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Logo of nihpaAbout Author manuscriptsSubmit a manuscriptHHS Public Access; Author Manuscript; Accepted for publication in peer reviewed journal;
 
J Am Chem Soc. Author manuscript; available in PMC 2010 April 8.
Published in final edited form as:
PMCID: PMC2735118
NIHMSID: NIHMS103475

Slow Hydrogen Transfer Reactions of Oxo— and Hydroxo— Vanadium Compounds: the Importance of Intrinsic Barriers

Abstract

Reactions are described that interconvert vanadium(IV) oxo-hydroxo complexes [VIVO(OH)(R2bpy)2]BF4 (1a-c) and vanadium(V) dioxo complexes [VVO2(R2bpy)2]BF4 (2a-c) [R2bpy = 4,4′-di-t-butyl-2,2′-bipyridine (tBu2bpy), a; 4,4′-dimethyl-2,2′-bipyridine (Me2bpy), b; 2,2′-bipyridine (bpy), c]. These are rare examples of pairs of isolated, sterically unencumbered, first-row metal-oxo/hydroxo complexes that differ by a hydrogen atom (H+ + e). The VIVtBu2bpy derivative 1a has a useful 1H NMR spectrum, despite being paramagnetic. Complex 2a abstracts H• from organic substrates with weak O–H and C–H bonds, converting 2,6-tBu2-4-MeO-C6H2OH (ArOH) and 2,2,6,6-tetramethyl-N-hydroxy-piperidine (TEMPOH) to their corresponding radicals ArO• and TEMPO, hydroquinone to benzoquinone, and dihydroanthracene to anthracene. The equilibrium constant for 2a + ArOH [left harpoon over right harpoon] 1a + ArO• is (4 ± 2) × 10−3, implying that the VO–H bond dissociation free energy (BDFE) is 70.6 ± 1.2 kcal mol−1. Consistent with this value, 1a is oxidized by 2,4,6-tBu3C6H2O•. All of these reactions are surprisingly slow, typically occurring over hours at ambient temperatures. The net hydrogen-atom pseudo-self-exchange 1a + 2b [left harpoon over right harpoon] 2a + 1b, using the tBu- and Me-bpy substituents as labels, also occurs slowly, with kse = 1.3 × 10−2 M−1 s−1 at 298 K, ΔH = 15 ± 2 kcal mol−1, and ΔS= 16 ± 5 cal mol−1 K. Using this kse and the BDFE, the vanadium reactions are shown to follow the Marcus cross relation moderately well, with calculated rate constants within 102 of the observed values. The vanadium self-exchange reaction is ca. 106 slower than that for the related RuIVO(py)(bpy)22+ / RuIIIOH(py)(bpy)22+ self-exchange. The origin of this dramatic difference has been probed with DFT calculations on the self-exchange reactions of 1c + 2c and on mono-cationic ruthenium complexes with pyrrolate or fluoride in place of the py ligands. The calculations reproduce the difference in barrier heights and show that transfer of a hydrogen atom involves more structural reorganization for vanadium than the Ru analogs. The vanadium complexes have larger changes in the metal–oxo and metal–hydroxo bond lengths, which is traced to the difference in d-orbital occupancy in the two systems. This study thus highlights the importance of intrinsic barriers in the transfer of a hydrogen atom, in addition to the thermochemical (bond strength) factors that have been previously emphasized.

Introduction

The interconversion of metal-oxo and reduced metal-hydroxo complexes (eq 1) is of fundamental importance in chemistry, biology, and the environment. This is also perhaps the paradigmatic example of processes that involve transfer of both a proton and an electron.15 In

Mn+=O+H (e+H+)M(n1)+OH
(1)

that context, reactions (1) have been called proton-coupled electron transfer (PCET), concerted proton electron transfer (CPET) and/or hydrogen atom transfer (HAT);17 distinctions among these various terms are discussed below. Reaction (1) is particularly important in metal-mediated oxidation reactions, where metal-oxo complexes are able to abstract a hydrogen atom (H• [equivalent] e + H+) from a substrate, forming the corresponding hydroxo complex. This is thought to be the key step in hydrocarbon C–H bond oxidations by metalloenzymes such as cytochromes P450,8 TauD,9 soluble methane monooxygenase,10 and class I ribonucleotide reductases.11 This step is also involved in the oxidation of butane to maleic anhydride by an oxovanadium(V) catalyst (to cite one industrial example)12 and in smaller-scale reactions of permanganate, oxochromium(VI) compounds, activated manganese dioxide, and other reagents.13 Oxo-hydroxo interconversions also occur widely in the aqueous speciation of metals from vanadium to osmium.14

Given the importance of reactions that interconvert oxo and hydroxo species, there are few studies of systems where both the metal-oxo and metal-hydroxo compounds have been isolated and their hydrogen atom transfer reactivity studied. Meyer and others have developed an extensive reaction chemistry of polypyridyl ruthenium-oxo and -hydroxo complexes, where the ruthenium hydroxo species can be observed but are typically not isolable.15 Bakac et al. have studied the hydrogen transfer reactions of transient chromium–oxo and –superoxo complexes.16 Systems with isolable oxo and hydroxo species have been developed by Borovik (Fe, Mn) and by Theopold (Cr), using bulky supporting ligands and hydrogen bonding to prevent formation of dimers (which may also inhibit reactions with substrates).17 In most other systems, the metal-oxo compounds are strong oxidants and/or the hydroxo compounds are unstable, for instance condensing to form µ-oxo species.

