamic potential for hydrogen evolution. Caveats for three acids are listed. Adapted with permission from ref 74. Copyright 2014 American Chemical Society.

Figure 5. Kinetic zone diagram and simulated CV waveforms for the catalytic EC′ mechanism of the one electron reduction of a substrate via a redox catalyst mediator. Reprinted with permission from ref 32. Copyright 2014 American Chemical Society.

to be the result of the added water stabilizing the in situ generated conjugate bases and so freeing additional acid molecules tied up in homoconjugation for direct electrode reduction. Overall, this study allows for easier and mindful selection of a suitable acid for electrochemical PCET reactions. 2.4. Homogeneous to Heterogeneous Transformation of Catalysts. The past decade has witnessed increased awareness of the possibility of transformation from a homogeneous precatalyst to heterogeneous catalyst.75−78 There has been a recent surge in the identification of molecular species which decompose to heterogeneous material deposited on the electrode surface; known examples are depicted in Table 2.11,79−88

λ=

γ=

⎛ RT ⎞⎛ ke[Cat]0 ⎞ ⎜ ⎟⎜ ⎟ ⎝ F ⎠⎝ ⎠ υ

(10)

[CA ]0 [Cat]0

(11)

By changing these experimental parameters, waveforms conforming to different zones can be accessed.32 3649


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Table 2. Metal Complexes Known to Electrochemically Degrade in Organic Solventsa under Catalytic Conditions To Form Electrode-Adsorbed Heterogeneous Materialsb

a All examples were in either acetonitrile or dimethylformamide. bThis table does not include examples where the homogeneous catalyst only degrades after harsh prolonged catalysis; e.g., ref 90.

Diagram could be qualitatively extended to multielectron, multisubstrate reactions.92 In addition to υ, [CA]0, and [Cat]0, the pKa of added acids could also be used to traverse the zone diagram as it was found that the observed rate constant is linearly dependent on acid strength.71,92 Using the methods detailed above in section 2.2.1, global rate constants can be extracted from the plateau region of S-shaped voltammograms of Zone KS (and the related Zone KD) for both EC′ and various multielectron, multisubstrate reactions, such as the ECEC pathway.26,28 Further, it has long been known that

Until recently, it was not clear if the EC′ Zone Diagram could be directly extended to the more complex multielectron, multisubstrate reactions associated with fuel production.92 However, voltammograms have now been reported for the catalytic reduction of protons from a series of eight parasubstituted aniliniums to hydrogen by the catalyst Co(dmgBF2)2(CH3CN)2. Notably, the voltammetric responses recorded for varying kinetic and excess factors apparently spanned all of the waveform shapes associated with the EC′ Zone Diagram. Detailed analysis confirmed that the EC′ Zone 3650


