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Thermodynamic Hydricity of Transition Metal Hydrides Eric S. Wiedner,† Matthew B. Chambers,‡ Catherine L. Pitman,‡ R. Morris Bullock,† Alexander J. M. Miller,*,‡ and Aaron M. Appel*,† †

Pacific Northwest National Laboratory, Richland, Washington 99352, United States Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-3290, United States

‡

S Supporting Information *

ABSTRACT: Transition metal hydrides play a critical role in stoichiometric and catalytic transformations. Knowledge of free energies for cleaving metal hydride bonds enables the prediction of chemical reactivity, such as for the bond-forming and bondbreaking events that occur in a catalytic reaction. Thermodynamic hydricity is the free energy required to cleave an M−H bond to generate a hydride ion (H−). Three primary methods have been developed for hydricity determination: the hydride transfer method establishes hydride transfer equilibrium with a hydride donor/acceptor pair of known hydricity, the H2 heterolysis method involves measuring the equilibrium of heterolytic cleavage of H2 in the presence of a base, and the potential−pKa method considers stepwise transfer of a proton and two electrons to give a net hydride transfer. Using these methods, over 100 thermodynamic hydricity values for transition metal hydrides have been determined in acetonitrile or water. In acetonitrile, the hydricity of metal hydrides spans a range of more than 50 kcal/mol. Methods for using hydricity values to predict chemical reactivity are also discussed, including organic transformations, the reduction of CO2, and the production and oxidation of hydrogen.

CONTENTS 1. Introduction 1.1. Hydricity of Transition Metal Hydrides 1.2. Nomenclature and Conventions 1.3. Scope of the Review 2. Thermochemical Cycles for Determining Hydricities 2.1. Conversions 2.2. Determining Hydricity Values Relative to Reference Hydrides 2.3. Determining Hydricity Values Relative to H2 2.4. Determining Hydricity Values Using Electrochemical Potentials and Acidity 2.5. Alternate Methods To Measure Hydricity 3. Hydricity Values for Metal Hydrides 3.1. Hydricity Values in Acetonitrile 3.2. Hydricity Values in Water 3.3. Comprehensive Thermochemical Schemes 4. Applications of Thermodynamic Transition Metal Hydricity for Small Molecule Activation 4.1. Introduction: Using Hydricity Values To Interpret and Predict Hydride Transfer Reactions 4.2. Hydrogen 4.3. Carbon Dioxide 4.4. Boranes 4.5. Nicotinamides and Pyridinium Derivatives 5. Applications of Thermodynamic Hydricity Involving Metal−Ligand Cooperation 5.1. Introduction © 2016 American Chemical Society

5.2. Hydride Transfer to Substrates Activated by Metal Coordination 5.3. Hydride Transfer Involving Ligands That Act as Hydride Shuttles 5.4. Separated Hydride Transfer Involving Proton-Coupled Electron Transfer 6. Summary and Outlook Associated Content Supporting Information Author Information Corresponding Authors Notes Biographies Acknowledgments References

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1. INTRODUCTION 1.1. Hydricity of Transition Metal Hydrides

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Metal hydrides play a key role in a wide range of chemical processes, from organometallic catalysis to energy conversion and hydrogen storage applications. For example, hydrogenation catalysis involves H−H bond cleavage to generate a metal hydride intermediate that facilitates hydride transfer to an unsaturated substrate.1,2 Metal hydride intermediates are also

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1.2. Nomenclature and Conventions

prevalent in processes that involve C−H bond activation, such as transfer hydrogenation,1,3−6 olefin isomerization,7,8 and C−H functionalization.9−12 In energy conversion schemes, electrochemical reduction and protonation of catalysts often generates metal hydrides that can undergo subsequent hydride transfer reactions with proton sources or carbon dioxide to generate H213−15 or carbon-based fuels, respectively.14,16,17 Many catalysts capable of reversible H2 storage proceed via metal hydride intermediates,18−21 and metal hydrides have also been investigated as bulk H2 storage and transport materials.22,23 Understanding how changes in steric and electronic properties influence hydride transfer is critical in the design of improved stoichiometric and catalytic reactions involving metal hydrides. The hydricity of a metal hydride provides an experimentally measurable parameter that can help predict the reactivity of a metal hydride. Hydricity connects fundamental bonding and reactivity studies with catalyst design principles and applications. To the best of our knowledge, the term “hydricity” to describe reactions of metal hydrides was first used in the literature in 1992,24 and two more papers referred to hydricity in 1994.25,26 A stronger linguistic justification might be made for the term “hydridicity” as a parallel to “nucleophilicity”,27,28 or “hydrididity” as a parallel to “acidity”. Nonetheless, hydricity is the most commonly used descriptor of hydride donor ability, and it will be used herein. Early studies of transition metal hydride reactivity focused on kinetic hydricity, in which the relative rates of hydride transfer to a reference substrate are measured for a range of metal hydrides.29−37 These pioneering studies provided early indications of periodic trends: Labinger observed that early metal hydrides reduce ketones faster than late transition metal hydrides,29 and Bullock observed that third-row transition metal hydrides undergo hydride transfer reactions to Ph3C+ faster than first-row analogues.33−36 Kinetic hydricity has inherent limitations, however. As illustrated in eq 1, the rate of a specif ic hydride transfer from a metal hydride (LnMH) to a hydride acceptor (A) is considered. In such reactions, the identity of A will influence the rate.31 Kinetic measurements can also be influenced by the detailed mechanism of hydride transfer, with some hydrides undergoing concerted hydride ion transfer and others proceeding by stepwise electron/hydrogen atom transfer.34 A more general measure of the inherent hydride donor ability of a metal hydride would help alleviate these limitations. LnM−H + A → LnM+ + HA−

(1)

LnM−H → LnM+ + H−

(2)

ΔG°H −

The terminology surrounding metal hydrides warrants a formal introduction, starting with the term “metal hydride” itself. Longstanding conventions of nomenclature dictate that transition metal complexes bonded to hydrogen are called metal hydrides, because hydrogen is more electronegative than most metal ions and is thus assigned a formal negative charge. The formal naming convention certainly does not imply that transition metal hydrides can only undergo hydride transfer reactions: metal hydrides can also undergo proton transfer42,43 and hydrogen atom transfer reactions.44−47 There are even examples in which the same metal hydride can undergo all three of these modes of cleavage48,49 in reactions with different organic substrates, showcasing the remarkable versatility of metal hydride reactivity. As shown in Scheme 1, the metal hydride bond can be cleaved to a proton (H+), a hydrogen atom (H•), or a hydride (H−). Scheme 1. Cleaving LnM−H Bonds to Yield H+, H•, or H−

Thermodynamic parameters have been determined for all three M−H bond-breaking reactions. As a thermodynamic measure of the energy required to break a bond, any of the reactions in Scheme 1 can be considered bond strengths. The heterolytic bond cleavage with proton dissociation is the acidity of the metal hydride, reported here as a pKa value. Metal hydride pKa values span roughly 43 decades of Ka (∼59 kcal/mol) in acetonitrile,50 as detailed in prior reviews42 and elsewhere in this special issue.43 The homolytic bond-dissociation free energy, or BDFE, involves formation of radicals M• and H•. Often the first “bond strength” that comes to mind, metal hydride BDFE values vary over a relatively narrow 24 kcal/mol range in acetonitrile.51 The focus of this Review is hydricity, which involves heterolytic bond cleavage with hydride dissociation. The hundreds of reported thermodynamic hydricity values tabulated in section 3 span an impressive 50 kcal/mol in acetonitrile. Thermodynamic hydricity, ΔG°H−, is defined as the free energy required to remove a hydride anion (H−) from a species, as shown in Scheme 1. Heterolytic cleavage of an M−H bond to generate H− is endergonic, and the magnitude of ΔG°H− indicates how much energy is needed for bond cleavage. Species with large values of ΔG°H− are therefore weak hydride donors, and species with small values of ΔG°H− are strong hydride donors. Metal hydrides that are strong hydride donors can be described as being hydridic. The term thermodynamic hydricity is considered synonymous with hydride donor ability.39,52,53 The hydride acceptor ability is the reverse of the reaction in eq 2, −ΔG°H−. There is a conjugate relationship between between hydride donors and acceptors. The hydride donor in eq 2 is LnM−H, and the conjugate hydride acceptor is LnM+.37 Solvation has a dramatic impact on thermodynamic hydricity. In principle, ΔG°H− could be defined as a gas-phase thermodynamic parameter, but to our knowledge no experimental gasphase transition metal hydricity values have been reported. The majority of hydricity values have been reported in acetonitrile solvent, facilitating comparisons between the compounds tabulated in section 3. In this Review, the solvent is typically specified after the hydricity is given, as in the phrase “ΔG°H− in

The development of methods for the determination of thermodynamic hydricity has enhanced the understanding of metal hydride bonding and reactivity. Pioneering work by DuBois and co-workers provided experimental thermodynamic values (ΔG°H−) for the heterolytic cleavage of an M−H bond, as shown in eq 2.38,39 The reaction in eq 2 is defined independently of a specific hydride acceptor and, therefore, is strongly solventdependent, as discussed below. As a thermodynamic property, ΔG°H− describes the propensity of a metal hydride to release H−, which can define reaction driving forces and predict reactivity. The utility of thermodynamic hydricity values has been documented by the predictive power they have in designing new catalytic reactions.40,41 8656

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single concerted hydride ion transfer, in a two-step sequence involving electron transfer and hydrogen atom transfer, or in a three-step sequence involving proton transfer and two electron transfers. A mechanistic study of hydride transfer to Ph3C+ illustrates how thermodynamic understanding can be paired with kinetic analysis to identify which mechanism is likely operating.34 The second-order rate constant (kH−) for hydride transfer from Mo(Cp)(CO)3H to Ph3C+ in CH2Cl2 at 25 °C was determined to be 380 M−1 s−1.34 A mechanism proceeding by initial one-electron oxidation of Mo(Cp)(CO)3H by Ph3C+ would be unfavorable by ∼0.9 V, producing a metal hydride radical cation with greatly increased acidity and a weakened M−H bond.59 Even an exceedingly fast hydrogen atom transfer step would only provide an overall maximum rate constant estimated to be 10−2 M−1 s−1. Because the observed rate constant exceeds the estimated rate constant by more than a factor of 104, the reaction likely proceeds by concerted hydride ion transfer, rather than the electron transfer/hydrogen atom transfer reaction shown in Scheme 3. A primary deuterium kinetic isotope effect is further

acetonitrile”. Hydricity values have a strong solvent dependence, based on differences in solvation of the metal species and especially the free hydride ion. Hydricity values are only directly comparable when both species are studied in the same solvent medium. Labinger emphasized early on that hydride transfer leaves a vacant coordination site at the metal center, which can be ligated by solvent or other Lewis bases. Measured equilibria would thus conflate the hydricity with a metal−ligand binding affinity. This might make it “virtually impossible to find equilibria from which hydridic character” could be directly measured, especially with electrophilic early transition metal complexes under most active study at the time.54 DuBois took advantage of the unique electronic structure afforded by square-planar d8 coordination complexes to overcome this challenge, ushering in the broader use of thermodynamic hydricity.14,16,38,39,55 A wide range of fivecoordinate, 18e−, d8 metal hydride complexes can be prepared. Hydride transfer leaves a four-coordinate, square-planar, 16e−, d8 metal complex that does not bind solvents or other ligands strongly (eq 3).

Scheme 3. Electron Transfer/Hydrogen Atom Transfer Mechanism

In cases where solvent coordinates to the metal center following hydride transfer, this interaction is considered part of the overall solvation process. It is common practice for hydricity values determined in coordinating solvents to include general solvation effects as well as any metal−solvent binding. In computational models, this would require “explicit” solvation involving the solvent molecule in the calculation, as well as “implicit” solvation involving a dielectric continuum model. The inclusion of solvation in the hydricity value is another reason why comparisons of hydricity values are most meaningful when compared in the same solvent. For cases where a ligand other than solvent binds to the metal center after hydride transfer, the metal−ligand binding affinity will be inherently included in the measured thermodynamic value. The ef fective hydricity, ΔG°H−(Y), is defined as the free energy required to remove a hydride anion (H−) from a species in the presence of a ligand (Y) that subsequently binds the vacant site on the metal. As shown in Scheme 2, the effective hydricity

evidence of a mechanism involving M−H bond breaking in the transition state.35 The phosphine-substituted analogue Mo(Cp)(CO)(dppe)H is more electron-rich and therefore easier to oxidize. As a result of this substitution, a change in mechanism is observed: hydride transfer to Ph3C+ proceeds via initial one-electron oxidation of Mo(Cp)(CO)(dppe)H.60 In some cases, hydride transfer reactions are referred to as “nucleophilic” reactivity,54,61,62 implying a specific mechanism of concerted hydride ion transfer. For example, Kao and Darensbourg described the conversion of n-BuBr to n-butane by a series of anionic metal carbonyl hydrides as nucleophilic attack.63 The nucleophilic character of metal hydrides can be identified without thermodynamic measurements: for example, on the basis of a kinetic analysis, Norton and co-workers proposed that metal hydrides (M−H) react by nucleophilic attack at a rhenium acyl complex, Et(CO)Re(MeCN)(CO)4, to produce the aldehyde EtC(H)O and M−Re(CO)4(MeCN).64 Several other mechanistic studies of hydride transfer have been carried out,37,49 including analysis of the ionic hydrogenation of CO bonds under stoichiometric65,66 and catalytic2,67−72 conditions. In ionic hydrogenation, the key product-forming step involves hydride transfer from a neutral metal hydride to a protonated ketone or aldehyde. Smith, Norton, and Tilset reported a thorough study of the kinetics of the ionic hydrogenation of acetone by HOTf and Mo(Cp)(CO)2(PPh3)H, and their studies provided evidence that the hydride transfer occurred in a single step.73