The recent synthesis of mononuclear vanadium(IV)-oxo-hydroxo-bipyridine compounds by Kabanos et al.,18 coupled with the known vanadium(V)-dioxo-bipyridine compounds,19 provide a rare example of an isolable 1st row metal-oxo/hydroxo system. Oxovanadium compounds, including polyoxometallates, have been extensively studied for a variety of reasons, including their insulin mimetic activity,20 their presence in haloperoxidase enzymes,21 and their utility for the catalytic oxidation of organic substrates.12,2226 Vanadium oxides are also valuable heterogeneous catalysts, as in butane oxidation mentioned above.12,

Reported here are net hydrogen transfer reactions involving these relatively simple vanadium oxo and hydroxo complexes. This is part of recent emphasis, in our lab2732 and others,13,1517,3337 on PCET/HAT/CPET reactions of metal coordination compounds, including iron, ruthenium, osmium, cobalt, nickel, manganese, copper, and chromium species. When the species that differ by a hydrogen atom (one e and one H+) can be isolated and examined in detail, new understanding of such H· transfer reactions can be derived. We have emphasized the importance of thermochemistry in these reactions (the strength of the H–X bonds that are broken and formed) and we have found that the rate constants for many of these reactions roughly follow the Marcus cross relation.2,29 Extending that work to this vanadium system, we have examined both cross reactions (eq 2, N–N = a bipyridine ligand) and self-exchange reactions, Ln V=O + HO–VLn → Ln V–OH + O=VLn. The vanadium reactions are remarkably slow compared to other metal-mediated hydrogen transfer reactions. Most notably, the self-exchange occurs roughly a million times slower than the self-exchange between structurally related ruthenium oxo/hydroxo-bipyridine complexes. DFT calculations on models for the vanadium and ruthenium self-exchange reactions reproduce this dramatic disparity and indicate that it results from differences in inner-sphere reorganization energies.

equation image
(2)

Results

1. Vanadium-oxo and -hydroxo complexes

have been prepared with 4,4′-di-t-butyl-2,2′-bipyridine (tBu2bpy, a series), 4,4′-dimethyl-2,2′-bipyridine (Me2bpy, b), and 2,2′-bipyridine (bpy, c). The vanadium(IV) oxo-hydroxo complexes [VIVO(OH)(R2bpy)2]BF4 (1a–c) reported by Kabanos et al.18 have been made, using a revised procedure, from V(O)SO4, HBF4, BaCO3, and the R2bpy ligand. We also describe (see Experimental Section) an improved route38 to the known19 dioxovanadium(V) compounds [VVO2(R2bpy)2]BF4 (2a–c), by adding two equivalents of ligand to aqueous HBF4 solutions of NaVO3 at pH = 1. The X-ray structure of 2c (Figure 1; Table 1 below)39 shows a cation with a distorted octahedral geometry. The two short V=O bonds of 1.627(5) Å are typical of VVO2+ compounds and are similar to the two previously reported crystal structures of this cation (in compounds with different counterions prepared in less direct fashions).19,40,41 The V–N bonds trans to an oxo are 0.15(3) Å longer than the V–N bonds trans to a bpy nitrogen, illustrating the strong trans influence of the oxo ligands.

Figure 1
ORTEP of the cation in [VVO2(bpy)2]BF4 (2c), with hydrogen atoms omitted for clarity.
Table 1
Calculated and solid-state bond lengths (Å) and angles (°) for [VIVO(OH)(bpy)2]+ (1c), [VVO2(bpy)2]+ (2c), and the transition structure for self-exchange.

Compounds 1 and 2 have been characterized by mass spectrometry, IR and 1H NMR spectra, in addition to comparison of their optical spectra with those of reported derivatives. UV-Vis spectra of 2a–c show strong peaks with λmax ≤ ~350 nm (ε ~ 20 000 M−1 cm−1) that tail into the visible. Compounds 1a–c also show two d-d bands at ca. 520 and 750 nm (ε ~ 30 M−1 cm−1), as expected for octahedral oxovanadium(IV) complexes.42 All the IR spectra show strong vanadium-oxo stretching modes, with the one stretch for the vanadyl compounds 1a–c (vV[equivalent]oca. 970 cm−1) at higher energy than the symmetric and antisymmetric stretches for 2a–cVO2ca. 930, 910 cm02212;1).40 Cyclic voltammograms of 1a in MeCN show a wave that is chemically reversible at high scan rates, with E1/2 = 1.065 V vs. Cp2Fe+/0(Figure S1 and Figure S2), and a number of features below −1 V.39 CVs of 2a show no reduction wave between 2 and −1 V vs. Cp2Fe+/0. Below −1 V, CVs of 2a have reduction waves identical to those of 1a even with efforts to remove all proton sources. Complex 2a reacts only very slowly with Cp2Fe in MeCN, indicating that its E1/2 is < 0 V vs. Cp2Fe+/0. The aqueous electrochemistry of the bipy complexes 1c and 2c, the only derivatives with appreciable solubility in water, showed an irreversible, pH-independent oxidation at 0.68 V vs. Cp2Fe+/0 (as well features below −1 V similar to those observed in MeCN).

The 1H NMR spectra of the dioxo complexes 2a–c at 295 K all have three or four very broad resonances in the aryl region, indicating that all the R-py rings are equivalent on the NMR timescale. Upon cooling, the peaks sharpen and separate into two sets of pyridine resonances, as shown in Figure 2 for 2a. Assignments were made using 2D COSY and NOESY experiments at 190 K in CD2Cl2. Raising the temperature of CD3CN solutions to 340 K causes the tBu resonances of 2a to coalesce and then sharpen, although the aryl peaks remain broad. The spectra indicate that the [VO2(R2bpy)2]+ cations have the expected C2 structure in solution (as found in the solid state, Figure 1) and have a fluxional process which equilibrates the pairs of R-py rings. Line shape analysis of spectra of 2a in CD2Cl2 from 190 to 300 K gives the Eyring parameters ΔH = 18 ± 2 kcal mol−1 and ΔS = 14 ± 2 cal mol−1 K−1 for the fluxional process. Complex 2b behaves similarly (ΔH = 16 ± 2 kcal mol−1, ΔS = 9 ± 2 cal mol−1 K−1), while 2c is insoluble in CD2Cl2. The spectra are similar in CD3CN for all three compounds, with broad peaks of similar linewidths at ambient temperature that sharpen but do not reach the low temperature limit at 235 K. The positive entropies and large trans influence, as evidenced by the structure of 2c, suggest a dissociative mechanism.

Figure 2
Variable temperature 1H NMR spectra of 2a in CD2Cl2, R = tBu.