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the [Co(dmgBF2)2(CH3CN)2]0/− couple. Notably, the Ecat/2 (the half-wave potential of the sigmoidal response) shifts to values positive of E°′([Co(dmgBF2)2(CH3CN)2]0/−). For acids of pKa > ca. 9.5, the catalytic response currents approach a plateau shape as the acid concentration is increased; these catalytic plateau currents (ipl) were found to be first-order in both acid concentration and catalyst concentration, consistent with prior results.97 These three observations provide evidence to support assignment of the reaction mechanism (Figure 6A). Specifically, the shifting Ecat/2 indicates that the first chemical step (C1, protonation of [Co(dmgBF2)2(CH3CN)2]− to form a putative CoIII-hydride intermediate) is rapid, and the subsequent chemical step(s) is(are) rate-limiting. Previously, two independent theoretical124,126 and one experimental125 study have shown that the CoIII-hydride intermediate is more readily reduced than the parent CoII species, indicating that this intermediate will be spontaneously reduced. Further, based on the influence of acid and catalyst concentration on the catalytic plateau current, the rate-limiting step for catalysis is also first-order in both acid and catalyst, indicating a second (rate-limiting) protonation step (C2) and an overall heterolytic route for hydrogen production. Together, these data support a heterolytic ECEC pathway for H2 evolution from anilinium acids in acetonitrile. Notably, this assignment is consistent with that from theoretical work54,124 as well as conclusions from photocatalysis and laser flash photolysis experiments.112,125 Of note, a homolytic route involving the bimetallic reaction of two CoIII-hydride species has previously been recognized as a viable pathway for H2 production by both computational studies124,126 and thermochemical analysis.54 While clear evidence diagnosing a homolytic pathway for hydrogen production has not yet been reported, 29 recent work investigating the cleavage of H2 by Co(dmgBF2)2(CH3CN)2 under high H2 pressures suggests this reaction proceeds via a homolytic cleavage route.127−129 As such, a homolytic reaction pathway for hydrogen production may be accessible under alternative experimental conditions, or could be occurring as a minor parallel pathway. With this mechanistic assignment, kinetic details for elementary steps can be elucidated through analysis of the voltammograms. The second-order global rate constants for catalysis were determined from the analysis of the acid concentration-dependent plateau currents for acids with pKa > ca. 9.5 (Figure 6B). As the plateau current reflects the observed rate constant, kinetic details of the rate-limiting step can be gleaned directly. As described above, the second protonation is the rate-limiting step under these conditions and described by rate constant k2. Across the series of aniliniums with pKa > ca. 9.5, catalysis was found to accelerate as the strength of acid increased and a linear correlation is observed between log(k2) and acid pKa (slope = −0.77). For acids with pKa < ca. 9.5, catalytic voltammograms remained peaked and the catalytic responses were determined to be controlled by diffusion of the acid to the reaction layer. As such, second-order rate constants could not be determined for these acids; however, extrapolation of the linear relationship determined for weaker acids allows these rate constants to be estimated. Rate constants for the first protonation step (k1) were determined from foot-of-the-wave (FOWA) analysis. Like k2, a linear correlation between log(k1) and pKa was observed (slope = −0.97). Values for k1 are ca. 3 orders of magnitude greater than

kinetic information can also be extracted from waveforms in the “total catalysis” region (Zone KT2) for a simple EC′ reaction.32,91 The catalytic peak potential of a total catalysis voltammogram varies as a function of ke, υ, [CA]0, and [Cat]0 (see section 4.1).32 More recently, peak shift analysis has been extended to the ECEC mechanism and demonstrated, under limiting conditions, to provide information about the rate constant k1.71 Ultimately, the primary utility of a multielectron, multisubstrate zone diagram is its use as a navigational tool to modify experimental conditions in order to enter either the zones in which kinetic information can be extracted.

3. CASE STUDIES: REACTION MECHANISMS AND KINETICS OF H2-EVOLVING CATALYSTS The electrochemical and spectroscopic methods discussed in section 2 have been applied in concert to elucidate the reaction pathways of catalysts for hydrogen evolution. Recent work has focused on two highly active catalysts, Co(dmgBF2)2(CH3CN)2 (dmgBF 2 = difluoroboryldimethylglyoxime) and [Ni(P2PhN2Ph)2]2+ (P2PhN2Ph = 1,3,5,7-tetraphenyl-1,5-diaza-3,7-diphosphacyclooctane) (Scheme 2). In addition to identifying Scheme 2. Co(dmgBF2)2(CH3CN)2 (left) and 2+ Ph [Ni(PPh 2 N2 )2] (right)

dominant PCET reaction pathways, these studies have yielded new insight into the kinetics of elementary reaction steps and revealed linear free energy relationships between rate constants for these individual steps and acid pKa. The complexities of proton source reactivity in acetonitrile have also been established through the investigation of reaction kinetics as a function of various acid parameters. 3.1. Mechanistic Study of Electrocatalytic H2 Evolution by a Cobaloxime. Cobaloximes and their derivatives have been incorporated as catalysts in both electrochemical and photochemical H2 production cycles.93−123 While their operating mechanisms have been explored in depth under certain conditions using both experimental and theoretical methods,54,93,94,96−98,124−126 electroanalytical methods were recently applied to explore the reaction pathways and kinetics of Co(dmgBF2)2(CH3CN)2 (dmgBF2 = difluoroboryldimethylglyoxime) as a function of acid strength.71,92 Through the use of a series of analogous acids, specifically para-substituted aniliniums, complicating factors resulting from different homoconjugation constants across dissimilar acids were reduced, allowing studies to focus on the direct effect of acid pKa on reaction kinetics and mechanism. It should be noted that Co(dmgBF2)2(CH3CN)2 decomposes in the presence of strong acids; in the extreme case of strong acid and reducing conditions cobalt-containing nanoparticles form, as noted in Table 2 above. In our studies of Co(dmgBF2)2(CH3CN)2, great care was taken to limit the extent of decomposition.71 Cyclic voltammograms of Co(dmgBF2)2(CH3CN)2 recorded in acetonitrile in the presence of para-substituted aniliniums exhibit an increase in cathodic current and loss of reversibility at 3651