Scheme 2. Role of Ligand Binding in Hydricity

includes both the hydricity and the metal−ligand binding affinity. If the association energy (ΔG°assoc) can be measured, the thermodynamic hydricity, ΔG°H−, can also be determined. As discussed in section 3.2, effective hydricity is particularly pertinent in water, where hydroxide is inherently present, and other species such as buffers or salts are also present in many cases.56 The situation is similar to changes in effective acidity encountered when homoconjugation occurs after proton loss.57,58 While the thermodynamic driving force for hydride transfer between two species does not depend on the detailed mechanism of hydride transfer, we provide a short mechanistic aside. Hydride transfer from a donor to an acceptor can occur as a 8657

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1.3. Scope of the Review

in which the hydricity value of a metal hydride is determined relative to the hydricity of H2. The other primary approach relies on thermochemical cycles for hydricity determination. Hydricity is a state function, so any pathway of proton and electron transfers that results in heterolytic release of hydride can be used in a thermochemical calculation. As shown in Scheme 4, acidity values, BDFE values,

This Review is focused on the thermodynamic hydricity of transition metal hydrides. Methods for hydricity determination are introduced, and comprehensive tables of transition metal hydricity values provide a definitive compilation of the thermochemical information involving metal hydride thermodynamics. Guidance in utilizing knowledge of hydricity to understand reactivity and aid predictive catalyst design is also provided. Although a number of main group and organic hydrides are discussed in the context of hydricity applications, the scope of this Review is focused on transition metal hydrides. A detailed review of kinetic hydricity studies is also outside the scope of this thermochemistry-focused Review.71 Finally, this Review is focused on experimental hydricity values, with leading examples of the many computational studies of hydricity provided in the appropriate context. Section 2 introduces the various thermochemical approaches for the determination of thermodynamic hydricity values. Using these experimental methods, scores of reported thermodynamic hydricity values are compiled in the tables of section 3, organized by solvent and by location on the periodic table. To our knowledge, this is the first comprehensive collection of thermodynamic hydricity values for transition metal hydrides, complementing prior reviews that were focused on related aspects of hydride transfer reactions.54,61,71 Accuracy and self-consistency were both considered in compiling the tables, which led us to provide new hydricity values in some cases; corrections were made if improved thermodynamic values became available after the original publication, if the choice of standard states differed from typical convention, or if the thermochemical cycle employed constants that have since been updated, such as those related to solvated H+, H•, or H−. Methods for using hydricity values to predict the reactivity of transition metal hydrides with small molecules are discussed in section 4. A noncomprehensive collection of main group and organic hydricity values is included, along with case studies that utilize thermodynamic hydricity to understand and design stoichiometric and catalytic reactions. In section 5, applications of hydricity in the study of hydride transfer reactions involving metal−ligand cooperation are discussed. This section includes cases in which hydride transfer occurs directly to metal-bound organic groups as well as cases of “separated” hydride transfer where electrons are transferred to the metal and a proton is transferred to a ligand site. Section 6 briefly summarizes the state of the field and provides an outlook for thermodynamic hydricity studies.

Scheme 4. Relationship between Acidity (Red), Homolytic Bond Dissociation Free Energy (BDFE, Blue), Thermodynamic Hydricity (Green), and Reduction Potentials (Black for M+/•/− and Gray for H+/•/−)

and reduction potentials can be converted to free energies and combined to determine hydricity. The necessary reduction potentials for H+ and H• (shown in gray in Scheme 4) are solvent-specific “constants” for which estimates are available. The potential−pKa method38 for determining the hydricity of a metal hydride, discussed in section 2.4, utilizes reduction potentials and acidity parameters. This approach provides net hydride transfer thermochemistry based on the individual thermodynamics of proton transfer and two electron transfers. The H2 heterolysis and potential−pKa methods require one or more of the thermodynamic constants that relate H+, H•, H−, and H2, shown in Table 1. The constants in Table 1 provide the Table 1. Thermodynamic Constants for H+, H•, H−, and H2 in Acetonitrile and in Watera ΔG° in MeCN (kcal/mol)

reaction

eq

2H+(solv) + 2e− ⇌ H 2(g)

3.6

H 2(g) ⇌ 2H•(solv)

103.6

105.7

5

53.6

b

52.8

c

6

26.0

b

−18.6

79.6

b

c

8

34.2

9

H+(solv)

−

+e ⇌

H•(solv)

H•(solv)

−

H−(solv)

H+(solv)

2. THERMOCHEMICAL CYCLES FOR DETERMINING HYDRICITIES Experimental methods for thermodynamic hydricity determination are introduced in this section. Thermochemical background, experimental considerations, and common pitfalls are included, providing a foundation for new researchers and context for the data tables of section 3. Three experimental methods are widely used for the determination of thermodynamic hydricity values: “hydride transfer”, “H2 heterolysis”, and “potential−pKa”. The hydride transfer method38 determines a hydricity value by measuring the equilibrium constant for the reaction of a hydride donor with a hydride acceptor, as will be discussed in section 2.2. This approach can provide the relative hydricity of two species, indicating the difference in hydride donor ability between two hydrides. The H2 heterolysis method,39 discussed in section 2.3, is essentially a special case of determining relative hydricity values,

ΔG° in H2O (kcal/mol)

+e ⇌ −

+ 2e ⇌

H 2(g) ⇌

H+(solv)

H−(solv)

+

H−(solv)

b

76.0

0.0

c

34.2

4

c

7

a

These values correspond to a 1 atm standard state for H2 and a 1 M standard state for H+, H•, and H−. bReferenced to the FeCp2+/0 couple in MeCN. cReferenced to normal hydrogen electrode (NHE).

basis for thermochemical cycles that can be used to quantify the favorability of each of the possible M−H bond-cleavage reactions. The values of these constants in acetonitrile74−77 and water78 are shown in Table 1, and the origins of these values have been previously described. Equation 9 is of particular importance as an expression for three different thermodynamic parameters: the free energy to heterolyze H2, the acidity of H2 (given as a free energy), and the hydricity of H2. The values of the constants shown in Table 1 rely on estimates of the free energy of formation of solvated species, and direct experimental observation of the solvated species is not always 8658

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2.2. Determining Hydricity Values Relative to Reference Hydrides

possible. While the precise values of the constants in Table 1 have been debated,76,78−80 the absolute accuracy of the constants is less important than the universal adoption of the same constant values. For example, the free energy of solvation of H− is experimentally challenging to accurately determine,79 so there is substantial uncertainty in the absolute values of the constants for a reaction involving H−. As long as the hydricity values are determined using the same constants for the reactions in Table 1, any error in the constants will cancel, and the relative hydricity differences will be accurate. The hydricity values are estimated to be accurate within ∼1 kcal/mol. Therefore, the resulting collection of hydricity values can be used to predict equilibria and reactivity, as described in section 4.

The hydride transfer method determines the hydricity of a metal hydride by measuring the equilibrium constant for the reaction of a hydride donor of unknown hydricity (MH) reacting with a reference hydride acceptor (A, where HA− has a known hydricity), as illustrated in Scheme 5. NMR spectroscopy and UV−vis spectroscopy are the two most common techniques for measuring the equilibrium constant for eq 12. The equilibrium constant can be converted to a free energy value according to eqs 10 and 11, as shown in eq 12. The free energy of eq 12 is combined with the free energy of eq 13 to yield the hydricity of MH (eq 14). A reliably quantifiable equilibrium constant can typically only be achieved when the two hydride donors have hydricity values within ∼3 kcal/mol of each other. This difference in hydricity is based on the assumption that a 1:10 ratio of the concentrations of two species is readily quantifiable. If equimolar amounts of a hydride donor and a hydride acceptor react to form a 1:10 equilibrium mixture, then Keq is 100, and thereby ΔG° = −1.364 log(Keq) ≈ 2.7 kcal/mol. For spectroscopic measurements, equilibria should be confirmed to be reversible by either (a) measuring each equilibrium constant from both the forward and reverse reactions or (b) measuring the equilibrium constant a second time after shifting the reaction back toward the reactants through addition of one of the products. Either of these approaches provides evidence that the reaction is at least close to equilibrium, or else the reverse reaction would not be observable under similar conditions. It is also important to monitor the reaction long enough to ensure that the relative concentrations of species are not changing. Hydride transfer reactions can take days or weeks to equilibrate, depending on the steric hindrance of the two metal complexes.37,38,52 If equilibrium is established, the uncertainty in the relative hydricity can be minimal. For thermodynamic values within a series, relative accuracy can often be within 0.1 kcal/mol. The absolute accuracy of the hydricity value is limited by the accuracy of the hydricity value that was determined for the reference hydride using the methods from section 2.3 or 2.4. While this Review is focused on experimental hydricity values, computational methods have been used extensively to predict hydricities. The leading computational methods for hydricity determination also use a reference hydride method. Calculations of accurate absolute hydricity values can be challenging because the energy of the free hydride needs to be considered. Benchmarking computations against reference hydricity values using “isodesmic reactions” provides an approach to determining each hydricity with sufficient accuracy to predict chemical equilibria. Examples of hydricity values from computation are increasing84−92 and provide great promise for continued extension beyond the existing experimental results. The hydride transfer method is intuitively simple and is straightforward to study experimentally, but it requires a reference hydride of known hydricity. If one is working in a solvent with few other hydricity values, or in a hydricity range that is not populated with readily available hydride donors, other methods are needed that provide a hydricity value without the need for reference complexes, as discussed in sections 2.3 and 2.4.

2.1. Conversions

Hydricity values are best expressed as solution f ree energies. The experimental methods involved in thermochemical calculations include electrochemical half-wave potentials, pKa values, and spectroscopically determined equilibrium constants. Thermochemical interconversion between potentials or equilibrium constants and free energy can be accomplished using eq 10, where R is the universal gas constant, T is the temperature in Kelvin, K is the equilibrium constant, n is the number of electrons, F is Faraday’s constant, and E° is the standard reduction potential. At 25 °C eq 10 is equivalent to eq 11, which is expressed in units of kcal/mol for a one-electron process. Equation 11 is expressed in the base 10 log form to facilitate the use of the log scale for pKa values. ΔG° = −RT ln(K ) = −nFE°

(10)

ΔG° = −1.364 log(K ) = −23.06E°

(11)

As with all thermodynamic values, the standard states should be considered carefully and kept consistent. The constants expressed in Table 1 for H+, H•, H−, and H2(g) are estimated using a 1 M standard state for H+, H•, and H− and a 1 atm standard state for H2(g). The choice of 1 atm standard for gases provides both experimental convenience and a simple connection to the hydrogen electrode in a specific solvent (eq 4). Electrochemical potentials are reported relative to a reference electrode or to an internal standard. Aqueous reduction potentials are reported relative to the hydrogen electrode in water. Reduction potentials in organic solvents are referenced to the potential of the ferrocenium/ferrocene (FeCp2+/0) couple.81,82 The different electrochemical reference conventions are responsible for eqs 8 and 9 in Table 1 having the same value in water but different values in acetonitrile because the H+/H2 potential is not equal to the ferrocenium/ferrocene potential. The solution free energies describing thermodynamic hydricity include both enthalpic and entropic contributions. In contrast, homolytic M−H bond cleavage is often reported as a bond dissociation enthalpy (BDE) value in the gas phase.83 The difference between homolytic BDE and BDFE values is often approximated as a constant (∼4.8 kcal/mol) that is largely composed of the estimated solvation and entropic contribution of H•, as the similarity in size and charge of the species M−H and M• is expected to result in very little contribution.76 Thermodynamic hydricity values, on the other hand, have substantial changes in solvation and entropy terms, as the heterolytic cleavage leads to a change in charge for the donors and acceptors in the relevant equilibrium expressions.

2.3. Determining Hydricity Values Relative to H2

The H2 heterolysis method for hydricity determination relies on measuring the equilibrium constant for the reaction of a hydride acceptor, a base, and H2 (eq 15 in Scheme 6) to form a metal 8659

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Scheme 5. Determination of Hydricity by Hydride Transfer

Scheme 6. Determination of Hydricity by Heterolysis of H2

Scheme 7. Determination of Hydricity from H2 Addition and a pKa

Scheme 8. Determination of Hydricity from Two E° Values and a pKa

hydride of unknown hydricity (MH).39 The free energy for H2 heterolysis can be combined with the pKa value of the acid (reverse of eq 16) and the free energy for the heterolysis of H2 in the same solvent (eq 17) to yield the hydricity of a metal hydride (eq 18). The H2 heterolysis method is conceptually related to the relative hydricity method, in that the hydricity of a metal hydride is determined relative to the hydricity of H2. The equilibrium constant for eq 15 is usually determined spectroscopically, typically using NMR spectroscopy. Using a 1 atm standard state for H2 is recommended, given both the experimental convenience and that the constants in Table 1 are based on a 1 atm standard state for H2. In addition to the choice of standard state, researchers should verify that the pressure of H2 is unchanged during equilibration. Verifying the pressure of H2 is especially important when performing equilibrium reactions in NMR tubes, which have a small headspace that can lead to a substantial change in the partial pressure of H2. The other key thermochemical parameter in the H2 heterolysis method is the pKa of an organic acid, the conjugate base of which will cooperate with the metal complex to cleave H2. Aqueous pKa values are widely available, but acidity is solvent-specific. A wealth of acidity data is available for the most commonly used acid/base pairs in acetonitrile, tetrahydrofuran (THF), and dimethyl sulfoxide (DMSO).93−97 The effects of homoconjugation of acids with their conjugate bases can also have a large influence on the observed equilibria,98,99 and therefore it is recommended to work at low concentrations and avoid acids with large homoconjugation constants, as previously described.58 The related pKa-H2 method involves the direct measurement of the equilibrium for binding of H2 to the metal, forming either a dihydrogen or dihydride complex (eq 19 in Scheme 7). By separately measuring the pKa value of the metal dihydrogen or dihydride complex (Scheme 7, eq 20), eq 15 is essentially split into the stepwise eqs 19 and 20, which can be combined with

eqs 16 and 17 (rewritten below as eqs 21 and 22). The same concerns for the stability of the H2 pressure and the choice of standard states for eqs 15−18 apply to eqs 19−23. Equation 20 also involves the experimental determination of the acidity of a metal dihydrogen complex, which is most commonly accomplished based on NMR spectroscopy (see section 2.4 for more on acidity determination). 2.4. Determining Hydricity Values Using Electrochemical Potentials and Acidity