[VIVO(OH)(tBu2bpy)2]BF4 (1a) has a useful and characteristic 1H NMR spectrum despite being a d1 paramagnet. Four peaks are observed between 1 and 4 ppm (Figure 3), assigned to the four inequivalent tBu groups. No other resonances are discernable between +100 and −100 ppm. The observation of 4 tBu resonances indicates that exchange of the proton between the hydroxo and oxo moieties is slow on the NMR timescale. While proton transfer between oxygens is usually very rapid, slow exchange has previously been described for other oxo-hydroxo compounds.43 The one broad tBu resonance, labeled 4 in Figure 3, is tentatively assigned to the tBu on the pyridine ring that is coplanar with the vanadyl group, because only this pyridine has overlap of its π system with the vanadium dxy orbital that holds the unpaired electron.18 The 1H NMR spectrum of 1b shows only two very broad peaks between 2 and 5 ppm (likely due to the Me groups) and compound 1c does not show a 1H NMR spectrum.

Figure 3
1H NMR of 1a in CD2Cl2 with the tBu peaks noted 1–4.

Solutions of the dioxo compounds 2a–c in organic solvents decompose over the course of days at ambient temperatures. The decomposition appears to be faster at higher concentrations and elevated temperatures, and is evidenced by a change in the color of solution from pale yellow to green. Decomposition appears to give primarily [VIVO(OH)(R2bpy)]+ (1a–c), by NMR. While the other products have not been identified, oxo-bridged vanadium dimers are often green.44,45 Compounds 1a–c are stable in solution for a week at room temperature.

2. Hydrogen atom self-exchange

1H NMR spectra of CD3CN and CD2Cl2 solutions containing both 1a and 2a are essentially the sum of the spectra of the two species. There is no evidence for any chemical exchange process, even at 340 K when the spectrum of 1a is fairly sharp (Figure 2). This indicates that degenerate hydrogen atom self-exchange between the cis-oxo-hydroxo-vanadium(IV) and cis-dioxovanadium(V) compounds 1a and 2a is slow on the NMR timescale, too slow to measure by NMR line broadening. This result is surprising because we have used NMR line broadening to measure self-exchange rate constants in a number of related systems.46

We have therefore examined the reaction between 1a and 2b to give 1b and 2a, and the reverse (Scheme 1). This can be considered a pseudo self-exchange reaction because the alkyl substituents are electronically similar (Me vs. tBu) and are distant from the reacting centers. The reaction was monitored by 1H NMR because three of the four species have characteristic resonances (1a, 2a and 2b). Because of the broad spectra of 2a and 2b near 300 K (Figure 2), these experiments must be done either at elevated or lower temperatures, to sharpen the tBu peaks of 2a, which would otherwise obscure the tBu peaks of 1a. Figure 4 shows a portion of the NMR spectra for a reaction at 320 K, in which the three narrower tBu resonances of 1a convert to the single tBu peak of 2a at δ1.45. Over the course of these reactions, a small amount of free tBu2bpy ligand (1–5%) is observed as a sharp peak at δ1.40, its diamagnetic signal higher than the broader peaks despite its low concentration. The integrals of the NMR resonances were determined using the line-fitting tool of MestreC.47 Plots of concentrations vs. time, such as Figure 5a, illustrate the good mass balance in these reactions: in all of the reactions, the sum of the concentration of tBu2bpy species ([1a] + [2a]) remains constant within error.

Figure 4
tBu region of 1H NMR spectra of the of pseudo self-exchange reaction between 1a and 2b at 320 K in CD3CN; [filled triangle] denotes free tBu2bpy.
Figure 5
(a) Plot of the concentrations vs. time for the pseudo-self-exchange reaction of 1a and 2b in CD3CN at 320 K. (b) Eyring plot of pseudo self-exchange rate constants.
Scheme 1
Hydrogen atom pseudo self-exchange reaction

The kinetic data fit well to a second-order approach to equilibrium model, which yields both equilibrium and rate constants Kse and kse.48 The same Kse and kse were found from reactions done in either direction and with various starting conditions, including beginning with three species present initially. Over the entire 280–320 K temperature range, the equilibrium constants are 1.0 ± 0.6 indicating that these are very close to true self-exchange reactions. kse at 320 K in anhydrous CD3CN is (1.0 ± 0.3) × 10−1 M−1 s−1. Eyring analysis of rate constants at 280 and 310–330 K gives ΔH = 15 ± 2 kcal mol–1 and ΔS = −16 ± 5 cal mol−1 K−1.49 These values give a rate constant at 298 K (which is not directly measurable, see above) of kse = (1.3 ± 0.1) × 10−2 M−1 s−1. This is a remarkably slow reaction, more than a million times slower than the sterically similar self-exchange reaction of RuIVO(py)(bpy)22+ + RuIII(OH)(py)(bpy)22+;28 this comparison will be developed below. At 320 K, the reaction proceeds very similarly in CD2Cl2 giving kse(CD2Cl2) = (1.7 ± 0.2) × 10−1 M−1 s−1, 1.7 times faster than in CD3CN. Addition of H2O up to a concentration of 0.5 M (100 equivalents) slowed the rate in CD3CN by a factor of 1.6. The same rate within error was observed when D2O was added, indicating a kinetic isotope effect of close to unity. There is no indication, from either the NMR or the kinetic data, for initial formation of a hydrogen-bonded adduct between the hydroxo complexes 1 and the oxo complexes 2. A stable complex involving paramagnetic 1 would likely broaden the resonances for 2, but no broadening is observed in mixtures of 1 and 2. The negative activation entropy of ΔS = −16 ± 5 cal mol−1 K−1 also suggests that the reaction proceeds from separated 2b + 1a.