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determined from this limiting current value (ca. 125 s−1) likely reflects either an H−H bond formation step or H2 release. Kinetic and mechanistic insight gained from comprehensive analysis of the catalytic voltammograms contributes to a more detailed picture of hydrogen evolution from cobaloximes. Specifically, it was learned that upon reduction of the CoII species, rapid protonation (k1) yields a CoIII-hydride species (that may undergo tautomerization128−130). This intermediate is readily reduced to form a CoII-hydride (k2). The subsequent second protonation step is 3 orders of magnitude slower than the first step. The resulting CoII−H2 species releases H2 to close the catalytic cycle via an acid-independent step with a first-order rate constant of 125 s−1. 3.2. Mechanistic Study of Electrocatalytic H2 Evolution by [Ni(P2PhN2Ph)2]2+. [Ni(P2PhN2Ph)2]2+ (P2PhN2Ph = 1,3,5,7tetraphenyl-1,5-diaza-3,7-diphosphacyclooctane) and its family of hydrogen production/oxidation catalysts have been thoroughly studied over the past decade, with the notable result of clearly demonstrating the importance of the secondary coordination sphere in catalytic reactions.131−136 While original reports assumed a single catalytic mechanism, the possibility of multiple catalytic routes has been recently examined.24,39,137 Two generic pathways have been identified: an ECEC pathway, whereby the first protonation occurs on the NiI species after a single reduction, and an EECC pathway, where double reduction to form a Ni0 species occurs prior to both protonations. Two recent experimental studies probing these parallel reaction pathways approached the problem very differently. Wiedner and co-workers examined the reactivity of [NiPh 2+ (PPh using dimethylformamidium (DMFH+) as the 2 N2 )2] proton source.39 DMFH+ is a relatively strong acid in acetonitrile with a pKa of 6.1, resulting in cyclic voltammograms with both the ECEC and EECC pathways readily apparent. Using a mixture of electrochemical techniques, primarily based around FOWA, the overall contributions of the two different pathways were deconvoluted. In work in our lab, we utilized a weaker acid (anilinium, pKa = 10.6) to isolate the EECC mechanism and then used several of the previously discussed methods to determine the individual rate constants of the elementary reaction steps.24 3.2.1. Electrochemical Studies To Elucidate Kinetics of Ni0 Protonation and Subsequent Reactivity. While analysis of cyclic voltammograms under pseudo-first-order conditions is useful for extracting kinetic information, experiments carried out under more stoichiometric catalyst:substrate conditions can provide additional information on how the reaction proceeds. In 2+ Ph the analysis of [Ni(PPh 2 N2 )2] , stoichiometric addition of anilinium confirmed rapid reactivity of acid with the Ni0 species by a signature kinetic peak shift (discussed further in section 4.2) and the appearance of a new oxidation feature on the return trace. Given that two electrons and a single proton had been transferred to the catalyst, it was natural to assume hydride formation. Independent synthesis and cyclic voltammetric analysis of a NiII-hydride confirmed that the new oxidation feature observed did correspond to the hydride species (Figure 7).24 With the first chemical step established as hydride formation, we next examined the details of this elementary step under catalytic conditions. First, a rate constant was determined by analyzing the kinetic peak shift of the NiI/0 wave in the presence of acid. Conveniently, catalytic turnover could be prevented on the electrochemical time scale, even under pseudo-first-order conditions, by the addition of excess base. This rate constant was corroborated with FOWA. FOWA uniquely provides the rate