The potential-pKa method for hydricity determination of a metal hydride involves measuring the pKa value of MH (Scheme 8, eq 24) and the reduction potential of the conjugate hydride acceptor (eqs 25 and 26).38 Combining these experimental free energies with the free energy for the two-electron reduction potential for reduction of H+ to H− (eq 27) provides the hydricity (eq 28). For eqs 25 and 26, each reduction potential E° is most commonly determined using cyclic voltammetry (CV), a method that can readily identify reversible reductions that reach equilibrium at the electrode surface. The E1/2 value from a CV experiment is a good approximation of E° if the electrochemical wave is reversible.100,101 Other voltammetric techniques such as square-wave voltammetry or differential pulse voltammetry are also occasionally used, especially to resolve closely spaced 1e− reductions. If a reduction is irreversible on the CV time scale, an open circuit potential technique such as redox potentiometry can be used.56 Similarly, a chemical redox titration could theoretically be used to establish the reduction potential of a metal complex. The reduction potentials of the metal complex and H+ must be given relative to the same reference couple (we recommend FeCp2+/0 in organic solvents).81,82 For the values shown in Table 1, the reference potential in aqueous solution is the hydrogen electrode,78 and the reference potential in acetonitrile is the ferrocenium/ferrocene couple.77 8660

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Scheme 9. Determination of Hydricity from a Two-Electron E° Value and a pKa

known hydricity. By measuring the heat of reaction, relative hydricity enthalpies ΔH°H− can be determined. Hydricity values can even be obtained for irreversible reactions.52,77,106 To make valid comparisons to hydricity values that are given as free energies, the entropic contribution must be taken into account. For example, DuBois and co-workers measured both ΔG°H− and ΔH°H− for a single tungsten hydride and then assumed TΔS°H− (3 kcal/mol) would be conserved across a series of structurally related complexes.52 Support for this assumption was found by comparison to semiempirical estimations of TΔS° in electron transfers.107 In addition to traditional thermodynamic methods, kinetic approaches can also be harnessed to determine thermodynamic hydricities. Creutz and co-workers determined the hydricites of two Ru hydrides through an approach-to-equilibrium method with CO2 and formate.37,108 The equilibrium constant for the process is the ratio of the rate constants for the forward and reverse reactions.

For complexes that display a two-electron redox couple, eqs 25 and 26 can be replaced by a single equation, shown as eq 30 in Scheme 9. As with chemical equilibria, it is important to achieve equilibrium in electrochemical experiments. As mentioned earlier, either two reversible 1e− reduction potentials or one reversible 2e− reduction potential is required. Application of these expressions to irreversible electrochemical waves is not advised, unless the source of the irreversibility is understood sufficiently to provide an accurate estimate of E°.102 Electrochemical measurements represent a potentially significant source of uncertainty in hydricity determination. One large source of relative errors stems from possible discrepancies in reference potentials for electrochemical measurements. Different reference electrodes might be used, or different methods of converting from a reference electrode to the recommended FeCp2+/0 reference couple81,82 could be used. Furthermore, relatively small changes in reduction potential lead to significant differences in hydricity, so accurate measurement of E1/2 values is important. These errors can be avoided by adding ferrocene, or another suitable internal reference molecule, to the solution in order to directly calibrate the reference electrode. In addition to reduction potentials, the acidity of the metal hydride also needs to be assessed in the potential−pKa method. The pKa can be determined in several ways. A spectrophotometric titration can be carried out, adding aliquots of an acid of known pKa (or its conjugate base), or changing pH in water, while monitoring changes in UV−vis spectra. NMR spectroscopy methods can be used as well, typically by adding an acid of known acidity (and/or its conjugate base) and measuring the relative ratios of the species involved. The latter method would only provide a single-point measure of the equilibrium constant and should thus be carried out at several relative concentrations, which should all provide the same pKa value. Many other methods for acidity determination are known, which are beyond the scope of this Review.

3. HYDRICITY VALUES FOR METAL HYDRIDES 3.1. Hydricity Values in Acetonitrile

The majority of themodynamic hydricities have been measured in MeCN. One reason that MeCN is a popular solvent for hydricity determination is the large number of reduction potential and pKa values available in that solvent. Chart 1 depicts the structure of the ligands and abbreviations for the complexes. Table 2 lists first-row transition metal hydrides, while Table 3 provides data on second- and third-row transition metal hydrides. Both tables are arranged in order of decreasing ΔG°H− value (the first entries are the least hydridic). Each hydricity calculation was examined for consistency in the choice of standard states, thermochemical constants, and other parameters. Many values reported here vary slightly ( 1 M−1 s−1 (eq 56). Despite the same free energy of reaction (ΔG°rxn = 0 kcal/mol), the analogous reaction of [Ru(terpy)(bpy)H]+ and [Ru(terpy)(bpy-d8)(NCMe)]2+ is immeasurably slow, with an estimate of k < 3 × 10−6 M−1 s−1.37 The examples below demonstrate that thermodynamic hydricity is proving valuable in predicting single-step hydride transfer reactivity under various reaction conditions. These reactions are often integral steps of catalytic cycles, and many of them are relevant to sustainable synthesis of commodity chemicals and fuels. In addition to stoichiometric hydride transfer reactions, examples of catalytic reactions that were developed based on knowledge of hydricity are also included. The examples are not meant to be exhaustive but rather to provide leading examples of how hydricity values can be used to understand and predict reactivity.

by changing the ligands present in solution, as described in section 1. Following reactions are common in multistep catalytic reactions, where an uphill hydride transfer might be followed immediately by a favorable step that drives the overall sequence. For example, reductions of organic carbonyls (aldehydes, ketones, esters, etc.) often involve stepwise (or concerted) transfer of a proton to the carbonyl oxygen and hydride to the carbonyl carbon, and coupling these reactions can enhance reactivity.5,158−161 Metal hydrides with a d6 configuration often bind a ligand at the unsaturated 16e− species that is generated upon hydride transfer. For example, hydride transfer from Fe(Cp*)(CO)2H to [Ph3C][BF4] in dichloromethane solution produces the adduct Fe(Cp*)(CO)2FBF3, in which the BF4− counterion binds to the iron (Scheme 22, eq 54).36 In contrast, hydride transfer from the analogous complex Ru(Cp*)(CO)2H to [Ph3C][BF4] under the same conditions leads to a bridging hydride complex (Scheme 20; eq 55); in this case, Ru(Cp*)(CO)2H exhibits reactivity as a nucleophile, leading to the observed cationic product with a bridging Ru hydride. The solvent can also act as the trapping ligand: when M(Cp*)(CO)2H (M = Fe or Ru) reacts with [Ph3C][BF4] in MeCN, the solvent species [M(Cp*)(CO)2(NCMe)]+ are formed, where the MeCN solvent binds to the metal cation after hydride transfer. The free energy of a hydride transfer reaction provides valuable information on the propensity for a reaction to occur, but thermodynamic favorability does not guarantee that a reaction will occur at an appreciable rate. The kinetic barriers to hydride transfer can be substantial and hard to predict from one group of hydrides to another. Kinetic barriers for such transfers are often evaluated via “self-exchange” reaction rates.

4.2. Hydrogen

A proton is the simplest hydride acceptor, producing dihydrogen upon reaction with a hydride.13 The propensity of a metal hydride to react with a proton and release H2 will depend on (a) the hydricity of H2, (b) the hydricity of the metal hydride, and (c) the acidity of the proton source. The propensity of a metal hydride to release H2 at 1 M H+ can be estimated using the solvent-specific H2 hydricity values (see Table 1 in section 2). In acetonitrile, for example, any hydride with ΔG°H− < 76 kcal/mol is thermodynamically capable of H2 evolution at 1 M H+ in MeCN. This is illustrated in the thermochemical cycle shown in Scheme 23, where ΔG°rxn < 0 indicates favorable H2 release. In water, any metal hydride with ΔG°H− < 34 kcal/mol is thermodynamically capable of H2 release at pH 0. The free energy of H2 release depends strongly on the acid source, as acids of varying acidity will each have a different driving force for reaction with a metal hydride. This variation can be taken into account using the adjusted thermochemical cycle in Scheme 24, which is simply a rearrangement of Scheme 5 that was used to determine hydricity based on the equilibrium heterolysis of H2. For example, the reaction of [Ir(Cp*)(bpy)H]+ (ΔG°H− = 62 kcal/mol) with methanesulfonic acid (pKa = 10.0)162 is predicted to be favorable by 0.4 kcal/mol, while the analogous reaction with triethylammonium (pKa = 18.8)93 would be predicted to be unfavorable by 11.6 kcal/mol in acetonitrile solvent. Indeed, treatment of [Ir(Cp*)(bpy)H]+ with methanesulfonic acid triggers H2 release, whereas no reaction is observed with triethylammonium (eq 65).128 Considering the large number of acidity values that have been compiled in various organic solvents,163,164 H2 release and heterolysis thermodynamics are 8669

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Scheme 23. Free Energy for Protonation of a Metal Hydride

Scheme 24. Free Energy for Protonation of a Metal Hydride

Table 6. Acid Dependence of Hydricity Required to Release H2 acid water phenol acetic acid benzoic acid HNEt3+ pyridinium anilinium (PhNH3+) p-toluenesulfonic (tosylic) acid trifluoromethanelsulfonic (triflic) acid (CF3SO3H; HOTf)

pKa in MeCN

ΔG°H− to release H2 in MeCN (kcal/mol)

ref

38−41 29.1 23.5 21.5 18.8 12.5 10.6 8.6 2.6

20−24 36 44 47 50 59 62 64 72

165 166 97 94 93 93 93 167 94

Figure 3. Tuning of H2 addition for [Ni(PR2NR′2)2]2+ complexes.174

to alter reaction conditions to improve catalysis. For the [Ni(PR2NR′2)2]2+ catalysts shown in Figure 3, the hydricity of the metal hydride must be matched with the acidity of the protonated pendant base. For example, [Ni(PCy2NBn2)2H]+ has ΔG°H− = 60.7 kcal/mol, and the pKa of the pendant acid in the protonated form is 13.4. Using a general form of the thermochemical cycle shown in Scheme 13, this combination will favor H2 addition by 3.1 kcal/mol. On the other hand, [Ni(PPh2NBn2)2H]+ has ΔG°H− = 57.1 kcal/mol and the pendant acid has pKa = 11.8 in the protonated form, which instead favors H2 evolution by 2.7 kcal/mol. Finally, [Ni(PPh2NPh2)2H]2+ has ΔG°H− = 59.0 kcal/mol and the pKa of the pendant acid of the protonated form is estimated as 6, leading to highly favorable H2 release by ∼9 kcal/mol. On the basis of the H2 release/ heterolysis thermochemistry, one would predict that [Ni(PCy2NBn2)2]2+ would be the best H2 oxidation catalyst and that the H2 evolution activity would increase in the order of catalyst [Ni(PCy2NBn2)2]2+ < [Ni(PPh2NBn2)2]2+ < [Ni(PPh2NPh2)2]2+. Precisely this trend is observed.115 Knowledge of hydricity can also help explain changes in mechanism as a function of acid choice. In electrocatalysis, a metal hydride formed at a mild reduction potential may react only with strong acids to release H2. If the chosen acid is so weakly acidic that there is no thermodynamic driving force for H2 release, the hydride intermediate must undergo an additional reduction process at more negative potentials, which generates a more hydridic intermediate that is then able to react with the weak acid.175−178 The Co(II) catalyst [Co(CpC5F4N)(PtBu2NPh2)]+ offers an illustrative example, where initial reduction and protonation forms a catalytically inactive Co(III)−H species. The hydricity of the Co(III)−H was determined to be 73 kcal/mol, putting this species among the poorest hydride donors reported.105 Under the conditions of electrocatalysis, H2 evolution by

readily predicted for any species with a well-defined hydricity value. Table 6 lists several organic acids along with the thermodynamic hydricity that a metal hydride would need to favor H2 evolution. Strongly hydridic metal hydride intermediates will readily release H2, whereas weakly hydridic metal hydrides will not be thermodynamically capable of releasing H2, favoring instead H2 splitting to form a stable metal hydride. Homoconjugation can complicate predictions by perturbing the effective acidity, as described in section 2.3.