3. Hydrogen atom transfer cross reactions

Reactions of 1a and 2a with hydrogen atom donors and acceptors have been explored, taking advantage of their good solubility and distinctive optical and 1H NMR spectra. Complex 2a oxidizes hydroquinone in CD3CN to give benzoquinone and 1a, in ~80% yield by 1H NMR (eq 3; N–N = tBu2bpy). With stoichiometric amounts of reactants at 10 mM concentrations, the reaction requires several hours to reach completion at ambient temperatures. UV-vis spectra of reaction 3 show formation of 1a, benzoquinone, and an unidentified species with a strong absorbance extending from 380 nm into the visible region.39 This unknown species is not formed from 2a plus the product benzoquinone, nor from 1a plus hydroquinone, as neither of these reactions show any discernable change over 24 h by 1H NMR and UV-Vis. Monitoring the kinetics by 1H NMR, the rate of disappearance of hydroquinone under pseudo-first order conditions of excess 2a suggests a rate constant of (4.0 ± 0.5) × 10−2 M−1 s−1 for reaction 3, assuming a bimolecular rate-limiting step.39

equation image
(3)

Complex 2a also slowly oxidizes 2,6-di-tert-butyl-4-methoxyphenol (ArOH), with 1a appearing over a period of hours at NMR concentrations in CD3CN at 298 K (eq 4). Optical spectra of the reaction show the distinct spectrum of the phenoxyl radical, which is quite stable (Figure 6).50 The formation of 1a and ArO• indicates that net H-atom transfer has occurred. Much less than a stoichiometric amount of radical is formed, however. At [2a] = 1.1 mM, 3.0 mM ArOH gives a 3.6% yield of ArO• (in 4000 seconds), while 45 mM ArOH gives 25% ArO•. Less phenoxyl is formed when 1a is present initially: 1.9 mM 2a plus 5.4 mM ArOH gives 3% ArO•, but only 0.6% ArO• when 1.9 mM 1a is added. These and fifteen other measurements, spanning an order of magnitude in [2a] and [ArOH], indicate that equilibrium is reached with K4 = (4 ± 2) × 10−3 in CH3CN at 298 K, ΔG4 = 3.3 ± 0.5 kcal mol−1. ArOH has a weak O–H bond, with a bond dissociation free energy in MeCN (BDFE) of 73.9 ± 1 kcal mol−1,51 from this the VO-H BDFE of 1a can be calculated as 70.6 ± 1.2 kcal mol−1. The kinetics, monitored by UV-vis spectroscopy and analyzed using SPECFIT,52 fit well to a second-order approach to equilibrium model with k4 = (1.4 ± 0.5) × 10−3 M−1 s−1.

Figure 6
(a) Optical spectra and (b) absorbance vs. time plots for the reaction of 2a (1.6 mM) with tBu2(MeO)C6H2OH in CH3CN at 298 K.

equation image
(4)

In contrast, 2a does not react with 2,4,6-tri-tert-butylphenol (tBu3ArOH) in MeCN. Neither the UV-vis spectrum of 1a nor the distinct sharp UV spectrum of the stable radical tBu3PhO• are observed in millimolar solutions of these reagents. The reverse reaction, tBu3PhO• plus 1a, does proceed slowly. For instance, a solution of 3 mM 1a and 30 mM tBu3PhO• in CD3CN showed a 20% yield of 2a by NMR after several days, although many side products were also observed. This is the only case where H-atom abstraction from 1a has been observed (reactions of tBu2(MeO)C6H2O• are problematic because the radical solutions contain excess phenol). The different position of the equilibria for 2a + tBu3ArOH vs. 2a + tBu2(MeO)C6H2O• is consistent with the 3 kcal mol−1 higher O–H BDFE of tBu3ArOH in MeCN, 77 ± 1 kcal mol−1.53,54

The hydroxylamine TEMPOH (2,2,6,6-tetramethyl-N-hydroxy-piperidine), which has an O-H BDFE of 66.5 ± 0.5 kcal mol−1,29,55 also slowly reduces 2a to 1a. This reaction is however complicated by the further reaction of 1a with TEMPOH. 1H NMR spectra of 7 mM 2a plus 73 mM TEMPOH at 275 K showed complete disappearance of 2a within 3 h but the maximum yield of 1a was only 30% and other products are observed, including free tBu2bpy.39 With a larger excess of TEMPOH, 2a is consumed faster but the yield of 1a is no higher. These issues have prevented quantitative studies but the data are qualitatively consistent with a bimolecular rate constant of ca. 10−1 M−1 s−1. TEMPOH and hydroquinone seem capable of coordinating and displacing tBu2bpy, while this seems to be prevented by the 2,6-di-t-butyl groups of the phenols.

Complex 2a is also slowly reduced by xanthene and dihydroanthracene (DHA). A sealed NMR tube containing 13 mM 2a and 51 mM DHA showed stoichiometric 20% conversion to anthracene and a corresponding amount of 1a after 2.5 days at 298 K by 1H NMR, but further reaction is complicated by the decomposition of 2a. Xanthene reduces 2a more quickly, with 45% conversion to 1a in 30 hours under similar conditions. Within 7 days, 2a is quantitatively reduced to 1a, and 5% free ligand is observed. No xanthone or anthrone are observed in these reactions.56 The kinetics of 2a + xanthene were followed by the stoichiometric growth of the d-d band of 1a at 754 nm (Figure S1039). The pseudo-first order kobs varied linearly with [xanthene], yielding kxan = (7.6 ± 2.4) × 10−5 M−1 s−1. The more limited data for the DHA reaction are consistent with kDHA ~ 2 × 10−6 M−1 s−1. The xanthene C-H BDFE is 70.9 kcal mol1,28,57 so H• transfer from xanthene to 2a is essentially thermoneutral; the reaction is presumably driven by further reactions of the xanthyl radical.

4. DFT Calculations

Hydrogen atom self-exchange between 1c and 2c has been examined using DFT calculations, which were performed with the B3LYP functional58 and the 6-31G* basis set.59 The polarizable continuum model (PCM)60 was used to approximate the effects of solvation in acetonitrile. All of the calculations were carried out with the Gaussian03 package of programs.61

Gas phase structures of the 1c and 2c cations were calculated, and their bond lengths and angles are listed with the experimental solid-state structures of the BF4 salts in Table 1. For 1c, the calculated V=O is close to that expected for a VIV=O distance.62 The experimental V=O and V–OH distances are unusually long and short, respectively (as was noted in reference 18a), perhaps reflecting some disorder in the crystal.