Figure 6. (A) Reaction mechanism for hydrogen evolution catalyzed by Co(dmgBF2)2(CH3CN)2 as supported by detailed analysis of catalytic CVs. (B) Rate constants k1 and k2 determined from FOWA and plateau current analysis, respectively, of acid concentration-dependent catalytic voltammograms of Co(dmgBF2)2(CH3CN)2 and a series of parasubstituted aniliniums. k2 values for acids 6, 7, and 8 were extrapolated from the linear relationship established at higher pKa values. kΩ represents an acid-independent step and was determined from catalytic voltammograms recorded at high acid concentrations in which the voltammetric responses were acid concentration and pKa-independent. Acids identity: (1) 4-methoxyanilinium, (2) 4-t-butylanilinium, (3) anilinium, (4) 4-chloroanilinium, (5) 4-trifluoromethoxyanilinium, (6) 4-(methylbenzoate)anilinium, (7) 4-trifluoromethylanilinium, and (8) 4-cyanoanilinium. Adapted with permission from ref 71. Copyright 2016 American Chemical Society.

those of k2, consistent with the interpretation of an ECEC pathway with a rate-limiting second chemical step. At high acid concentrations, the plateau currents reach a limiting value that is acid concentration and pKa-independent. This observation extends to both the weaker acids and the stronger acids described above which exhibit only peak-shaped voltammograms in their acid-dependent region. Across the range of acids explored, this limiting plateau current is constant, though the concentration of the acid required to reach this limiting value increases as acid strength decreases. The first-order rate constant 3652


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chemical step of H2 productionshould control the overall reaction kinetics. If this is the case, the plateau currents should directly reflect this reactivity and spectroscopic monitoring of the reaction of NiII-hydride with anilinium should result in a similar observed rate constant. The NiII-hydride is a stable complex with a distinct absorption Ph 2+ spectrum from [Ni(PPh 2 N2 )2] , allowing us to observe its reactivity with anilinium directly, using stopped-flow spectroscopy. As stopped-flow spectroscopy is a time-resolved technique compared to the steady-state nature of electrochemical methods, complementary chemical reactivity information can be gathered by employing this experimental technique. Reaction of the NiIIhydride with anilinium affords a kinetic trace with two distinct regions: a fast component on the order of seconds, and a slow component on the order of minutes (Figure 9). Under pseudo-

Ph 2+ Figure 7. Cyclic voltammograms of 1.5 mM [Ni(PPh and 2 N2 ) 2] approximately 1 equiv of anilinium (blue) compared with the cyclic voltammogram of 0.6 mM of the independently synthesized NiIIhydride (black) in CH3CN. The new oxidation observed at −0.4 V for [Ni(P2PhN2Ph)2]2+ in the presence of acid matches that of the independently synthesized hydride. Voltammograms recorded at 100 mV/s in 200 mM [NBu4][PF6] CH3CN solutions. Reprinted with permission from ref 24. Copyright 2015 American Chemical Society.

constant for only the first step in the catalytic reaction. Reasonable agreement (1.2 × 106 M−1 s−1 for peak shift analysis, and 6.5 × 106 M−1 s−1 for FOWA) was found between the resulting values.24 Subsequent reactivity of the NiII-hydride was then investigated. Comparison of the value estimated for kobs from the catalytic plateau currents and the value determined for rate constant k1 indicates that the overall rate of catalysis is not governed by hydride formation. As such, kinetic information obtained from the plateau current represents steps subsequent to the first protonation. At low acid concentrations the reaction is approximately first-order in acid; however, at higher acid concentrations, an acid concentration-independent region is reached (with anilinium the maximum rate constant was ∼20 s−1 as measured by the ECEC plateau equation, Figure 8).24 3.2.2. Investigation of Hydride Reactivity with Acid via Stopped-Flow Spectroscopy. The cyclic voltammetry studies described above establish that the formation of the hydride is rapid and these kinetics do not affect the overall catalytic rate described by kobs. With this assumption, it was asserted that the reactivity of the hydride with the proton sourcethe second

Figure 9. Stopped-flow kinetics trace for a solution 0.39 mM NiIIhydride and 100 mM anilinium. 88% of the change in absorbance is accounted for in the “fast” kinetics (shown in inset), which occurs over 2 s followed by a “slow” process occurring over the course of 5 min.

first-order conditions, the fast component was fit to single exponential kinetics. The rate constant obtained was found to be first-order with respect to the anilinium concentration.24 At first glance, the first-order response with anilinium appeared to be in conflict with the acid-independent response obtained from cyclic voltammetry at higher acid concentrations.