Electrocatalytic H2 evolution, in particular, relies on hydride transfer to a proton source for solar fuel synthesis.13,146,168−170 If the hydricity of a catalytic intermediate is well-defined, one can predict which acid will be needed to ensure favorable H2 release. In an informative case study, a family of Ni diphosphine complexes has evolved over the last 15 years from H2 oxidation catalysts with low activity (1 000 000 s−1).171−173 The addition of pendant bases was found to play an important role in accelerating the kinetics of proton transfer to and from the metal center. In general the pendant amines do not dramatically alter the hydricities of the catalysts but enable proton transfer reactions and positioning of protons, both of which are critical for the observed catalytic activity. A detailed thermochemical understanding of several catalytic intermediates provided insight into reactivity trends and how 8670

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quantitative release of H2 from solutions of [Ir(Cp*)(bpy)H]+ and pyridinium or acetic acid irradiated with 460 nm light. Table 6 indicates that a hydride must have ΔG°H− < 44 kcal/mol in order to react with acetic acid to release H2, suggesting that light absorption is generating an intermediate that is at least 18 kcal/mol more hydridic than the ground state of [Ir(Cp*)(bpy)H]+. Photochemical H2 release is also observed in neutral water, which enabled the development of a molecular photoelectrocatalyst for proton reduction in water.135 In the dark, [Ir(Cp*)(bpy)H]+ is a poor H2 evolution electrocatalyst, consistent with the experimentally determined aqueous hydricity value.128 Upon visible light illumination, however, electrocatalysis proceeds smoothly in neutral water with rates in the range of 0.1 s−1. Catalytic activity was observed even at zero electrochemical overpotential. Although the detailed mechanism of light-triggered H2 evolution is unknown, the influence of light absorption on the thermodynamic hydricity has been estimated by extending a ground state thermochemical scheme to include the excited state intermediates (Scheme 25).128 Light absorption by a ground state metal hydride generates an excited state intermediate that has been previously characterized and implicated in catalysis.128,184,185 Excited state hydride transfer coupled with relaxation to the ground state would provide an “excited state hydricity” (ΔGoH−*). The excited state hydricity, ΔGoH−*, was experimentally determined based on the energy difference between the excited state and the ground state (estimated as E00, the energy difference between the lowest vibrational energy levels, respectively, for the ground state and excited state). The free energies of the E00 value, the acidity of the metal hydride, the reduction potential of the conjugate hydride acceptor, and the free energy of the H+/H− couple (eq 72) can be combined to give ΔGoH−*, as illustrated in Scheme 26 (the factor of 1/350 converts the units of wavenumbers to kcal/mol). The excited state hydricity of [Ir(Cp*)(bpy)H]+ was thus determined as ΔGoH−* = 14 kcal/mol, one of the most potent thermodynamic hydride donors available. While still a developing area, the ability to modulate hydricity through the absorption of light represents a promising strategy for generating reactive metal hydride intermediates. Hydrogen evolution reactions have been emphasized in this section, but hydricity is also critical for reactions in which hydrogen is cleaved to form metal hydrides. Many hydrogenation catalysts generate metal hydride intermediates by heterolytic cleavage of H2, and this reaction step will depend strongly on the properties of the hydride acceptor and proton acceptor. Hydricity can also play a role in the subsequent reduction of a substrate, and the free energy of H2 heterolysis (favored by a strong hydride acceptor ability of the catalyst) must be balanced with the free energy of hydride transfer to substrate (favored by a strong hydride donor ability of the catalyst). Hydrogenations are often carried out at high pressure, which will shift the H2

protonation of the Co(III)−H is endergonic by >13 kcal/mol using p-anisidinium as the acid (pKa = 11.86 in MeCN).93 Only extremely strong acids (pKa < 2.2 in MeCN) would be capable of reacting with the Co(III)−H to release H2. On the other hand, reduction of the Co(III)−H by one electron to form a Co(II)−H complex leads to a much more hydridic species, with ΔG°H− = 41.9 kcal/mol. Upon formation of the Co(II)−H, protonation by p-anisidinium to release H2 is exergonic by almost 18 kcal/mol. The electrocatalytic response of the Co catalyst is fully consistent with the hydricity predictions: no catalytic current enhancement is observed at the potential at which the Co(III)−H is formed, and upon subsequent reduction to the Co(II)−H a large catalytic response is observed with a turnover frequency of ∼350 s−1. It should be noted that simply increasing hydricity will not universally improve electrocatalysis because of a strong correlation between ΔG°H− and E° (see Figure 1 in section 3). Boosting the hydricity of an electrocatalyst will be balanced by a negative shift in the reduction potential, leading to a trade-off that results in higher activity at higher overpotential.99,179 If H2 release is not involved in the turnover-limiting process, tuning hydricity will not have much impact on the rate of the reaction. Light absorption can also modulate hydride transfer reactivity for H2 evolution. The photochemistry of metal hydrides has received significant attention, especially in the context of solar water splitting.180−183 A recent thermochemistry-guided investigation found that [Ir(Cp*)(bpy)H]+ releases H2 when illuminated with visible light.128 As mentioned above, [Cp*Ir(bpy)H]+ readily releases H2 upon addition of methanesulfonic acid (pKa = 10),162 but no H2 release is observed upon addition of pyridinium salts (pKa = 12.5)93 or acetic acid (pKa = 23.5).94 These observations are in accord with the hydricity of [Ir(Cp*)(bpy)H]+ in acetonitrile, ΔG°H− = 62 kcal/mol. Under illumination, however, a striking enhancement in reactivity is observed, with rapid and Scheme 25. Thermochemical Scheme for [Ir(Cp*)(bpy)H]+, Including Photoexcitation, in MeCN

Scheme 26. Determination of an Excited State Hydricity in MeCN

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hydride transfer to CO2 has also been investigated. The Ni hydride [Ni(dmpe)2H]+ (ΔG°H− = 50.8 kcal/mol) is unable to reduce CO2 under typical conditions, but hydride transfer occurs in the presence of the Lewis acid BEt3. This reaction is rationalized in terms of stabilization of formate as a borane adduct, which increases the effective hydride acceptor ability of CO2 (Scheme 27).154

activation equilibrium toward production of a metal hydride according to eq 53. For example, heterolytic cleavage of H2 by complex [Ir(Cp*)(bpy)(H2O)]2+ (ΔG°H−(H2O) = 31.5 kcal/mol) under 1 atm H2 pressure is predicted to yield only 1.1% conversion to the hydride in pH 0 water. Under 100 atm of H2, however, 51% conversion to the metal hydride would be expected. This is generally consistent with the experimental observation that the rate of hydrogenation of neat acetic acid by [Ir(Cp*)(bpy)(H2O)]2+ levels off around 50 atm.161

Scheme 27. Solvent and Additive Dependence of Hydride Transfer from [Ni(dmpe)2H]+ to CO2

4.3. Carbon Dioxide

Carbon dioxide reduction has been considered in the context of hydricity by several groups, motivated by alternative energy schemes involving fuel synthesis.17,186−194 The hydricity of the formate anion in acetonitrile has been estimated as ΔG°H− = 44 kcal/mol, based on a hydride transfer equilibrium between CO2 and [Pt(depe)2H]+.16,154 Thus, hydrides with ΔG°H− < 44 kcal/mol are poised to produce formate in the presence of 1 atm CO2. This was demonstrated by DuBois and co-workers for [Pt(diphosphine)2H]2+ complexes: the weak hydride donor [Pt(dmpp)2H]+ (ΔG°H− = 51 kcal/mol) does not react with CO2, while the stronger hydride donor [Pt(dmpe)2H]+ (ΔG°H− = 42 kcal/mol) readily produces formate.16,39 Similarly, Creutz and co-workers demonstrated that [Ru(terpy)(bpy)H]+ (ΔG°H− = 39 kcal/mol) is sufficiently hydridic in MeCN to reduce CO2 to formate. Hydride transfer from Ru to CO2 proceeds readily (k = 1.8 × 10−2 M−1 s−1) to afford the formato complex [Ru(terpy)(bpy)(O2CH)]+, in which binding of formate to the product provides an additional driving force.37,195 Several approaches have been developed for promoting hydride transfer to CO2, beyond simply preparing a more hydridic metal hydride. The most commonly employed tactic involves high pressures of CO2 (and/or H2). Carbon dioxide hydrogenation is often carried out under elevated pressures of both H2 and CO2.188 The hydricity of formate is defined at a standard state of 1 atm CO2, so the elevated pressures are expected to influence hydride transfer reactions involving CO2. Increasing the initial CO2 pressure in eq 72 will lead to more complete hydride transfer from the hydride donor, proportionately to the change in pressure. ⎛ [M+] [HCO −] ⎞ 2 ⎟⎟ ΔG°rxn = −RT ln⎜⎜ ⎝ [MH] PCO2 ⎠

Changes in solvent have a large influence on hydride transfers involving CO2. The hydricity of formate is solvent-dependent, with the value in water (ΔG°H− = 24.1 kcal/mol)78 roughly 20 kcal/mol smaller than that in MeCN or DMSO (ΔG°H− = 42 kcal/mol).118,197 Because the hydricity of formate does not decrease as dramatically as most transition metal hydrides in moving from organic solvents to water, CO2 becomes a better hydride acceptor in water relative to many transition metals. The different solvent dependence is attributed to differences in solvation of the formate anion relative to transition metal hydride complexes.37,145 Solvent-dependent hydricities lead to a situation where some metal hydrides are thermodynamically incapable of reducing CO2 in acetonitrile but readily transfer hydride to CO2 in water.118,197 This phenomenon has been explored using [Ni(dmpe)2H]+, which does not readily react with CO2 in MeCN solvent but produces formate in water (Scheme 27).145 The family of Fe cluster electrocatalysts developed by Berben and co-workers offers an elegant example of tuning CO2 catalysis through the choice of acid and the solvent (Chart 3). Electrocatalytic CO2 reduction hinges on control of selectivity. A hydride intermediate can either react with CO2 to produce formate or react with a proton source to generate H2.186,193 Electrolysis of CO2-saturated acetonitrile solutions containing

(72)

Carbon dioxide reduction is most often performed in the presence of base. The thermodynamic reasons for basic reaction conditions are manifold. First, the overall thermodynamics of H2 and CO2 reacting to form HCO2− and H+ become more favorable as more base is present to react with the released proton.78,145 Furthermore, precipitation of formate salts from organic solvents can help drive reactions.196 Individual hydride transfer steps can also be influenced by solution basicity. Hydride transfers involving formate are independent of pH, in contrast to reactions involving H2 that release H+. The hydricity of most metal hydrides will also be unaffected by pH, so the driving force for hydride transfer to CO2 will be constant across the pH range. In water, some hydride donors do have pH-dependent reactivity, however. For example, under basic conditions, the product of hydride transfer from [Ir(Cp*)(bpy-COO)H]− is the hydroxide complex [Ir(Cp*)(bpy-COO)(OH)]−. The concentration of hydroxide is governed by solution pH, so hydride transfer to CO2 is almost 10 kcal/mol uphill at pH 7, but favorable by ∼1 kcal/mol at pH 14.56 The use of additives to shift the thermodynamics of

Chart 3. Structural Representations of Fe4 Clusters

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the Fe catalyst [Fe4N(CO)12]− produced H2 as the sole product when toluenesulfonic acid was the proton source, but some formate was detected when the weaker proton source benzoic acid was employed.198 A high selectivity for formate (∼95% Faradaic efficiency) was observed in both a 95/5 MeCN/H2O solvent mixture and in water above pH 5.121 The high formate selectivity of [Fe4N(CO)12H]− under aqueous conditions prompted a careful thermochemical study.121,123 The hydricity of [Fe4N(CO)12H]− was determined in both acetonitrile (ΔG°H− = 49 kcal/mol) and water (ΔG°H− = 15.5 kcal/mol).121 The Fe hydride is a moderately strong hydride donor in acetonitrile, but transfer to CO2 is not predicted to be favorable (ΔG°rxn = +5 kcal/mol). On the other hand, as one of the most hydridic species known in water, the very same Fe cluster is predicted to easily reduce CO2 to formate (ΔG°rxn = −9 kcal/mol).123 The hydricity of formate in the 95/5 MeCN/ H2O solvent mixture has not been measured. However, the acidity of formic acid starts to shift toward the aqueous value at ∼5% water content,199 so the hydricity of formate is likely shifting in this range as well. The increased selectivity for formate in aqueous solvent mixtures (or pure water) is attributed to the change in driving force for hydride transfer that results from this shift in formate hydricity.121,123 While the hydricity of formate is pH-independent, the driving force for protonating a hydride decreases with increasing pH, and therefore the thermodynamic bias for CO2 reduction relative to H+ reduction also increases as a function of pH. A hydride with ΔG°H− = 15.5 kcal/mol will maintain a constant −9 kcal/mol driving force for reduction of CO2, while the driving force for proton reduction shifts steadily until becoming thermoneutral near pH 14. Kinetic factors are surely also influencing this selective reduction, but the thermochemical analysis supports the experimental observations. Similar thermochemical analyses have now been extended to related clusters (Chart 3). The phosphine-substituted variant [Fe4N(CO)11(PPh3)H]− is a stronger hydride donor (ΔG°H− = 46 kcal/mol in MeCN) than [Fe4N(CO)12H]−,122 and the carbide-bridged cluster [Fe4C(CO)12H]2− is even more hydridic (ΔG°H− = 44 kcal/mol in MeCN), consistent with the increased negative charge on the complex.123 The increased hydricity of the carbide-bridged cluster is identified as a key reason for a loss of selectivity: only H2 evolution is observed with this cluster. The nitride-bridged cluster remains one of the most selective catalysts for electrocatalytic CO2 reduction to formate, with the best performance observed in neutral water.123 In another example, bis(diphosphine) cobalt complexes were developed into exceptional catalysts for CO2 reduction by leveraging knowledge of hydricity and thermochemical trends of metal hydrides.200 This system required a cobalt hydride of sufficient hydricity to reduce CO2 to formate. The hydride Co(dmpe)2H (ΔG°H− estimated at 36 kcal/mol in MeCN)134 was predicted to have ample driving force for hydride transfer to CO2 (formate ΔG°H− = 44 kcal/mol in MeCN).16,154 Although the hydricity of the Co(I)−H has not been experimentally determined, the estimate of 36 kcal/mol is based on both DFT calculations85 and known ligand trends in similar systems (see section 3).134 The initial product of H2 cleavage is the Co(III) dihydride [Co(dmpe)2(H)2]+, which is not hydridic enough to reduce CO2 to formate. Deprotonation of Co(III) dihydride (pKa = 33.7 according to DFT calculations)84 with an organic base generates the more potent Co(I) hydride and completes the catalytic cycle of Scheme 28. Specifically, catalytic studies were performed using Verkade’s base (2,8,9-triisopropyl-2,5,8,9tetraaza-1-phosphabicyclo[3,3,3]undecane), for which the