Bringing complexes 1c and 2c together forms a hydrogen-bonded ‘precursor complex,’ in which the hydroxo group of 1c acts as a hydrogen bond donor to one of the oxo groups of 2c. Formation of this precursor complex is calculated to be endothermic by 17.9 kcal mol−1 in the gas phase, due to the large Coulombic repulsion between the positively charged vanadium reactants. Solvation reduces this Coulombic repulsion, as formation of the hydrogen-bonded precursor complex is calculated to be exothermic by −4.7 kcal mol−1 in acetonitrile. This negative ΔH° in acetonitrile is counteracted by the positive entropy of formation, which is calculated (in the gas-phase) as ΔS° = −36.8 cal mol−1 K−1, so ΔG° is calculated to be positive, 6.8 kcal mol−1. Thus, the calculations find that the ground state of the system is the separated reactants, not the hydrogen-bonded complex, consistent with experimental observations.

Transition state structures for hydrogen exchange between 1c and 2c were also calculated. In acetonitrile, ΔH for formation of the C2 symmetric transition structure, in which the proton is halfway between the oxygen atoms, is calculated to be 12.4 kcal mol−1 above the separated reactants and 17.1 kcal mol−1 above the hydrogen-bonded complex. This value of ΔH = 12.4 kcal mol−1 for reaction of 1c and 2c is in reasonable agreement with the experimental value of ΔH = 15 ± 2 kcal mol−1 in MeCN.

Interestingly, the calculated gas-phase value of ΔH = 16.3 for degenerate hydrogen atom exchange from the hydrogen-bonded precursor complex is nearly the same as the PCM value. Thus, although the PCM and gas-phase values for formation of the hydrogen-bonded complex differ by 22.6 kcal mol−1, the PCM and gas-phase values of for ΔH for hydrogen exchange differ by only 0.8 kcal mol−1. The presence of a polar solvent is calculated to stabilize the precursor complex and transition structure almost equally, so the solvent has very little effect on the barrier for hydrogen exchange within the complex.

The calculated C2 transition structure63 has a nearly linear O–H–O bridge (see Table 1 and Figure 11 below). The two vanadium centers are twisted about this bridge, with a calculated O=V---V=O torsion angle of 124.2°. The V–O bonds involved in the H transfer are calculated to be 1.710 Å long at this transition state, intermediate between the calculated lengths of 1.859 and 1.595 Å for the V–OH and V=O bond lengths of isolated 1c and 2c, respectively.

Figure 11
Computed structure (left) and SOMO (right) of the 1c + 2c transition structure. The V-O-O-V dihedral angle is 118.4°.

As mentioned in the previous section, the reaction of 1a with 2b is much slower than the apparently similar self-exchange reaction of RuIVO(py)(bpy)22+ with RuIII(OH)(py)(bpy)22+.28 Unfortunately, computational comparison of the reaction between 1c and 2c and this ruthenium self-exchange reaction proved problematic, due to the dicationic nature of the ruthenium complexes. The dipositive charges on each of the separated Ru+2 reactants resulted in very strong electrostatic repulsion between them, so that neither a hydrogen-bonded complex nor a true transition state for the ruthenium self-exchange reaction could be located. Therefore, model systems were chosen,64 in which the neutral pyridine ligand was replaced with a fluoride or pyrrolate ligand.65 The resulting ruthenium reactants are, like their vanadium counterparts, monocationic; so hydrogen-bonded complexes and transition structures for hydrogen exchange could be located for both the ruthenium fluoride and ruthenium pyrrolate complexes.

Formation of these hydrogen-bonded precursor complexes in the gas phase is calculated to be endothermic by 15.5 kcal mol−1 for RuIV(O)F(bpy)2+ + RuIII(OH)F(bpy)2+ [henceforth abbreviated as Ru(O)F + Ru(OH)F], and by 12.6 kcal mol−1 for RuIVO(pyrr)(bpy)2+ plus RuIII(OH)(pyrr)(bpy)2+ [henceforth abbreviated as RuO(pyrr) + Ru(OH)(pyrr), pyrr = pyrrolate]. However, PCM calculations predict that in acetonitrile solution, formation of these hydrogen-bonded complexes should be exothermic with ΔH° = −5.2 and −8.7 kcal mol−1, respectively. Thus, the effect of acetonitrile solvation on the enthalpies of formation of the ruthenium precursor complexes is calculated to be ΔH° = −20.7 and −21.3 kcal mol−1. Both of these values are very similar to ΔH° = −22.6 kcal mol−1, computed for solvation of the vanadium precursor complex.

C2 symmetric transition state structures were also located for both ruthenium model complexes. The transition structures are calculated to be higher in enthalpy than the precursor complexes by 10.3 kcal mol−1 (F) and 12.1 kcal mol−1 (pyrr) in the gas-phase and by 11.1 and 12.6 kcal mol−1 in acetonitrile solution. Thus, as in the reaction of vanadium complexes 1c and 2c, solvation by acetonitrile is calculated to have a large enthalpic effect on formation of the ruthenium complexes, but only a very small effect on the barrier to degenerate hydrogen exchange reactions in the ruthenium complexes.

Figure 7 gives the differences between the calculated activation enthalpies for degenerate exchange in the hydrogen-bonded vanadium and ruthenium complexes in both acetonitrile solution and in the gas-phase. Starting either with the isolated reactants or with the hydrogen-bonded complexes, the activation enthalpy is much larger for hydrogen exchange between vanadium complexes 1c and 2c than between RuO(F) and Ru(OH)(F) or RuO(pyrr) and Ru(OH)(pyrr). For example, the enthalpic barrier to hydrogen exchange is 6.0 kcal mol−1 higher for 1c + 2c than for RuO(F) + Ru(OH)(F), even though in MeCN the enthalpies of formation of the hydrogen-bonded precursor complexes are only 0.5 kcal mol−1 smaller for 1c + 2c than for RuO(F) + Ru(OH)(F). The same 6.0 kcal mol−1 difference in barrier heights is also computed for the gas-phase reactions, indicating that the difference is not due to solvation. Nor is the difference in barrier heights due to zero-point energies or to heat capacities, because, if electronic energies are substituted for enthalpies, the 6.0 difference in activation enthalpies is reduced, but by only 0.2 kcal mol−1.

