Metal hydricity

Phosphorus—Hydrogen Bond. A hydrogen bound to phosphoms has Httie acidic or hydric character. Most of the reactions the bond undergoes are those of a reducing agent. P—H bonds are formed by hydrolysis of active metal phosphides or phosphoms haUdes, by the rearrangement of P—O—H or P—S—H linkages, or by the hydrolysis of P—P bonds (6,17). 

Although orbital hybridizations and molecular shapes for hypovalent metal hydrides of the early transition metals and the normal-valent later transition metals are similar, the M—H bonds of the early metals are distinctly more polar. For example, metal-atom natural charges for YH3 (+1.70), HfH4 (+1.75), and TaHs (+1.23) are all significantly more positive than those (ranging from +0.352 to —0.178) for the homoleptic hydrides from groups 6-10. Indeed, the empirical chemistry of early transition-metal hydrides commonly reveals greater hydricity than does that of the later transition-metal hydrides. 

DuBois et al. carried out extensive studies on the thermodynamic hydricity of a series of metal hydrides [13, 15-19]. The determination of thermodynamic hydricity generally requires several measurements (coupled with known thermochemical data) to constitute a complete thermochemical cycle. As with other thermodynamic cycles, obtaining reliable values in an appropriate solvent can be a difficult challenge, and this is sometimes coupled with problems in obtaining reversible electrochemical data. Scheme 7.2 illustrates an example in which the hydricity of cationic monohydrides have been determined. 

Eq. (8) requires determination of the two-electron oxidation potential of L M by electrochemical methods. When combined with the two-electron reduction of protons in Eq. (9), the sum provides Eq. (10), the AGh- values of which can be compared for a series of metal hydrides. Another way to determine the AGh-entails the thermochemical cycle is shown in Scheme 7.3. This method requires measurement of the K of Eq. (11) for a metal complex capable of heterolytic cleavage of H2, using a base (B), where the pK., of BH+ must be known in the solvent in which the other measurements are conducted. In several cases, Du-Bois et al. were able to demonstrate that the two methods gave the same results. The thermodynamic hydricity data (AGh- in CH3CN) for a series of metal hydrides are listed in Table 7.4. Transition metal hydrides exhibit a remarkably large range of thermodynamic hydricity, spanning some 30 kcal mol-1. 

M—H bond dissociation energies, 1, 287 photochemistry, 1, 251 single crystal neutron diffraction, 1, 578 stability toward disproportionation, 1, 301 Metal—hydrogen bonds bond dissociation energy in 1,2-dichloroethane, 1, 289 stable metal hydrides in acetonitrile, 1, 287 thermochemical cycle, 1, 286 in THF and dichloromethane, 1, 289 olefin insertion thermodynamics, 1, 629 in Zr(IV) bis-Cp complexes, 4, 878 Metal—hydrogen hydricity data, 1, 292... 

The M-H bonds of transition-metal hydride complexes may be cleaved heterolyti-cally (H, H transfer) or homolytically (H transfer). AG for the transfer in Equation 1.1 is readily quantified by pKj measurements (see Chapter 3). Analogous measurements for H transfer, or hydricities , are difficult because the loss of generates a vacant coordination site. However, AG for Equation 1.2 can be determined indirectly, from electrochemical and pJG measurements in the appropriate solvent [1, 2], and we can thus compare the hydricities of various hydride complexes (see Chapter 3). The lowest values of AG. (corresponding to the complexes most eager to transfer H ) are found for second- and third-row transition metals [3], which is why those (relatively expensive) metals are good donors and effective catalysts for reactions like ionic hydrogenation [5-10],... 

DuBois and coworkers carried out extensive studies of the thermodynamic hydric-ity of metal hydrides (3.2) in many cases the metal product resulting following hydride transfer will bind a CH3CN ligand (not shown in 3.2), so comparisons of hydricity wiU be solvent-dependent. 