Scheme 28. Proposed Catalytic Cycle for Hydrogenation of CO2 to Formate Using Co(dmpe)2H

conjugate acid in MeCN has a pKa of 33.6. Using this base, extraordinary rates of catalytic CO2 hydrogenation are obtained, with turnover frequencies (TOFs) up to 74 000 h−1 at ambient temperature and 20 atm of a 1:1 mixture of H2 and CO2.201 The importance of matching the pKa of the conjugate acid of the base with the pKa of the metal hydride intermediate is emphasized by the sharp drop-off in activity when weaker bases are employed. When a phosphazene base (tert-butyliminotris(dimethylamino)phosphorane) that is roughly 5 pKa units less basic than Verkade’s base was employed, the activity dropped by a factor of 60. This system also challenges the prevailing notion that precious metal catalysts are inherently more active than firstrow congeners. The analogous Rh hydride catalyst Rh(dmpe)2H is about an order of magnitude less active than Co(dmpe)2H under the same conditions. The poorer performance of the Rh catalyst is not attributed to a change in hydricityRh(dmpe)2H is almost 13 kcal/mol more hydridic than Co(dmpe)2H131but rather to the decreased acidity of [Rh(dmpe)2(H)2]+ (pKa = 36.7),131 which prevents Verkade’s base from readily deprotonating this species to generate the key hydride. Knowing the structure of catalytic hydride intermediates can be critical to understanding the role of hydricity in the reaction mechanism. In the Ir-catalyzed photochemical reduction of CO2 to CO, two hydride intermediates with two different geometric isomers have been isolated, and each species has a distinct hydricity. Following the report by Ishitani and co-workers that [Ir(terpy)(ppy)(Cl)]+ is a promising photocatalyst for CO2 reduction with good turnover numbers and quantum yields for CO production,202 a team led by Ertem and Fujita isolated the two hydrides C-trans-[Ir(terpy)(ppy)H]+ and N-trans-[Ir(terpy)(ppy)H]+ (where the phenyl and pyridyl groups of ppy are trans to the hydride, respectively, as in Chart 4).106 Chart 4. Structural Representations of cis and trans Isomers of [Ir(terpy)(ppy)H]+

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weak bases. On the other hand, weaker Lewis acids will be poor hydride acceptors, necessitating the use of a stronger base. The initial reports of H2 cleavage by FLPs composed of B(C6F5)3 and trialkylphosphines208,209 relied on substantial thermodynamic driving force. Using reaction calorimetry, Autrey and co-workers measured a reaction enthalpy of −31.4 kcal/mol for H2 splitting by B(C6F5)3 and P(tBu)3 in bromobenzene solution.210 Comparison of the experimental enthalpy value with the calculated free energy (ΔG° = −14.7 kcal/mol)207 suggests a large negative entropy of reaction (ΔS° ≈ −56 cal mol−1 K−1), as expected for a termolecular reaction that produces an ion-pair in a nonpolar solvent. These H2 cleavage reactions are also consistent with the calculated hydricity of [(C6F5)3BH]− (ΔG°H− ≈ 65 kcal/mol in acetonitrile).92 For example, a hypothetical FLP composed of B(C6F5)3 and PMe3 (conjugate acid pKa = 15.5, ΔG°H+ = 21.1 kcal/mol)211 favors H2 splitting to form the salt [HPMe3][(C6F5)3BH] by 10 kcal/mol (in MeCN solution). This ample driving force helps to explain the initial observations of this new mode of metal-free H2 cleavage. In recent breakthroughs in FLP-catalyzed ketone and aldehyde hydrogenation, boranes of similar Lewis acidity have been paired with very weak bases such as THF or the ketone substrate itself.212−214 In this case, because the proton acceptor is such a weak base, the H2 heterolysis thermodynamics are unfavorable at 1 atm, perhaps explaining the high pressures of H2 required for observing efficient retion rates. In analogy to proton transfer from acids, homoconjugation of borohydrides is observable and certainly affects thermodynamic considerations of hydride transfers.58,215−219 In the case of the equilibrium between Rh(dmpe)2H and BEt3, the observed 1 H NMR chemical shift was found to vary over a ppm based on the ratio of Rh(dmpe)2H/BEt3.155 It was suggested that the B−H resonance was controlled by the homoconjugation equilibrium between [HBEt3]−, BEt3, and [Et3B−H−BEt3]−, as shown in eq 73.

The isomer with pyridine trans to the hydride showed no reactivity with CO2, whereas the isomer with phenyl trans to the hydride immediately reacted with CO2 to produce formate. The hydricity of the more stable isomer N-trans-[Ir(terpy)(ppy)H]+ (ΔG°H− = 57) was determined in MeCN solvent, confirming that this species is far too weak a hydride donor to reduce CO2. The other isomer was too reactive for experimental hydricity determination, and DFT calculations suggested that the phenyl trans to the hydride leads to a large increase in hydricity relative to the N-trans isomer. The differences in hydride donor ability of the two hydride isomers, along with a photochemical process for interconverting between the isomers, is proposed to be responsible for the high selectivity toward CO. This system illustrates how altering the electron donor strength of the ligand trans to the hydride can be a powerful method for tuning hydricity. 4.4. Boranes

Boranes have received substantial attention in hydride transfer studies, in large part due to the prevalence of borohydrides as reductants in organic chemistry and as candidates for hydrogen storage applications. In spite of such widespread utility, borohydrides present challenges in hydricity determination. The propensity of borohydrides to react by concerted hydride ion transfers typically rules out potential-pKa thermochemical cycles, and even hydride transfer equilibria can be complicated by side reactions.134,203 Hydride transfer equilibria were used to estimate the hydricity of the “Super-Hydride” [Et3BH]− (ΔG°H− = 26 kcal/mol in MeCN).104,131,155 Rhodium(I) hydrides of the type Rh(diphosphine)2H were found to establish equilibrium with triethylborane.104,131,155 One of the strongest hydride donors among known transition metal hydrides, Rh(dmpe)2H (ΔG°H− = 26.4 kcal/mol in MeCN), transfers a hydride to BEt3 (in THF solution) to produce [Et3BH]−. The reverse reaction was observed, suggesting similar hydricity values for these species. As predicted by thermodynamic hydricity values, a similar equilibrium was observed between BEt3 and Rh(depe)2H (ΔG°H− = 28.1 kcal/mol in MeCN), but with less formation of [Et3BH]−. Complete hydride transfers from Rh(dmpe)2H in THF to a borane were observed with boranes having attenuated hydride acceptor abilities, such as BH3 or B(OPh)3. Calibrated by these experimental studies, Heiden and Lathem recently carried out an extensive computational study that estimates the thermodynamic hydricity values for a wide array of borohydrides.92 Hydrogen cleavage equilibria could also be used to determine borohydride hydricity. So far, this type of heterolytic H2 splitting reactivity has only been observed by boranes in “frustrated Lewis pair” (FLP) chemistry in conjunction with an organic Lewis base (phosphine, amine, ether, etc.).204−206 Understanding the hydricity of boranes has been of paramount interest in the context of predicting the reactivity of FLPs.207 This is analogous to heterolytic H2 cleavage by a transition metal and a base (see section 2.3). In general, stronger Lewis acids will be good hydride acceptors and are able to cleave H2 cooperatively with relatively

Et3B−H− + BEt3 ⇌ [Et3B−H−BEt3]−

(73)

In conditions where Rh(dmpe)2H was in excess relative to BEt3, the B−H resonance is assigned to [HBEt3]−. When an excess of BEt3 was present, however, the homoconjugation product [Et3B−H−BEt3]− was observed. Similar species were also observed in the equilibrium between Pt(dmpe)2H and BEt3.220 While the hydricity of borohydride homoconjugation products remains a largely unexplored field, the observation of such products in equilibrium with transition metal hydrides suggests that the bridging borohydrides are less hydridic than the parent borohydrides. Homoconjugation could potentially be used to drive an otherwise unfavorable hydride transfer from a transition metal to a borane by thermodynamic stabilization of the borohydride product (Scheme 29). The stability of bridging borohydrides is also evident in the hydride acceptor reactivity of 1,8-naphthalenediylbis(dimethylborane), a “hydride sponge” analogous to the commonly used “proton sponge”, 1,8-bis(dimethylamino)naphthalene.221 The extent of the thermodynamic stabilization of homoconjugation

Scheme 29. Stabilization of Et3BH− by Reaction with BEt3

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was explored through hydride transfer equilibria with other borohydrides (eq 77), which showed complete hydride transfer from monomeric borohydrides to 1,8-naphthalenediyl-bis(dimethylborane).221

Hydride transfer between transition metals and boranes is rarely incorporated into functional catalytic systems. One promising example involves the olefin polymerization catalyst [Zr(Cp)2H][HB(C6F5)3], a much more active species than the dihydride Zr(Cp)2(H)2.222 The [Zr(Cp)2H][HB(C6F5)3] catalyst can be prepared via hydride transfer from Zr(Cp)2(H)2 to B(C6F5)3. Extension of this synthesis to in situ activation under catalytic conditions would afford control over the initiation of catalysis. In a related reaction, it has already been demonstrated that treatment of Zr(Cp) 2 HCl with the aforementioned hydride sponge results in hydride transfer to the boranes, affording [Zr(Cp)2Cl]+.221 The design of catalysts that can be activated by hydride transfer would be greatly aided by knowledge of the relative hydricities of a transition metal catalyst and borohydride.

Figure 4. Correlation between ΔG°H− and ΔH°H− for a series of NADH derivatives in acetonitrile.

on the basis of their reported enthalpies (ΔH°H− = 69 kcal/mol)156 and the linear correlation shown in Figure 4. Organic hydride donors can establish hydride transfer equilibrium with transition metals, and such reactions have helped form a basis of observed selectivity in nicotinamide reduction chemistry. Treatment of nicotinamide derivatives with [Rh(Cp*)(phenylpyridine)H]+ or [Ru(terpy)(bpy)H]+ results in the selective generation of the products of a single hydride transfer, a mixture of the 1,4- and 1,6-dihydronicotinamide products.37,226 Both [Rh(Cp*)(phenylpyridine)H]+ (ΔG°H− = 49 kcal/mol in acetonitrile) and [Ru(terpy)(bpy)H]+ (ΔG°H− = 39 kcal/mol in acetonitrile) have hydricity values consistent with favorable reduction of nicotinamide derivatives.37,226 There exists a wide array of reports of hydride transfer to nicotinamide derivatives from various transition metal complexes invoking a metal hydride as a key intermediate, including [Rh(Cp*)(α-diimine)] 2 + , 22 4 , 22 7 −2 3 3 [Ru(η 6 -arene)(αdiimine)]2+,37,234,235 [Ru(terpy)(bpy)]2+,236−241 and [Ru(porphyrinato)]2+.242 Complete thermodynamic analyses are lacking, however, often because the metal hydride species has not been isolated. For example, [Rh(Cp*)(bpy)(solvent)]2+ and related complexes are widely studied for hydride transfer involving nicotinamides, but the proposed hydride intermediate [Rh(Cp*)(bpy)H]+ has never been isolated.227,228,231 Nonetheless, data on hydride transfers to nicotinamides can be rationalized in terms of expected hydricity trends. When a solution of [Rh(Cp*)(bpy)(H2O)]2+ in D2O/d8-THF is treated with formate to presumably generate [Rh(Cp*)(bpy)H]+ in situ, no hydride transfer to 1-benzylpyridinium or 3-methyl-1-benzylpyridinium was observed.228 Under identical conditions, hydride transfer to substituted 1-benzylpryidinium derivatives such as 3-carbamoyl-1-benzylpyridinium and 3-methoxycarbonyl-1-benzlpyridinium was observed (Table 7).228 These reactivity observations suggest that the hydricity

4.5. Nicotinamides and Pyridinium Derivatives

Nicotinamide derivatives are another family of small molecules that has received considerable attention in hydride transfer reactivity. Biological hydride transfer is carried out by nicotinamide coenzymes, such as 1,4-NADH (1,4-dihydronicotinamide adenine dinucleotide). The hydricity of 1,4-NADH itself has been determined in the aqueous environment in which it is typically utilized, although there is some discrepancy in the specific value obtained.80,103,223 This may stem from inaccuracy in the reduction potential of N,N,N′,N′-tetramethyl-p-phenylenediamine (TMPA), which is used as a reference.224 A potentialpKa thermochemical cycle can be constructed with literature values to obtain a hydricity ΔG°H− = 28.9 kcal/mol in water.224 This value is consistent with a pH-dependent redox potentiometry study in which the open-circuit potential of transferring 1H+/2e− was measured over a broad pH range.223 The insolubility of NAD+ in acetonitrile has precluded direct measurements of hydricity in this solvent. Several phosphate-free models, such as 1-methylnicotinamide (MNA) and 1-benzylnicotinamide (BNA), have been characterized in acetonitrile and other solvents, affording a wide range of hydricity values from 56 to 82 kcal/mol.77,88,149,156 It is important to emphasize that many of the experimental values for nicotinamide derivatives are hydricity enthalpies (ΔH°H−), so any comparisons to ΔG°H− values requires a correction factor estimating the entropic contribution.156,225 DuBois and co-workers77 and Creutz and coworkers37 have measured hydricity free energies (ΔG°H−) for several 1,4-dihydronicotinamide derivatives. Comparison of these ΔG°H− values to the corresponding hydricity enthalpies (ΔH°H−) from Zhu et al.150,156,225 affords a linear correlation (Figure 4). This correlation may not be due solely to entropic changes across the series, as Creutz suggested the presence of a systematic error in Zhu’s hydricity enthalpies.37 Regardless of the precise contributing factors, this trend provides a semiempirical method that can be used to obtain estimates for ΔG°H− values for a wide range of nicotinamide-derived species having known enthalpy estimates. For example, two common nicotinamide derivatives, PNAH and HEH, can be estimated to have ΔG°H− = 61 kcal/mol