Ph 2+ Figure 8. Observed rate constants obtained from the current plateau (blue circles) for the reaction of [Ni(PPh 2 N2 )2] and anilinium, and from stoppedflow kinetics traces (red circles) for the reaction of NiII-hydride and anilinium. Black circles represent the fit for the mechanism shown in the right panel, and the appropriate kinetics models are shown for the time-resolved (red) and steady-state (blue) kinetics. Data from ref 24.

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Changing the proton source structure significantly affects its reactivity (as will be shown in the following sections). The effect of added base on reaction kinetics was explored. When base is added to solution, the reaction kinetics were observed to slow. A trend between amount of base necessary to slow the reaction and acid pKa was observed. The higher the pKa, the less base required to lower the observed reaction rate (when pKa is greater than 8.6, reactivity of acids below a pKa of 8.6 do not have a pKa dependence as illustrated in Figure 10). 3.3.2. Molecular Association. Two different forms of molecular association were examined, homoconjugation and dimerization (aggregation and heteroconjugation were also considered for studies involving water). The list of proton sources with recorded homoconjugation constants in the Ph 2+ working range of [Ni(PPh (pKa < 12) is limited, but 2 N2 )2] the effect could be examined through the comparison of 4cyanoanilinium (pKa = 7, log(KHC) = 0.1),73 p-toluenesulfonic acid (pKa = 8.6, log(KHC) = 3),74 and dimethylformamidium (pKa = 8.6, log(KHC) = 1.7).74 While this series of acids had only small differences in their pKa values, the reaction rate constants were found to be highly dependent on the homoconjugation constants, with fast reaction rate constants observed for the acids with larger homoconjugation constants (p-toluenesulfonic acid (130 000 M−1 s−1) > dimethylformamidium (12 750 M−1 s−1) > p-cyanoanilinium (750 M−1 s−1 )).73 In past work, the acceleration of acid−base reactivity for acids with large homoconjugation constants has been attributed to a lowering of the “effective” pKa as a result of stabilization of generated conjugate base via homoconjugation with acid.69,72,136,143−145 However, rate constants for this series span more than 2 orders of magnitude, and this span is too large to be accounted for simply by homoconjugation. Likely, larger homoconjugation constant correlate with increased hydrogen bonding abilities, and these acids may better stabilize the proton transfer transition state. To our knowledge, proton source dimerization has yet to be considered in the PCET fuel forming literature. While there are no recorded dimerization constants for proton sources with a pKa less than 12, carboxylic acids are known to dimerize in acetonitrile and this generally results in a pKa for the dimerized species that is ∼3.5 units lower than the monomer.64 No dimerization constants are recorded for trifluoroacetic acid (pKa = 12.6) and trichloroacetic acid (pKa = 10.8). However, the reaction of trifluoroacetic and trichloroacetic acid with the NiIIhydride is second-order in acid, as would be expected for a highly reactive acid dimer with a very low dimerization constant. This is an important finding as trifluoroacetic acid is often used in electrocatalytic studies; indeed, many studies report secondorder kinetics with trifluoroacetic acid.136,146−151 3.3.3. Addition of Water to Acetonitrile Solutions. Because the target for most H2 evolution electrocatalytic systems is to operate in water, water is often intentionally added to acetonitrile solutions during catalysis to probe catalyst water stability and performance.140,152 Consequently, it is important to understand how the addition of water affects the proton source being employed. Two cases are considered here: (1) water becomes protonated and thereby hydronium is available to act as a proton source, or (2) water heteroconjugates with the proton source, affecting its reactivity. Hydronium has a very low pKa in acetonitrile (2.2), but the ability of water to form large aggregates with concurrent stabilization of the hydronium ion significantly increases its effective pKa with increasing water concentration. Figure 11 demonstrates the influence of water on the rate constant of reaction with NiII-hydride and acid over a range of