Figure 7
B3LYP/6–31G* PCM enthalpies (kcal mol−1), relative to the hydrogen-bonded complexes (b), of the reactants (a) and transition structures (c) in three degenerate hydrogen exchange reactions in acetonitrile solution. For the transition structures, ...

Discussion

I. Cross reactions involving transfer of a hydrogen atom

Compound 2a abstracts a hydrogen atom from the weak O–H or C–H bonds in a number of organic substrates. Complex 1a and the oxidized organic substrate are formed (Scheme 2, Table 2), although the yields are sometimes low due to the sluggish nature of the reactions and competing decomposition. The reactive substrates have low O-H or C-H bond dissociation free energies (BDFEs66), ranging from 66.5 to 78 kcal mol−1.28,29,51,5357,67,68 The O–H BDFE in 1a was determined to be 70.6 ± 1.2 kcal mol−1 from the equilibrium 2a + tBu2(MeO)ArOH [right harpoon over left harpoon] 1a + tBu2(MeO)ArO•.51

Scheme 2
Summary of the reactivity of 2a with hydrogen atom donors.
Table 2
Driving force and second-order rate constants for transfer of a hydrogen atom from organic substrates to[VVO2(tBu2bpy)2]+ (2a).

These vanadium reactions are remarkably slow. Similar hydrogen transfer reactions (eq 5) have been examined for a wide variety of metal coordination complexes, including Cr, Mn, Fe, Ni, Co, Cu, Ru and Os complexes.13,15,16,2737 These reactions all involve redox change at the metal center, coupled with protonation or deprotonation of a ligand (eq 6).69 A variety of ligands

X+HYXH+Y
(5)

LXMn+(X)+HYLXM(n1)+(XH)+Y
(6)

act as the proton acceptor: oxo (bridging or terminal), alkoxy, superoxo, anilide, hexafluoro-acetylacetonate, imidazolate, imidazolinate, and terpyridine-4-carboxylate. Hydrogen transfer reactions of these complexes are typically rapid at modest driving forces when H• moves between electronegative oxygen or nitrogen atoms, reaching completion in seconds or less at room temperature and common concentrations (k = 101 to 106 M−1 s−1). (Two exceptions are described below.) However, to our surprise, the reactions of the vanadium complexes 1a and 2a with oxyl radicals or O–H bonds are much slower, occurring over periods of hours to days at typical reactant concentrations, with rate constants of 10−4 to 10−1 M−1 s−1.

The mechanisms of the reactions of 1 and 2 could involve concerted (one kinetic step) or stepwise transfer of the two particles (cf. 1,456,70,71). Here, the stepwise paths would require high-energy intermediates. Stable vanadium-oxo compounds are predominantly either vanadium(IV) mono-oxo (vanadyl) or vanadium(V) polyoxo (many are cis-dioxo) complexes.72 Initial electron transfer to 2a would make [VIV(O)2(N-N)2], which is not easily generated from 2a, as described above. Initial proton transfer to 2a would give [VV(O)(OH)(N-N)22+], a high energy species based on its estimated pKa[VVO(OH)(tBu2bpy)2+] of −7 ± 2 (derived from a thermochemical cycle using the O-H BDFE and E1/2 of 1a, Scheme S1).39 In addition, none of the substrates in Scheme 2 are good outer-sphere reductants (E1/2 values from 0.7 to ~1.5 V versus Cp2Fe+/0,5,51c,73,74,75) or strong acids (pKa values from 28 to 4051b,54,76,77). The data thus indicate that initial proton or electron transfer from any of these substrates to 2a would be very unfavorable, and that the reactions must therefore proceed by concerted transfer of e and H+.

II. Pseudo Self-Exchange Reaction

Self-exchange rate constants provide a measure of the intrinsic ability of a reagent to undergo a particular reaction, in this case hydrogen transfer. In the vanadium-oxo system, self-exchange of a hydrogen atom is remarkably slow, based on the pseudo self-exchange reaction between the tBu2bpy and Me2bpy derivatives, 1a + 2b1b + 2a (Scheme 1). This reaction, with k = (1.3 ± 0.1) × 10−2 M−1 s−1 and Keq = 1.0 ± 0.6 in MeCN at 298 K, is much slower than related self-exchange reactions involving NH and OH moieties, as shown in Table 3. The closest analogy is the hydrogen atom self-exchange between RuIV(O)(py)(bpy)22+ and RuIII(OH)(py)(bpy)22+, which occurs at (7.6 ± 0.4) × 104 M−1 s−1.28a This value is 6 × 106times faster than the vanadium self-exchange measured here. The Ru and V systems are both oxo/hydroxo interconversions and both have bis(bipyridine) supporting ligands; they differ only in the sixth ligand (oxo for V, pyridine for Ru) and the d-electron count (d0/1 vs. d4/5). The origin of this difference, and the slowness of the V systems, is discussed below.

Table 3
Measured and estimated NH and OH hydrogen atom self-exchange rate constants, kse for XH + X [right harpoon over left harpoon] X + HX, statistically corrected.a

The VIVO(OH)/VVO2 self-exchange is slower than all but one of the hydrogen atom self-exchange reactions shown in Table 3 by a factor of ca. 105 or more. This large difference is seen in comparisons of reactions involving first-row and second-row metals, with terminal and bridging oxo complexes, in a variety of steric environments, with neutral and multiply charged complexes, and in purely organic reactions. For example, self exchanges between protonated ketyl radicals and ketones, which are similar to VIVO(OH)/VVO2 as Ė–OH + E=O processes, occur with rate constants of (3.7 – 8.6) × 103 M−1 s−1.78 The self-exchange rate constants for other oxo/hydroxo couples, crudely estimated using the Marcus cross relation and measured cross-reaction rate constants, all occur rapidly, with k [greater, similar]103 M−1 s−1. Self-exchange rate constants for reactions involving N–H bonds of metal-containing compounds range from 3.2 × 105 to 9.7 × 102 M−1 s−1,30,31,70 with two exceptions. Self-exchange involving CoII– and CoIII–tris(biimidazoline) complexes is slow, apparently due to the spin-change between low-spin CoIII and high-spin CoII and the attendant large reorganization energy.30a H-transfer between Os(IV)-anilide and Os(III)-aniline complexes is also very slow (kse ≤ 1.5 × 10−3 M−1 s−1), apparently due to strong steric crowding about the aniline hydrogen atoms.31