Table 7. Reactivity of Nicotinamide Derivatives as a Function of Substituents

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of [Rh(Cp*)(bpy)H]+ lies in the range of 48−58 kcal/mol, bracketed between the hydricities of the electron-rich and electrondeficient 1-benzylpryidinium derivatives (Table 7). The analogous [Ir(Cp*)(bpy)H]+ complex is isolable and has ΔG°H− = 62 kcal/mol in acetonitrile.128 Ranges of ΔG°H− with variations of the bpy ligand were recently reported to be between 22.9 and 33.4 kcal/mol in water.56,128 The attenuated hydricities of iridium complexes relative to rhodium analogues are reflected in the few reports of hydride transfers from [Ir(Cp*)(αdiimine)H]+ species to nicotinamide derivatives.231 Dihydronicotinamide derivatives more commonly act as the hydride donor in reactions with [Ir(Cp*)(α-diimine)(L)]2+ species. Clean formation of iridium hydride complexes was observed upon hydride transfer from NADH to [Ir(Cp*)(phen)(H2O)]2+ and [Ir(Cp*)(phenylpyridine)(H2O)]+ in neutral water.243,244 The hydride transfer reactivity of [Ir(Cp*)(bpy)H]+ can be tuned using visible light. In the ground state in acetonitrile, [Ir(Cp*)(bpy)H]+ (ΔG°H− = 62 kcal/mol) is thermodynamically incapable of hydride transfer to MNA (ΔG°H− = 56 kcal/mol).128 However, upon irradiation with 460 nm light, hydride transfer is observed. The expected product, 1-methyl-1,4-hydronicotinamide, is observed by 1H NMR spectroscopy along with a rarely observed doubly reduced product, 1-methyl-1,4,5,6-tetrahydronicotinamide.128 Nicotinamides have been used as reversible hydride carriers, delivering hydride equivalents derived from H2 or facilitating the release of H2. Maenaka and co-workers used [Ir(Cp*)(LCN)(H2O)]+, (LCN = N-4-pyrazolbenzoic acid, Scheme 30) to facilitate the reaction of eq 78, with the direction of the equilibrium tunable based on pH.245 Under acidic conditions, NADH is converted to NAD+ with concomitant H2 evolution,

in accord with the equilibrium expression of eq 78. Under neutral pH 7 conditions and H2 pressure, NAD+ is reduced to 1,4-NADH. Successful reversibility indicates that all of the hydricity values are closely spaced in this system. In water, 1,4NADH and H2 have ΔG°H− values of 28.9 and 34.2 kcal/mol, respectively, indicating that under standard conditions formation of H2 should be favored by 5.3 kcal/mol. The free energy difference will decrease as the pH rises. The hydricity of the proposed Ir hydride intermediates must also be similar to that of 1,4-NADH in order to mediate hydride transfer in both directions, although the protonation state of the benzoate portion of the ligand will affect the hydricity. 1,4−NADH + H+ ⇌ NAD+ + H 2 PH2 [NAD+] Keq = [H+] [NADH]

(78)

Metal hydrides can couple reductants, such as H2 or electrochemically generated H+/e−equivalents, with a biocompatible reducing agent through the catalytic regeneration of 1,4-NADH for use by various reductase enzymes.149,246,247 A representative example involves the use of H2 as a reductant to produce catalytic amounts of 1,4-NADH, which mediates the reduction of thermoanaerobium brockii alcohol dehydrogenase (TdADH) for the reduction of ketones to chiral alcohols in water.248 The catalyst [Rh(Cp*)(bpy)Cl]+ reacts with H2 to form an intermediate that undergoes hydride transfer to NAD+, forming 1,4-NADH, which in turn transfers a hydride equivalent to TdADH for the catalytic transformation of 2-heptanone to (S)-2-heptanol.248 Knowledge of thermodynamic hydricity allows for rational designs of such organometallic-biochemical systems to perform specialized reactions under benign conditions. The use of hydricity values to predict reactivity can be extended to a large number of additional systems and not just the handful of case studies mentioned here. Knowledge of thermochemistry has helped guide the development of reactions in many systems to date and remains an active area for expansion. The use of hydricity values for transformations involving ligands is discussed in the next section, and in section 6, an outlook for future work in the area of hydricity will be presented.

Scheme 30. Dependence of Reactivity upon pH for Ir(Cp*)(LCN)H

5. APPLICATIONS OF THERMODYNAMIC HYDRICITY INVOLVING METAL−LIGAND COOPERATION 5.1. Introduction

In section 4 we considered outer-sphere hydride transfer reactions involving a metal hydride and a small molecule substrate. In this section, three modes of metal−ligand cooperation in hydride transfer are presented. In section 5.2, examples of hydride transfer from metal hydrides to organic ligands bound to a transition metal are described. Substrate activation by metal coordination can dramatically change the hydride acceptor ability, leading to implications in numerous catalytic cycles involving metal hydride transfer reactions. Section 5.3 includes several examples where hydride transfer reactions have been mediated through the secondary coordination sphere of transition metal complexes. A “hydride shuttle” mechanism moves H− from a metal hydride to a pendant group on the ligand, and then to an exogenous hydride acceptor. This strategy can confer a kinetic advantage in hydride transfer. In section 5.4, we consider examples where net hydride transfer is achieved by formally “separating” the hydride ion into proton and electron 8676

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equivalents. Such “separated” hydride transfers often result in transfer of a proton to a ligand site and electrons to the metal center. Leading examples of hydride transfer involving ligands or metal−ligand cooperation are highlighted in this section, with an emphasis on systems that have been considered in thermochemical detail. Not every example in this section involves a well-defined metal hydride complex, but all of the examples are governed by the same thermodynamic constructs. Furthermore, many of the highlighted cases have some ambiguity regarding the structure of a transition metal complex that is reduced by two electrons and protonated. A net hydride transfer to the metal complex has occurred, so it is helpful to consider hydricity. However, the product is not always a metal hydride.

The more electrophilic carbonyl ligands will be more susceptible to hydride donation to generate a metal formyl. For example, the Re complex [Re(Cp)(CO)2(NO)]+ (with formyl conjugate hydride donor ΔG°H− = 55.0 kcal/mol) reacts with relatively weak hydride donors, including [Ni(dmpe)2H]+ (ΔG°H− = 50.8 kcal/mol).255 Substitution of one CO for the strong sigma electron donor PMe3 in [Re(Cp)(CO)(NO)(PMe3)]+ (with formyl conjugate hydride donor ΔG°H− = 44.1 kcal/mol) leads to more than a 10 kcal/mol change in hydride acceptor ability. The latter complex requires substantially stronger hydride donors for CO reduction: [Ni(dmpe)2H]+ does not react with this carbonyl complex, but [Pt(dmpe)2H]+ (ΔG°H− = 42 kcal/mol) readily generates the formyl product.254 Building on this work, Wiedner and Appel performed a comprehensive study of the various intermediates involved in subsequent CO reduction steps in the [Re(Cp)(PPh3)(CO)(NO)]+ system in MeCN,148 as illustrated in Scheme 31. Protonation of a

5.2. Hydride Transfer to Substrates Activated by Metal Coordination

Carbon monoxide reduction illustrates the advantage of using metal−ligand binding to promote hydride transfer. Carbon monoxide itself is a gas that is poorly soluble in organic solvents and relatively unreactive: even very strong borohydride reductants do not react with CO gas. On the other hand, metal carbonyl complexes are quintessential hydride acceptors in organometallic chemistry, and borohydrides can readily transfer hydride to the bound carbon monoxide to produce a metal formyl (M−CHO) complex.249 Hydride transfer routes to metal formyl complexes have long been explored in the search for selective molecular catalysts for reductive coupling of CO.250−253 A broad study of metal formyl hydricity values was performed by DuBois and co-workers.52,254 As shown in Table 8, metal

Scheme 31. Thermochemical Scheme for [CpRe(PPh3)(CO)(NO)]+ in MeCN

Table 8. Hydricity of Metal Formyls in Acetonitrile, Determined by Hydride Transfera hydride donor

hydricity of (kcal/mol)

ref

Re(Cp)(NO)(CO)(CHO) Re(Cp*)(NO)(CO)(CHO) Re(CpDMEG)(NO)(CO)(CHO)c Ru(Cp*)(CO)2(CHO) cis-Mn(PPh3)(CO)4(CHO) cis-[Ru(bipy)2(CO)(CHO)]+ Re(Cp)(NO)(PPh3)(CHO) Re(CO)4(Ph2P-NMe2)(CHO)c Re(CO)4(Ph2P−NC(NMe2)2)(CHO)c Re(Cp)(NO)(PMe3)(CHO) Re(Cp)(NO)(PEt3)(CHO) Re(Cp*)(NO)(PMe3)(CHO)

55.0 52.6 51.7 50.6 50.2 49.6 46.5 ∼45b ∼45b 44.1 43.8 42.1

254 254 256 254 254 254 254 257 257 254 254 52

formyl would give a hydroxycarbene complex, [ReCH(OH)]+ (pKa = 10.6). Hydride transfer to the hydroxycarbene would generate a hydroxymethyl, Re-CH2OH (ΔG°H− = 58 kcal/mol). Product release could be accomplished at this stage by protonolysis of the Re−C bond to release methanol, or by CO insertion into the Re−C bond to produce C2 products. CO reduction by sequential hydride and proton transfers has long been considered a promising route to homogeneous CO reduction catalysis, but until these studies little thermochemical knowledge of the late-stage intermediates was available.250 The resulting free energy landscape indicates that radical intermediates are thermodynamically unrealistic and that initial net hydrogenation of CO to form the [ReCH(OH)]+ species is the thermodynamically most demanding step. With metal formyl hydricity values in hand, one can readily predict whether a particular metal hydride will be thermodynamically capable of performing the first step in a CO reduction sequence. Miller, Labinger, and Bercaw employed such thinking to guide the design of cooperative CO reduction systems

a

The metal formyl hydricites are anchored to [Pt(dmpe) 2H]+ having ΔG°H− = 42.5 kcal/mol. Adjustment of the hydricity for [Pt(dmpe)2H]+ will shift the hydricities of the formyls by an equivalent amount. bHydricity bracketed between [Pt(dmpe)2H]+ and [Ni(dmpe)2H]+. cSee Chart 5 for structural representation.

formyl complexes feature hydricity values in the same range as many metal hydride hydricity values. Some metal formyl complexes are quite potent thermodynamic hydride donors. The hydricity of metal formyl complexes tracks broadly with the electrophilicity of the metal center, as reported by the CO stretching frequency of the corresponding metal carbonyl complexes. Metal complexes with poor σ-donor ligands or multiple π-backbonding ligands (such as CO and NO+) support less back-donation into the CO π-system, leading to higher CO stretching frequencies and more electrophilic CO ligands. 8677

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involving late transition metal complexes to activate H2 and transfer a hydride to a midtransition metal carbonyl complex. In cases where metal hydrides were not hydridic enough to reduce the metal carbonyl of interest, Lewis acids were employed to facilitate the kinetics and thermodynamics of hydride transfer.220,258,259 In the presence of an appropriate base, H2 itself is thermodynamically capable of donating a hydride to a CO ligand.254 To circumvent the need for an exogenous metal hydride donor, Teets, Labinger, and Bercaw designed Re carbonyl complexes with pendant bases in order to facilitate direct H2 heterolysis at the CO ligand (Chart 5). The hydricity of

Scheme 32. Change in Hydricity of BNAH upon Coordination to a Ru Complex

Chart 5. Re Carbonyl Complexes Bearing Pendant Bases

[Ru(terpy)(bpy)(BNAH)]2+ to yield [Ru(terpy)(bpy)(BNAH)]+ with coordination through the amide.260 Cyclic voltammetry of [Ru(terpy)(bpy)(BNAH)]+ in acetonitrile reveals a shift in the potential for oxidation of the coordinated and deprotonated BNAH ligand of over −0.3 V relative to the oxidation potential of free BNAH.260 A potential−hydricity relationship can be used to relate the shift in potential to a shift in hydricity. Assuming the BDFE of BNAH•+ does not change upon coordination to Ru, then a 300 mV potential shift would make BNAH more hydridic by ∼7 kcal/mol.

the Re formyl must be matched to the acidity of the pendant acid, just as the [Ni(diphosphine)2]2+ catalysts with pendant bases balanced hydricity and acidity in H2 heterolysis. In the case of the monophosphine pentacarbonyl rhenium cations, the formyl species had ΔG°H− ≈ 45 kcal/mol and pendant acids with pKa ≈ 15, which leads to H2 heterolysis being unfavorable by 8−11 kcal/mol.257 The free energy for H2 addition was improved to ∼3 kcal/mol by moving to a [Re(CpDMEG)(NO)(CO)2]+, a rhenium carbonyl platform with an improved hydride acceptor ability (ΔG°H− = 51.7 kcal/mol for the formyl complex).256 Despite the thermodynamic accessibility of the H2 addition product shown at the bottom of Chart 5, no evidence for H2 heterolysis was observed at PH2 = 3 atm over prolonged time periods, indicating a substantial kinetic barrier for the process. Nicotinamide derivatives also exhibit shifts in thermodynamic hydricity upon coordination to a transition metal center. The role of 1,4-NADH and NAD+ coordination to [Ru(terpy)(bpy)]2+ was studied by Ishitani and co-workers.236−238,240,260 The influence of coordination on thermodynamic hydricity is evident in the equilibrium between BNAH and B(Et)NA+ (Scheme 32). In acetonitrile, B(Et)NAH is expected to be more hydridic than BNAH,225 and indeed B(Et)NAH was not detected upon mixing BNAH and B(Et)NA+.225,236 A Ru complex of BNAH was synthesized and allowed to react with B(Et)NA+, as shown in Scheme 32. When BNAH is coordinated to the Ru center, complete hydride transfer to B(Et)NA+ is observed. Metalcoordinated BNAH must be much more hydridic than free BNAH, such that the extent of the equilibrium with B(Et)NA+ can be completely altered, a trend attributed to proton loss upon coordination.260 A base such as NEt3 is required for the Ru-promoted hydride transfer. The base deprotonates an amide proton from