However, the discrepancy was reconciled by considering the possibility of an off-cycle intermediate, which has been extensively explored in several previous studies.39,136,138−142 These works reveal that the [Ni(PR2 NR2 ′)2]2+ molecules can be protonated in an “exo” position on the pendant amine which prohibits catalytic turnover. This effectively inactivates the catalyst prior to catalytic turnover. The “exo” position proton must be removed, followed by protonation in the “endo” position for productive catalysis to occur. Under the conditions accessed in cyclic voltammetry, an increase in the acid concentration results in an increase in the concentration of the inactive catalyst, reducing the apparent rate constant and giving the appearance of acid concentration independence. The time-resolved (stoppedflow) kinetics do not exhibit acid independence because no equilibrium can be established between the inactive and active forms of the catalyst. However, the generation of the inactive catalyst affects the kinetics analysis of the fast component and the slow component observed corresponds conversion of the inactive species to the active species followed by hydrogen release.24 3.3. Reactivity of NiIIH with Various Proton Sources. As the reaction of the NiII-hydride with acid was established to be the rate-limiting step for hydrogen production, this elementary reaction step provided a convenient platform to take a deeper look at the influence of the proton source on reaction mechanisms in acetonitrile. Reactivity of the NiII-hydride with a series of acids were examined using the same stopped-flow procedure described for the reaction with anilinium (section 3.2.2). The effect of various parameters, including pKa, presence of base, molecular association, and interaction with water, were explored.73 3.3.1. Acid pKa. As discussed above, pKa is the most commonly considered parameter for acids in PCET studies. In the reaction of NiII−hydride with acid, the anticipated linear free energy relationship is obtained between the log of the reaction rate and the acid pKa (Figure 10). With a sufficiently strong driving force for proton transfer the reaction rate constant becomes substrate pKa independent (pKa < 8.6). The data in Figure 10 is limited to a series of structurally related para-substituted aniline derivatives.

Figure 10. Second-order rate constants vs acid pKa for the reaction of NiII-hydride with para-substituted aniliniums. The linear region between pKa 12 to 8.6 has a slope of −0.7. Reprinted with permission from ref 73. Copyright 2016 American Chemical Society. 3654


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4. CASE STUDIES: MODEL SYSTEMS TO INVESTIGATE ELEMENTARY PCET REACTION STEPS As detailed above, catalysis involves elementary steps whose kinetics may be investigated by a variety of spectroscopic and electrochemical methods. A key question for driving the design of more energy-efficient catalysts is whether the elementary steps may be controlled through rational design of catalysts or experimental conditions. Toward this, efforts in our group have focused on furthering cyclic voltammetry techniques for the study of PCET systems. Shifts in the applied potential at which reactivity is observed are common in PCET systems (or all catalytic systems for that matter), yet the reason for potential shifts can differ greatly. Three separate factors, all dependent upon proton addition, are presented herein: First, the appearance of total-catalysis responses;71,92 second, a kinetic shift in the reduction potential and loss of reversibility;24 and third, a thermodynamic shift in the reduction potential.11 Distinguishing between these cases is not always straightforward; we hope that presentation of these three case studies can help aid researchers in elucidating the factors that perturb peak location. Special emphasis is placed on the thermodynamic potential shift, which in this case was found to correspond to a concerted proton electron transfer 4.1. Total Catalysis Gives Rise to a Second Peak. As noted above, cobaloxime has been reported to yield total catalysis responses under specific conditions (Figure 12). This occurs

Figure 11. Concentration of water versus observed rate constant for 1.25 mM triflic acid (pKa = 2.6), 1.3 mM dimethylformamidium (DMFH+, pKa = 6.1), and 2.5 mM 4-cyanoanilinium (4-CN, pKa = 7). Observed rate constants obtained from reaction with NiII-hydride to form hydrogen. Reprinted with permission from ref 73. Copyright 2016 American Chemical Society.