The very slow VIVO(OH)/VVO2 self-exchange rate could have a number of origins. Hydrogen bond accepting solvents have been shown to significantly slow rates of hydrogen atom transfer, particularly for phenols, by forming an unreactive hydrogen-bonded complex XH (...)solvent.71 However, this is not significant for VIVO(OH)/VVO2, as the self exchange is only 1.7 times faster in methylene chloride than in acetonitrile. This small solvent effect also supports a symmetrical transition structure, with little change in charge, rather than a stepwise mechanism of initial e or H+ transfer.71

Steric crowding can slow rates of hydrogen transfer,31,69c but that seems unlikely here because the O-H and O moieties of VIVO(OH) and VVO2 are uncrowded (see Figure 1). In addition, Ru(OH)(bpy)2(py)2+ + Ru(O)(bpy)2(py)2+ self-exchange is much more facile, despite having more steric congestion around the oxygen atoms.

Small rate constants could also be a result of the vanadium reaction being nonadiabatic, such that the probability of crossing from reactant to product surfaces is small (κ [double less-than sign] 1 in transition state theory).79 This should result in a small Arrhenius prefactor A, which is equivalent to an unusually large negative activation entropy ΔS. However, the values found for 1a + 2a self exchange, A = 4 × 109 s−1, ΔS = −16 ± 5 cal mol−1 K−1, are typical of an adiabatic bimolecular reaction.80 Therefore, large non-adiabatic contributions to the slow self-exchange are unlikely.

The slow vanadium self-exchange rate appears to be due almost entirely to a large activation barrier, ΔH = 15 ± 2 kcal mol−1 (Arrhenius Ea = 15.4 ± 2 kcal mol−1). In terms of Marcus theory, the high barrier indicates an adiabatic self-exchange reaction with a large intrinsic barrier, λ. As discussed below, this interpretation is supported by both the experimental results described in the next section, and by the DFT calculations that follow it.

III. Application of the Marcus Cross Relation

We have shown that the Marcus cross relation (eq 7) holds for many although not all hydrogen transfer reactions, 2831,46b even though it is a simplistic treatment.81Equation 7 predicts the rate constants for XH + Y → X + HY cross reactions (kxy) in terms of its driving force (Keq) and the self-exchange rate constants (kxx, kyy).82,83

kxy=kxxkyyKeqfxy
(7)

Cross rate constants have been calculated for the five vanadium reactions reported here (kcalc) and are compared with the measured kxy in Table 4. As detailed in the notes to Table 4, the Keq values are derived from the relevant BDFEs [except for the directly measured Keq for tBu2(MeO)ArOH], and the kyy values are estimated from pseudo-self exchange reactions, by application of the cross relation in other systems, or from values for related substrates. In applying eq 7, no correction has been made for the work required to assemble the [Ln V–OH(...)O=VLn ]2+ precursor complex;84 including the electrostatic component of this work (wr [congruent with] +1.5 kcal mol−1)39,82,85 would raise the kcalc values by a factor of five. We estimate that the rate constants, kcalc, calculated from eq 7 have an uncertainty of 1–2 orders of magnitude, due to the uncertainties in the BDFEs, in the estimated substrate kyy values, and in the work terms. The kcalc for the reaction of 2a with tBu2(MeO)ArOH is more precise because the Keq was measured directly rather than from BDFEs.

Table 4
Calculated and observed hydrogen atom transfer rate constants (statistically corrected) and equilibrium constants for reactions of VO2(tBu2bpy)22+ (2a) with organic substrates.

The calculated rate constants for transfer of a hydrogen atom are in fair agreement with the experimental values, with kxy/kcalc varying from 0.05 to 80. While some of the values are outside the estimated errors, the cross relation captures the major trends in the rate constants for a range of organic substrates and for values of kxy that range over almost six orders of magnitude. For example, the reaction of 2a with DHA is ~105 times slower than reaction of 2a with hydroquinone, at a similar driving force, because kyy for DHA is much slower than that of hydroquinone.

The Marcus analysis indicates that the slow hydrogen transfer reactions of the vanadium compounds can be traced to the very slow VIVO(OH)+/VVO2+ self-exchange rate constant. For example, 2a oxidizes TEMPOH 104 times more slowly than FeIII(Hbim)(H2bim)22+ (at the same driving force) because hydrogen-transfer self-exchange is much slower for vanadium than iron (H2bim = 2,2′-biimidazoline). All of the slow rates reported here can largely be attributed to an unusually high intrinsic barrier for hydrogen atom self-exchange for VIVO(OH)/VVO2. The origins of this high barrier for hydrogen atom self-exchange are discussed below.

IV. Computational Insights

As described above, DFT calculations at the B3LYP/6-31G(d) level, with a polarizable continuum solvent model appropriate for MeCN, give ΔH of 12.4 kcal mol−1 for the 1c + 2c self-exchange reaction. This calculated value is in good agreement with the experimental activation enthalpy of 15 ± 2 kcal mol−1 for 1a + 2b in MeCN. This agreement supports the conclusion that this self-exchange reaction occurs by a concerted transfer of the electron and proton with a comparatively high enthalpic barrier. Since the calculations did not include any proton tunneling or non-adiabatic effects, the agreement also suggests that such effects are not of great importance in this reaction.

Computational studies have also examined the ruthenium oxo/hydroxo H-atom self-exchange reactions RuIVO(L)(bpy)2+ + RuIIIOH(L)(bpy)2+ for L = fluoride and pyrrolate (pyrr), as models for the experimentally studied dicationic compounds with L = pyridine.28a Starting from the hydrogen-bonded complexes, the calculated enthalpic barriers for RuO(pyrr) + Ru(OH)(pyrr) and RuO(F) + Ru(OH)(F) are 4.2 and 6.0 kcal mol−1 lower, respectively, than for 1c + 2c. These differences in activation enthalpies correspond to differences in rate constants of 103 and 104. Thus the majority of the experimentally observed 106 difference in self-exchange rates between the vanadium and ruthenium complexes is mirrored by our calculations.