5.3. Hydride Transfer Involving Ligands That Act as Hydride Shuttles

The secondary coordination sphere can be used to mediate hydride transfer with pendant “hydride shuttle” groups (Scheme 33). Scheme 33. Incorporation of a Hydride Shuttle

Hydride shuttles can improve the kinetics of hydride transfer for intramolecular reactions or facilitate metal hydride formation. If direct hydride transfer from a donor to an acceptor is slow, introducing a hydride shuttle can facilitate a two-step transfer process in which rapid hydride transfer from the donor to the pendant shuttle is followed by another rapid hydride transfer from the shuttle to the acceptor. This situation can be advantageous if hydride transfer between two bulky species is slow, or if there is there is some other large kinetic barrier to direct hydride transfer. Boranes have found wide utility on the periphery of a range of transition metal complexes, and this topic is comprehensively reviewed by Maity and Teets in this special issue.261 Borane hydride shuttles facilitate CO reduction in Re carbonyl complexes. When paired with a strong phosphazene base, frustrated Lewis pair (FLP) chemistry forms a fleeting borohydride intermediate (ΔG°H− near 26 kcal/mol)131,155 that readily 8678

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Scheme 34. Reactivity of a Re Complex Containing a Pendant Borane

Chart 6. Borane-Appended Metal Complexes Capable of H2 Activation

Hantzsch ester group.276,277 The proposed mechanism involved metal−ligand cooperation, with the Rh center binding the imine to render its carbon more electrophilic and the Hantzsch ester acting as the hydride transfer reagent (Scheme 35).

reduces the nearby carbonyl ligands on Re (Scheme 34).259 The borane removes hydride from H2 and transfers it to the Re−CO moiety, mediating the hydride transfer and opening a new mechanistic pathway for H2 heterolysis and reduction of CO. Pendant boranes can also shuttle hydrides from a transition metal center to the Re carbonyl.220,258 Hydride transfer from a Pt hydride to the pendant borane is moderately unfavorable, but once the borohydride is formed, it has a strong thermodynamic driving force for carbonyl reduction to produce a formyl. Once the borohydride is formed, the pendant group plays an additional role of promoting rapid intramolecular CO reduction, with product selectivity dictated by the strength of the Lewis acid, the ring size of the hydride transfer intermediates, and the formation of strong B−O bonds after hydride transfer.220 Several other borane-containing ligands have showed interesting hydride transfer chemistry in cooperation with transition metals. Bourissou and co-workers pioneered the development of phosphinoborane complexes of late metals, building new bonding motifs and activating small molecules.262−266 Peters and co-workers reported hydrogen cleavage in Ni phosphinoborane complexes such as Ni(MesDPBPh) (Chart 6), wherein a M−BR3 interaction is broken and a B−H bond is formed.267−269 The borohydride was proposed as a key intermediate in olefin hydrogenation using this hydride shuttle strategy. Cooperative chemistry involving hydride shuttling through a pendant borane in small molecule activation and catalysis has since expanded to include examples on Fe and Co supported by phosphinoborane ligands from Peters and co-workers,270−273 as well as on Pt supported by the “FcPPB” phosphinoborane ligand by Cowie and Emslie274 and “BIM” iminoborane ligands from Figueroa and co-workers (Chart 6).275 Several ligands containing nicotinamide, Hantzsch ester, or acridinium groups have also been prepared and carefully studied as hydride shuttles. An excellent review on the transition metal coordination chemistry of these organic hydride shuttles by McSkimming and Colbran was recently published.149 As discussed in section 4, nicotinamides are ubiquitous hydride (and electron) carriers in biology. In one striking example of a pendant hydride shuttle, Colbran and co-workers reported transfer hydrogenation of imines catalyzed by a Cp*Rh complex supported by an iminopyridine ligand containing a pendant

Scheme 35. Rh-Catalyzed Imine Hydrogenation with a Pendant Pyridinium Hydride Shuttle

Other complexes containing organic hydride donor/acceptor sites have been designed to incorporate transition metal sites capable of excited state electron transfer to promote formation of the organic hydride donor.149,278 For example, Fujita and Tanaka led a team of researchers that prepared a Ru complex bearing a pendant acridine group that could be reduced by 2H+/2e− to give the dihydroacridine derivative.279 The Ru polypyridyl core of the molecule can act as a light harvester, with visible light absorption triggering excited state electron transfer to generate a pendant acridine radical that can undergo proton-coupled electron transfer to afford a pendant hydride equivalent.280,281 Subtle structural changes to the organic ligand can have dramatic effects in this system. Addition of a methyl in the polypyridyl backbone shuts down productive photoreduction chemistry,282 and 8679

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one proton to the oxo (and loss of another proton to solvent) to produce a Ru(II)−OH complex and the organic carbonyl R2CO.296,297,299−301 Although the authors did not report hydricity values, published Pourbaix (potential−pH) diagrams already contain the pKa, potential, and BDFE values needed to determine aqueous ΔG°H− values (Scheme 36).78 As an example,

changing the position of the nitrogen so that the C−H bonds point away from the metal center leads to 1e− reduction and C−C bond formation instead of hydride formation via 1H+/2e− reduction.283 The hydricity of Ru/dihydroacridine complexes has been estimated by experimentally calibrated DFT calculations, with Ru(II) species calculated to have ΔG°H− = 89 kcal/mol.88 This calculated hydricity is quite weak but in accord with experimental observations that the pendant hydride can be transferred to Ph3C+ (conjugate donor ΔG°H− = 99 kcal/mol) and that the Ru complex can catalyze acetone hydrogenation.283,284 The calculations predict that addition of an electron to this Ru(II) hydride donor would produce a reduced species that is almost 40 kcal/mol more hydridic. If accessible, these species would be in the range required for reduction of CO2 or metal carbonyl complexes.

Scheme 36. Thermochemical Scheme for Ru Complex of a Bisimidazole Pyridine Ligand in H2O

5.4. Separated Hydride Transfer Involving Proton-Coupled Electron Transfer

The examples of metal−ligand cooperation described above have all featured hydride ion transfer from one species to another. Net hydride transfer can be accomplished by concerted hydride ion (H−) transfer reactions or by separate transfer of H+ and 2e−. Appropriate metal−ligand combinations can act cooperatively to mediate individual proton and electron transfer steps, often with the metal acting as a redox center for electron transfer and the ligand acting as a proton transfer site. Although these reactions do not always involve metal hydride bonds, metal hydrides were often the expected product. Few examples of “separated” hydride transfers have been proposed, although biological alcohol oxidation by galactose oxidase is thought to proceed by single electron transfer to a Cu center and H atom transfer to a nearby cysteine.285 The use of metal−ligand cooperation has recently emerged as a major theme in catalyst design, with examples of ligands that are redox-active, ligands containing proton relays, and various other modes of “non-innocent” behavior in which the ligand and metal both play key roles in reactivity.146,286−290 In some instances, such as Milstein’s pyridyl-based pincer ligands, a metal hydride can be transferred as a proton to ligand backbone.291 This section provides illustrative examples of metal−ligand cooperation in hydride transfer where the H+ and e− equivalents are distributed across the metal center and the ligand. Metal chalcogenide complexes can participate in net hydride transfer chemistry through metal−ligand cooperation. Although high degrees of covalency can complicate bonding descriptions, the hydride transfer can be formally considered as involving electron transfer to/from the metal center and proton transfer to/from the chalcogenide ligand. Oxo and sulfido metal complexes are frequently encountered in both molecular and materials electrocatalysis, motivating thermochemical understanding of these structural motifs. Metal hydrides are not typically formed in this case, because protonation/reduction pathways are thermodynamically more favorable. Nonetheless, M−SH and M−OH species can be assigned hydricity values using the thermochemical methods described in this Review. Metal oxo complexes are key intermediates in many chemical transformations, including electrocatalytic oxidation of water and alcohols.292−295 Alcohol oxidation has been proposed to proceed via hydride transfer from the alcohol H−CR2OH bond to the metal oxo fragment in Ru polypyridyl complexes.296−299 Depending on the catalyst system, Ru(IV), Ru(V), and Ru(VI) oxo species can trigger the initial alcohol oxidation process.297−299 In the case of Ru(IV)O species, hydride abstraction from an alcohol involves 2e− reduction of Ru and transfer of

we include a thermochemical cycle that contains new values for the hydricity of Ru(II)−OH species of a recently reported catalyst supported by bisimidazole−pyridine ligands.297,299 The high-valent oxo species are exceptional hydride acceptors, based on our estimated hydricity of both the Ru(II) hydroxo (ΔG°H− = 73 kcal/mol in H2O) and Ru(III) hydroxo (ΔG°H− = 102 kcal/mol in H2O) complexes. The metal oxo species in Scheme 36 are much better hydride acceptors than the transition metal complexes discussed in section 3.3, consistent with the observed reactivity involving hydride abstraction from an alcohol C−H bond. High-valent Mn oxo complexes can activate C−H bonds via numerous pathways, including hydride abstraction.302,303 The oxo-bridged dimer [(bpy)2Mn(μ-O)2Mn(bpy)2]4+ reacts with toluene to generate diarylmethane products that implicate a benzylic carbocation intermediate formed by hydride abstraction. Mayer and co-workers constructed a thermochemical cycle in acetonitrile based on the reduction potentials of [(bpy)2Mn(μO)2Mn(bpy)2]4+ and [(bpy)2Mn(μ-O)2Mn(bpy)2]3+ along with the pKa of [(bpy)2Mn(μ-O)(μ-OH)Mn(bpy)2]3+ (Scheme 37). Scheme 37. Themochemical Scheme for Mn(bpy)2n+ in MeCN

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The bridging hydroxide complex [(bpy)2Mn(μ-O)(μ-OH)Mn(bpy)2]3+ has ΔG°H− = 122 kcal/mol and is one of the weakest hydride donors reported. Due to its strong hydride acceptor ability, [(bpy)2Mn(μ-O)2Mn(bpy)2]4+ is capable of removing hydride from triphenylmethane (ΔG°H− = 99 kcal/mol) or even the relatively unactivated C−H bonds of toluene. The Mn product is a Mn(III)/Mn(III) hydroxy-bridged dimer, indicating that a separated hydride transfer has occurred by two 1e− reductions of Mn and a one H+ transfer to the bridging oxo. Molybdenum sulfide complexes and (nano)materials serve as models for both molybdenum sulfide containing enzymes304 and industrial hydrodesulfurization catalysts305 and are among the most promising H2 evolution electrocatalysts.15,306−309 The dinuclear sulfide-bridged species CpMo(μ-S)2(μ-SCH2S)MoCp was one of the earliest molecular Mo sulfide electrocatalysts for H2 evolution.309 Extensive thermochemical studies were undertaken on Cp*Mo(μ-S)4MoCp* to better understand the possible mechanisms for H2 formation.147,310 As is evident from the thermochemical cycle of Scheme 38, the dimolybdenum

Figure 5. Free energy landscape for [Cp*Mo(μ-S4Hm)MoCp*]n+ at 1 atm H2 and pH 5.

H+/H2 couple (−0.37 V vs FeCp2+/0) under the same conditions (1 atm H2 and a pH of 5).310 With a free energy landscape in hand, mechanistic pathways can be vetted. Pathways to H2 evolution can be imagined by tracing the addition of 2e− and 2H+ to an oxidized species (as in the left side of the scheme). In Figure 5 it is readily apparent that pathways that proceed via the anionic species (1− overall charge) will pay a high thermodynamic price under acidic conditions, and therefore these species and their analogues are unlikely to be energetically accessible intermediates at low overpotentials. Many other ligand classes, including some of those most commonly employed as supporting ligands, can support separated hydride transfer reactions that avoid the formation of a metal hydride. A series of studies led by Nocera and Hammes-Schiffer revealed the rich proton- and electron-storing capabilities of porphyrin ligands that are ubiquitous in both synthetic and biological catalytic species. Bediako and co-workers reported a Ni porphyrin complex fitted with a “hangman” carboxylic acid site that can facilitate proton transfer and proposed that H2 evolution electrocatalysis proceeds via a stepwise proton transfer/electron transfer (PTET) pathway to generate a hydridic intermediate.311 A detailed computational study found that the only pathway that correlated well with all of the experimental results did not involve a metal hydride intermediate but rather protonation at a meso position of the porphyrin to generate a phlorin intermediate that exhibited reactivity consistent with ΔG°H− = 42 kcal/mol.312,313 In an analogous Co porphyrin system, competing pathways involving either a traditional metal hydride intermediate or an organohydride phlorin intermediate are proposed, depending on the choice of acid strength.314 Hydrogen evolution catalyzed by transition metal complexes is nearly always proposed to proceed via metal hydride intermediates (see section 4), but the Co and Ni porphyrin complexes exhibit metal−ligand cooperation that reveals new H2 release pathways that involve net hydride transfer without a metal hydride. Even the venerable pentamethylcyclopentadienyl (Cp*) ligand has been shown to mediate hydride transfer reactions in recent studies. As discussed in section 4.5, the hydride [Rh(Cp*)(bpy)H]+ had long been proposed as the key intermediate in Rh-catalyzed transfer hydrogenation reactions and 1,4-NADH regeneration schemes.228,229,315,316 The iridium analogue, [Ir(Cp*)(bpy)H]+, is well-characterized structurally and thermochemically,128,317 but similar synthetic tactics do not lead to the expected rhodium hydride. In almost simultaneous publications, two groups revealed that proton transfer to the pentamethylcyclopentadienyl ligand occurs to yield a diene complex [Rh(Cp*H)(bpy)]+.224,318 The Cp*H complex can release H2 by reaction with acid sources and transfers hydride to

Scheme 38. Thermochemical Schemes for Cp*Mo(μ-S)4MoCp* in MeCN

system supports an extraordinary range of stable oxidation and protonation states. The stability of ten species was quantified through proton/electron transfers to Cp*Mo(μ-S)4MoCp*, enabling the determination of five pKa values, seven homolytic BDFE values, and four hydricity values. Each hydride transfer reaction involves breaking an S−H bond, but the electrons are transferred from different sites for different species, based on whether a disulfide bond is formed between two of the bridging sulfide ligands. Surprisingly, no metal hydride intermediates were detected. The formation and cleavage of the disulfide bond results in maintaining formal oxidation states of MoIV for one of the two metal centers, and MoIII, MoIV, or MoV for the other metal center in each of the species in Scheme 38. Therefore, for each hydride transfer, the source of electrons may be from sulfur or molybdenum. With so many possible pathways and so many thermochemical values represented in Scheme 38, alternative methods of visualizing the data can be helpful. The authors developed “free energy landscape” plots that map the relative free energy of each of the 10 species available in Scheme 38. The corresponding free energy landscape is shown in Figure 5, for which the reference state is the 8681

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NAD+, implicating this diene complex as an intermediate capable of separated hydride transfer (H+ from Cp*H and 2e− from the Rh(I) center) in a range of catalytic processes (Scheme 39).

thermochemical understanding of [Ni(PR2NR′2)2]2+ catalysts to develop a new system for formate oxidation, a reaction relevant to direct formic acid fuel cells. In an initial report, nine Ni catalysts were examined, with ΔG°H− varying from 56 to 68 kcal/mol. The relative hydricities indicate that formate is thermodynamically capable of hydride transfer to the dications [Ni(PR2NR′2)2]2+ to form the Ni hydride and release CO2. Consistent with the thermochemical predictions, treatment of Ni(II) precursors with formate produced the corresponding Ni(II) hydrides.112 Detailed structure−function studies showed that the electrocatalytic activity did not correlate directly with hydricity. Instead, an inverse correlation with hydricity was observed (Figure 6).