water concentrations for three acids. triflic acid (pKa = 2.6), dimethylformamidium (pKa = 6.1), and 4-cyanonanilinium (pKa = 7) were compared; in each case, addition of water resulted in hydronium formation, which showed a higher reaction rate than the original proton source. Eventually, further increases in water concentration began to shut down the reaction, as the excess water reacts as a base in solution to hinder turnover, as discussed above.73 Heteroconjugation of water to proton sources with higher pKa values was investigated through the addition of water to solutions of trifluoroacetic acid. The reaction rate decreased exponentially with the concentration of water, suggesting the acid heteroconjugates to water. This chaotropic activity breaks up the dimerization of trifluoroacetic acid. 3.4. Linear Free Energy Relationships in H2 Evolution. Ph 2+ For both the Co(dmgBF2)2(CH3CN)2 and [Ni(PPh 2 N2 ) 2 ] catalysts, linear free energy relationships have been observed between rate constants for elementary proton transfer steps and acid pKa in acetonitrile, with larger rate constants observed for stronger acids. Plots of log(k) vs pKa for these catalysts have slopes that span −0.7 to −0.97. The physical interpretation of these correlations is still under investigation, but the differences likely reflect the nature of the proton transfer transition state. As such, differences between the nucleophilic species reacting in these proton transfer steps are reflected in these relationships. As the thermodynamic potential for reduction of an acid to hydrogen in acetonitrile depends linearly on the pKa of the acid for structurally comparable acids,69,153 the larger rate constants observed for stronger acids also correspond to larger overpotentials for catalytic hydrogen production. As such, the linear free energy relationships observed provide an experimentally based visual representation of the catalytic activity in different conditions. These relationships could prove useful in benchmarking catalysts, and illustrate the power of detailed kinetic and mechanistic studies.

Figure 12. Example of a total catalysis waveform for the electrochemical hydrogen evolution with 5 mM catalyst Co(dmgBF2)2(CH3CN)2 and 5 mM 4-cyanoanilinium in acetonitrile at 100 mV/s with 0.25 M [Bu4N][PF6]. All of the acid is consumed near the reaction layer, resulting in the appearance of initial catalytic current followed by the CoII/I redox wave.

when the kinetics of catalysis are fast enough such that prior to the primary redox wave the small amount of reduced catalyst (as governed by the Nernst equation) consumes all of the available substrate to produce product and regenerate the catalyst. Consequently, the normal redox wave of the catalyst is observed after the catalytic peak. Extracting accurate kinetic information from the catalytic peak position in this case is possible but is significantly restricted to a small parameter range; therefore, more versatile methods like foot-of-the-wave analysis are generally preferred.24,71 Assignment of a feature as total catalysis can be confirmed via observation of the suspected catalytic peak location as a function of acid concentration, catalyst concentration, and scan rate.71 4.2. Kinetic Control over Peak Location. The mechanism 2+ Ph for the first elementary steps for the [Ni(PPh catalyst 2 N2 )2] discussed above was probed by cyclic voltammetry. While Ph 2+ [Ni(PPh is a catalyst for hydrogen evolution, catalytic 2 N2 )2] turnover can be shut off by the addition of excess conjugate base, 3655


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Scheme 3. NiP2S2

Ph 2+ Figure 13. (A) Titration of a solution of 1 mM [Ni(PPh 2 N2 )2] and 1 M aniline with increasing concentrations of anilinium results in a positive kinetic shift of the reduction wave. (B) Plotting the shift versus log[anilinium] gives a line with a slope of 33 mV/decade, consistent with the theoretical slope of 30 mV/decade for an EC mechanism. Data taken from ref 24.

positive shift of the NiI/0 reduction event due to the subsequent chemical step: protonation of the Ni0 species. Plotting log[BH+] vs the difference in peak potential yields a slope of 33 mV/ decade. This observation supports a stepwise bimolecular EC reaction with initial Nernstian electron transfer to the Ni species followed by irreversible reaction with solution acid BH+; the peak location potential as a function of [BH+] can be described as25,154 Epeak = EI0/0 −

+ RT RT ⎛ k1[BH ]RT ⎞ (0.78) + ln⎜ ⎟ F 2F ⎝ Fυ ⎠

(12)