Thermochemical studies of acidity and hydricity are extremely valuable, since they can help determine the energies of potential intermediates in catalytic cycles, and can thus guide the choice of complexes proposed as catalysts. But, since catalysis is a kinetic phenomenon, the kinetics of delivery of a proton or hydride are also important. The kinetics of proton transfer from metal hydrides to amines [21] or metal alkynyl complexes [22], as well as degenerate proton transfers between metal hydrides and metal anions [21] led to the conclusion that proton transfers from metal hydrides have a high intrinsic barrier. 

DuBois and co-workers introduced the thermochemical cycle in Scheme 4 as a means to determine the hydride donor power (AG°h-) or hydricity, of a cationic metal hydride. The method requires the knowledge of metal-hydride acidity (p- a) data and the electrode potentials for the oxidation of the metal-hydride conjugate base to two-electron oxidized counterpart, either by two successive one-electron processes (Equation (10), Scheme 4), or by one two-electron process (Equation (11), Scheme 4). The thermochemical cycle is derived from one that was introduced by Parker and co-workers for use in organic systems/ The accuracy of values on an absolute scale rests upon the... 

Scheme 4 Thermochemical cycle for determination of metal-hydride hydricities from pKa and electrode potential data. The quantity C in Equations (10) and (11), when electrode potentials and pKa data are obtained in the same solvent and E° data are referred to the Cp2Fe/Cp2Fe scale, is 333 kJ mor (solv = MeCN) and 387 kJ mor (solv = DMSO).

Scheme 5 Thermochemical cycle for the determination of metal-hydride hydricities by heterolytic cleavage of H2 in the presence of a metal hydride and a base. The term 318 kJ mor represents the hydricity of H2 in MeCN.

It has also been demonstrated that metal-hydride hydricities can be determined from equilibrium measurements of base-induced, metal-promoted heterolytic cleavage of H2. This method requires only the measurement of the equilibrium constant for heterolysis of H2 in the presence of a base B with known basicity (Equation (12), Scheme 5). This technique does not require electrode potential measurements, but the thermochemical cycle that is shown in Scheme 5 reveals an obvious relationship to the cycles which have already been discussed the constant term of 318kJmol in Equation (13), Scheme 5 which yields the hydricity is derived from the H /H and redox... 

Table 5 Metal-hydride acidities, conjugate base oxidation potentials, and hydricities in acetonitrile...

Table 6 Hydricities of metal hydrides determined by base-induced metal-promoted heterolytic splitting of H2 in acetonitrile...

Table 10 C-H BDEs and hydricities of metal formyl complexes...

The combination of acidity and hydricity data for metal hydrides can be used to evaluate and better understand factors that govern stabilities of metal hydrides with respect to disproportionation and H2 loss. It is obvious that a metal hydride will be unstable with respect to deprotonation in the presence of a base B if the of is greater than the piTa of the metal hydride. Similarly, Equation (13), Scheme 5 can be used to determine whether a metal hydride with a known hydricity will react with an acid BH of known to generate H2. Thus, the p a of the acid that causes Equation (12), Scheme 5 to be in equilibrium with 1 will be given by pA BH+ in Equation (29). 

DuBois has studied many examples of these reactions and from these data and related experiments has generated die "hydricity scale" shown in Table 12.2. Particularly electron-rich complexes and third-row metal complexes are stronger hydride donors than less-electron-rich complexes and first-row analogs. 

The tendency of a transition metal hydride to transfer H to a substrate is called hydricity [ 12]. It is possible to determine the Gibbs free energy of the splitting of the covalent polar M-H bond to afford a metal cation and the hydride ion in solution. The hydricity is not parallel to the polarity of the M-H IxMid, nor can it be predicted on the basis of the electronic structure of the metal atom. It is a complex property that can be modeled for transition metal hydrides using multiparameter approaches. The hydricity concept applies to the interaction of M-H bonds with CO2 as well [13]. The reactivity of M-H bonds toward CO2 is linked to reactions that may have industrial interest, such as the hydrogenation of CO2 to afford formic acid (4.2) and the electrochemical reduction of CO2 to other Cl or C1+ molecules (4.3). 

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