Scheme 39. Proposed Mechanism of Rh-Catalyzed NAD+ Reduction Involving Cp*H Intermediate; N∪N is 2,2′Bipyridine

Figure 6. Inverse correlation of hydricity of Ni(PR2NR′2)22+ with catalytic rate for oxidation of formate. For each data point, the substituents are shown at phosphorus (R, shown in red) and at nitrogen (R′, shown in blue).

Spectroscopic evidence for [Rh(Cp*)(bpy)H]+ itself was obtained at low temperature, and it is difficult to completely rule out the possibility of hydride transfer from a fleeting Rh−H species, but certainly both (thermodynamically equivalent) pathways are possible. The hydricity of this species in water, ΔG°H− = 23 ± 2 kcal/mol, reveals that the ligand-based hydride donor is substantially more hydridic than the metal-based iridium hydride (ΔG°H− = 31.5 ± 1 kcal/mol in water)and sufficiently hydridic to thermodynamically favor the experimentally observed hydride transfer to NAD+. Some metal−ligand combinations are poised to support either direct hydride transfer involving the metal center or separated hydride transfer involving metal−ligand cooperation. The large family of [Ni(PR2NR′2)2]2+ complexes discussed throughout this Review nicely illustrate this dichotomy, as shown in eq 79. The Ni(II) hydride can readily transform into a nitrogen-protonated Ni(0) complex through proton transfer.87,319,320 This tautomerization reaction enables the possibility of intermolecular hydride transfer occurring through separated transfer of a proton and two electrons.

The authors ultimately proposed a separated hydride transfer mechanism that completely avoids a traditional hydride transfer, wherein formate binding is followed by proton transfer to a pendant nitrogen and two-electron reduction of the Ni center (Scheme 40).321 An alternative pathway involving direct hydride transfer from formate to nickel was recently suggested on the basis of DFT calculations.322 Scheme 40. Proposed Mechanism of Electrocatalytic Oxidation of Formate

It is intriguing to imagine new net hydride transfer reaction pathways from species that contain separate proton and electron transfer sites. A pathway involving two-electron reduction of a Ni center occurring concertedly with proton transfer to a pendant base has been proposed for a [Ni(PR2NR′2)2]2+ system engaged in formate oxidation.112 The collaborative studies by Appel, Kubiak, and co-workers used detailed

6. SUMMARY AND OUTLOOK This Review provides an introduction and overview of the determination and application of the thermodynamic hydricity 8682

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of transition metal hydrides. Three main experimental approaches to determining hydricity values were presented, along with several less commonly used methods. Hydricity can be determined by establishing equilibrium with another species of known hydricity, by establishing equilibrium in the heterolytic cleavage of H2, or by combining equilibria for the transfer of a proton and two electrons. Experimental values are now available for many transition metal hydrides, and more than 100 transition metal hydricity values have been reported. Sections 4 and 5 explored how hydricity values can be used to understand and predict reactivity patterns of transition metal hydrides. The examples illustrate that researchers armed with a knowledge of relative hydricity can correctly predict hydride transfer reactions under a variety of conditions. This predictive capability leads to the incorporation of such reactions in catalytic transformations in which hydride transfer plays a key role. In this final section, we consider several emerging and exciting new areas where we believe that further research in thermodynamic hydricity of transition metal hydrides is warranted. As discussed in section 4, separate scales and thermochemical conventions have made comparisons between transition metal hydrides and organic or main group hydrides difficult. More small molecule thermodynamic hydricity values are needed, specifically free energies that are directly comparable with hydricity values for transition metal hydrides. Accurate hydricity free energies are not available for many of the most commonly employed substrates in catalysis, including olefins and organic carbonyls. Each new hydricity measurement would support the development of predictive reaction chemistry involving metal hydrides. The vast majority of the hydricity values in Tables 2−4 have been determined for monometallic complexes with a terminal hydride. The tetra-iron complexes developed by Berben and co-workers are striking as isolated examples of bridging hydrides.121−123,186,198 Bridging hydrides have appeared in the context of hydricity in several instances, but usually these species are formed as intermediates during hydride transfer,37,52 or as side products formed when the product of hydride transfer reacts with the starting hydride (especially with boranes).36 In sharp contrast to the prevailing notion that a bridging hydride would be less reactive, the iron clusters are among the strongest hydride donors collected here. Further studies comparing terminal and bridging hydrides are needed, and comparisons between dinuclear systems and small clusters could help define design parameters for future catalysts that mediate hydride transfer reactions. The influence of solvent on hydride transfer is a rapidly emerging area in thermodynamic hydricity. It is clear from Table 4 that far fewer hydricity values have been determined in aqueous solution than in acetonitrile. The ΔG°H− values are much lower in water, but the relationship between hydricity values in different solvents is not yet well-understood. Early reports suggested that changes in the solvation of hydride ion (H−) in MeCN and water play a major role.37,118 Developing an understanding of how various solvents influence hydricity values is a critical area for expansion. A better understanding of solvent effects could help guide the development of catalytic transformations performed in variety of different solvents (such as moving from MeCN solvent to more sustainable solvents such as water). One can also imagine changing solvents for the explicit purpose of tuning a hydride transfer stepfor example, it has been shown that CO2 reduction to formate is thermodynamically favorable in water for some metal hydrides while being unfavorable in MeCN.37,145

Most of the examples in sections 4 and 5 were single-step hydride transfer processes, but many exciting chemical transformations involve multistep processes. For example, the multielectron reduction of CO2 to CH3OH or other fuels involves multiple catalytic intermediates that each need to be reduced by hydride equivalents. Currently, the hydride acceptor ability of these intermediates are generally not well-understood, although the CO reduction chemistry discussed in section 5.2 provides a rare example of this type of analysis. Keeping with CO2 as an example, a recent report by Klankermayer, Leitner, and coworkers implicated hydride transfer to coordinated formic acid in the hydrogenation of CO2 to CH3OH.323 Knowledge of the hydride acceptor ability of metal-bound formic acid would be relevant to catalyst design. Understanding the thermochemical constraints for hydride transfer to these intermediates is critical, emphasizing the importance of elucidating the factors governing how hydride acceptor ability of an organic molecule changes upon binding a metal. Bond dissociation free energies (BDFEs) have been used extensively in correlations and predictions of organometallic reactivity, and Marcus theory successfully correlates driving force and kinetics, providing a powerful predictive tool for electron transfer (ET),324 proton-coupled electron transfer (PCET) reactions,325 and even some hydrogen atom transfer reactions.326,327 Analogous rate−driving force relationships for hydride transfer are underdeveloped. Creutz and co-workers reported an example of strong correlation between the hydride acceptor ability of pyridinium derivatives and the rate of hydride transfer from [Ru(terpy)(bpy)H]+ to the organic acceptor.37 Similar correlations of hydricity and rate have been observed with organic hydride donor/acceptor pairs.328−330 It will be interesting to see if structurally diverse hydrides, such as d8 species and d6 species (which bind solvent after hydride transfer), have similar driving force dependence. This Review has focused on experimental hydricity determination, but computational methods for hydricity determination are also likely to play an important role in future development. Prediction of thermochemical parameters using theoretical and computational techniques is increasing and has been extensively explored in some systems.84−87,89,92,331 Computational studies could address some of the challenges described above, such as the origin of solvent effects, shifts in hydride acceptor ability upon binding to metal centers, or the transfer of bridging hydrides. Rapid, general, and accurate computation-based hydricity predictions could become invaluable tools in catalyst design. Less than 20 years ago, the first thermodynamic hydricity values of metal hydrides were reported. Since then, over 100 hydricity values have been reported, providing insight into bonding and reactivity across a diverse range of metal complexes. Thermodynamic hydricity has emerged as a valuable predictive tool for reaction chemistry and a principle of catalyst design, conjuring a bright future for the development of new methods, expansion in new directions, and connections with chemical applications.

ASSOCIATED CONTENT S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.chemrev.6b00168. Spreadsheet containing hydricity values in MeCN and the corresponding thermochemical parameters from which the values were calculated (XLSX) 8683

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AUTHOR INFORMATION

Aaron M. Appel studied as an undergraduate at Washington State University, worked in the laboratory of Scot Wherland, and graduated in 2001. He studied at the University of Colorado, Boulder, under the supervision of Mary Rakowski DuBois and Dan DuBois. He received his Ph.D. in 2005 and thereafter started as a postdoctoral fellow at Pacific Northwest National Laboratory, where he worked with James Franz and was hired as a research scientist in 2008. His research interests include catalysis for the interconversion of energy and fuels, with an emphasis on thermodynamics.

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Notes

The authors declare no competing financial interest. Biographies Eric S. Wiedner graduated from the Missouri University of Science and Technology in 2004 with a B.S. in Chemistry, then received his Ph.D. in 2009 from the University of Michigan under the guidance of Marc Johnson. He subsequently studied as a postdoctoral fellow with Dan DuBois and Morris Bullock at Pacific Northwest National Laboratory and was hired as a research scientist in 2013. His research interests include the study of organometallic catalysis, understanding correlations between catalyst structure and thermodynamic properties, and using thermochemical considerations to predict and control chemical reactivity.

ACKNOWLEDGMENTS We dedicate this Review to the memory of Carol Creutz (Brookhaven National Laboratory), whose landmark insights had a profound influence on our understanding of hydricity in water. We thank Daniel DuBois for his extensive contributions to studies of thermodynamics of hydricity and for many helpful discussions. E.S.W. and A.M.A. were supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences & Biosciences. R.M.B. was supported by the Center for Molecular Electrocatalysis, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences. Pacific Northwest National Laboratory is operated by Battelle for the U.S. Department of Energy. M.B.C. and A.J.M.M. were supported by the Division of Chemical Sciences, Geosciences & Biosciences, Office of Basic Energy Sciences of the U.S. Department of Energy through Grant DE-SC0014255. C.L.P. was supported by the National Science Foundation Center for Enabling New Technologies through Catalysis (CHE-1205189) and is a Fellow of the Royster Society.

Matthew B. Chambers received his B.A. in Chemistry from Cornell University in 2007 while gaining research experience in the laboratory of Pete Wolczanski. He obtained his Ph.D. from the Massachusetts Institute of Technology under the supervision of Dan Nocera in 2013. Matt then moved to Paris to work for Marc Fontecave at the Collège de France as a postdoctoral research assistant for two years. In 2015, he returned to the United States and joined the group of Alex Miller at the University of North Carolina at Chapel Hill and is continuing his interest in mechanisms and organometallic catalyst design for energy storage transformations. Catherine L. Pitman received her B.S. in Chemistry from the University of Notre Dame in 2011, spending her junior year studying chemistry in New College at the University of Oxford. In 2012, she joined the lab of Alex Miller at the University of North Carolina at Chapel Hill. Her graduate work has focused on understanding the reactivity of transition metal hydrides in water in both the ground and excited states.

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Morris Bullock received a B.S. from the University of North Carolina at Chapel Hill (1979) and a Ph.D. from the University of Wisconsin Madison, where he worked in Chuck Casey’s group. Following a postdoctoral appointment with Jack Norton at Colorado State University, he moved to Brookhaven National Laboratory (Long Island, New York) in 1985. He moved to Pacific Northwest National Laboratory in 2006, where he is a Laboratory Fellow and the Director of the Center for Molecular Electrocatalysis (efrc.pnnl.gov). He has been studying reactions of metal hydrides for over 30 years, including hydride transfer reactions, hydrogen atom transfers, proton transfers, and development of molecular electrocatalysts. Alexander J. M. Miller was introduced to inorganic chemistry by Gregory Hillhouse as an undergraduate at the University of Chicago. After receiving his B.S. in 2005, Alex obtained a Ph.D. in Chemistry in the laboratory of John Bercaw and Jay Labinger at the California Institute of Technology in 2011, then continued his studies with Karen Goldberg and James Mayer as a Dreyfus Environmental Chemistry Postdoctoral Fellow at the University of Washington, Seattle. In July 2012, Alex joined the faculty at the University of North Carolina at Chapel Hill as an Assistant Professor of Chemistry. In pursuit of efficient and sustainable methods for the synthesis of fuels and chemicals from renewable resources, the Miller group is using thermochemical and kinetic mechanistic understanding to guide the development of catalysts with activity and selectivity tunable by changes to the surrounding environment. 8684

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