E0I/0

where is the potential of electron transfer in the absence of BH+, R is the gas constant, F is the Faraday constant, T temperature, k1 is the rate of reaction of BH+ with reduced catalyst, and υ is the scan rate. Analysis of the resulting shift vs [anilinium] permitted evaluation of the protonation rate constant as 1.2 × 106 M−1 s−1.24 Importantly, the slope of plots of log([BH+]) vs the difference in peak potential is predicted by eq 12 to be 30 mV/decadea close match for the experimentally observed slope of 33 mV/decade. Analysis of experimental plots of log([acid]) vs the difference in peak potential provides an experimental means of testing if an EC mechanism is operative. As the rate of the chemical step directly corresponds to the magnitude of the peak shift, this type of potential shift is often referred to as a kinetic shift. 4.3. Thermodynamic Control of Peak Location. The last example discussed here is that of thermodynamically controlled PCET. For a solution of a nickel bisphosphine dithiolate complex (NiP2S2, Scheme 3) with stoichiometric triethylammonium, reduction occurs at potentials much more positive and the magnitude of current passed doubles (Figure 14A,B).11 Assuming no nickel−nickel interactions, two electrons and one proton are added to each nickel molecule in this redox event. As two-electron transfer is not observed in the absence of acid, the

Figure 14. Analysis of reactivity of NiP2S2 with triethylammonium in acetonitrile; data taken from reference.11 (A) Cyclic voltammograms at 100 mV/s of solutions of NiP2S2 with and without one equivalent of triethylammonium. (B) Theoretical peak shift analysis based on a stepwise ECE mechanism compared with real peak shift; current integrations demonstrating two electron and one proton reactivity. (C) Observation of a kinetic isotope effect upon addition of either 0.24 M H2O or D2O to a solution of NiP2S2 with substoichiometric triethylammonium.

stepwise EEC mechanism was ruled out. Three possible mechanisms were then consideredECE, CEE, or the concerted (EC)E. No protonation or association between the nickel complex and acid was detected by either UV−vis or 1H NMR spectroscopy, suggesting that a CEE mechanism was not operative. A ECE mechanism with an irreversible chemical step could result in a large kinetic shift of the reduction event, especially with a fast kinetic step. A kinetic shift was observed for the second case 2+ Ph study above with [Ni(PPh 2 N2 )2] . Digital simulations of an ECE mechanism demonstrated that the magnitude of a reasonable kinetic shift was lower than that observed experimentally (Figure 3656


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14B). This left the (EC)E mechanism as a plausible candidate where the first proton and electron transfer together at a unique potentialin other words, a thermodynamic effect. In the literature, evidence for electrochemical CPET has been supported by observation of a kinetic isotope effect where a small amount of a protic solvent in the proteo or deutero form was added to the solution.20,155−159 Comparison of the cyclic voltammograms of solutions of NiP2S2 with either H2O or D2O showed a small but clear KIE (Figure 14C), supporting a (EC)E mechanism. It should be noted that while observation of an electrochemical KIE supports a concerted mechanism, the lack of a KIE does not rule out a concerted mechanism.

5. OUTLOOK: FUTURE DIRECTIONS FOR THE ELUCIDATION OF PCET PROCESSES IN MOLECULAR CATALYSTS It is our hope that with increased accessibility to tools for studying PCET processes, a greater research effort will be directed at investigating the PCET pathways involved in the catalytic transformations associated with fuel production. Both electrochemical and spectroscopic methods were presented, with a focus on applications of these methods and potential pitfalls. The case studies briefly discussed above illustrate the range of PCET mechanisms thus far explored by our group. Observation of apparent CPET in the reactivity of NiP2S2 has provided interesting directions for our future researchcan CPET be intentionally promoted by varying catalyst electronics or acid strength? As the CPET pathway circumvents high energy intermediates, inducing CPET may result in catalysts with decreased required operating potentials and thus higher efficiencies. Current work seeks to gain control over PCET reaction mechanisms through experimental parameters.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Author Contributions ‡

These authors contributed equally (N.E. and B.D.M).

Funding

This work was supported by the University of North Carolina at Chapel Hill and the David and Lucile Packard Foundation. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the University of North Carolina at Chapel Hill. J.L.D. acknowledges support from a Packard Fellowship for Science and Engineering.

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REFERENCES

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