A. Origin of the barriers

The difference in calculated electronic energies between the transition structures and hydrogen-bonded complexes is 5.8 kcal mol−1 larger for V than for Ru(O)F. The close correspondence between this ΔΔE and the computed values of ΔΔH = 6.0 kcal mol−1 (both in solution and in the gas-phase) has allowed us to understand the experimental difference between the self-exchange rates of the V and Ru complexes in solution by analyzing the origins of the differences between the electronic energies of the reactions in the gas-phase.

We have divided the passage from the optimized geometries of the hydrogen-bonded precursor complexes to the transition structures into three steps (ad in Figure 8). First, the O(...)O distances in the hydrogen-bonded complexes (a) were shortened to those in the transition structures, but with all other geometrical parameters reoptimized (b). As can be seen in Figure 8 the energies required for this step are the same to within 1.0 kcal mol−1, with the vanadium complex actually being the lowest. The similarity of these energy changes reflects the similar charge and steric environments around the metal oxo moieties in all three complexes.

Figure 8
Relative gas-phase energies (ΔE in kcal mol−1) of: (a) the optimized hydrogen bonded precursor complexes (E set equal to 0); (b) the O(...)O distances shortened to those in the transition structures, but all other geometrical parameters ...

In the next step (bc), all of the atoms except for the transferring hydrogen were moved to their positions in the transition structures, and the positions of the transferring hydrogens were optimized. This step is 5.1 to 5.2 kcal mol−1 more difficult for the vanadium than for the ruthenium complexes, and thus accounts for almost all of the differences in activation energies between the V and Ru compounds. The final step of the cycle involves moving the proton to the symmetrical position that it occupies in the transition structures (cd). Once again the energies required for this step are very similar among the three complexes, being only 1.2 kcal mol−1 more difficult for VO2/VOOH than for RuO(F)/Ru(OH)(F).89

The results in Figure 8 show that the difference between the barrier heights in the vanadium and ruthenium systems arises predominantly from the differences between the energies that are required to change the M=O, M–OH, and M–N bonds from their lengths in the precursor complexes to their lengths in the transition structures. In the Marcus model, the energies associated with these changes are part of the inner-sphere reorganization energies, λi.

Single-point calculations were also performed on the isolated vanadium and ruthenium oxo and hydroxo species, in order to probe the energetic consequences of the bond length changes between the reactant and transition state structures.39 Shortening the M–OH bond lengths of the hydroxo complexes, as occurs in the bc step of Figure 8, is more energetically costly for V(O)(OH)(bpy)2+ (1c) than for Ru(OH)(L)(bpy)2+, by 2.3 and 2.7 kcal mol−1 for L = pyrr and F, respectively. Slightly larger energy differences are calculated for stretching the M=O bonds of the oxo compounds, 3.2 and 3.0 kcal mol−1 greater for 2c than for the RuO(pyrr) and RuO(F) complexes. Together, these distortions are 5.5 – 5.7 kcal mol−1 more difficult for vanadium than for the ruthenium complexes, in good agreement with the calculated differences of 5.1 – 5.2 kcal mol−1 for step (b) to (c) in Figure 8. Thus, the difference in hydrogen exchange barriers, both observed and calculated, can be traced to the greater energetic cost for vanadium than for ruthenium of changing the M=O and M-OH bonds from their lengths in the hydrogen-bonded reactant complexes to their lengths in the transition structures for the self-exchange reactions.

Figure 9 provides the M–OH and M=O bond lengths in the reactants, hydrogen-bonded complexes, and transition structures. The M–OH and M=O bonds are calculated to be 0.1–0.2 Å shorter for M = V than for M = Ru, as expected.90,91 In going from the reactants or the hydrogen-bonded complexes to the transition structures, the changes in M–OH and M=O bond lengths are larger for vanadium than for ruthenium. The M–OH bonds contract by ca. 0.05 Å more for V and the M=O bonds lengthen by ca. 0.04 Å more for V. In addition, the shorter V–OH and V=O bonds have higher force constants. For example, the experimental frequency of the symmetric VO2 stretch is 927 cm−1 in 2avs. ν(Ru=O) = 800 cm−1 in RuIVO(L)(bpy)2+.40,92 The calculated frequencies are νsymm(VO2) = 1060 cm−1 in 2c and ν(Ru=O) = 795 and 798 cm−1 for, respectively, RuO(pyrr) and RuO(F). Therefore, the reorganization energies are higher in the vanadium complexes than in the ruthenium complexes for two reasons -- larger changes in M–OH and M=O bond lengths are required to reach the transition structure for M = V than for M = Ru, and these larger changes are magnified by the higher force constants for the V–O and V=O bonds.

Figure 9
M=O, M-OH, and O-H bond lengths (Å) in the reactants, hydrogen-bonded complexes, and transition structures for three hydrogen atom self-exchange reactions.

The metal-oxygen bond-length distortion energies for moving from the reactant to the transition structure are computed to be greater in the vanadium vs. the ruthenium reactions. This is a consequence of the structure of the reactants: the 0.264 Å difference between the V–OH and V=O bond lengths is ca. 0.10 Å larger than 0.167 – 0.168 Å difference between the Ru–OH and Ru=O bond lengths. Therefore, the key to understanding the greater reorganization energy in the self-exchange reaction of the vanadium versus the ruthenium complexes is understanding why the difference between the M–OH and M=O bond lengths is 50% larger for V than for Ru.

B. Differences between the Reductions of (bpy)2VO2+ (2c) and (bpy)2RuOX+

The larger M=O ➔ M-OH bond length change for M = V than for M = Ru can be understood in terms the differences between the electronic structures of the vanadium and ruthenium reactants and, in particular, the different d-orbital occupancies in these complexes. The vanadium-dioxo compounds 2a–c have a total of three formal metal-oxo π bonds. Two of these bonds come from electron donation from the filled 2pπ atomic orbital (AO) on each oxygen that is perpendicular to the O-V-O plane into an empty t2g orbital on the d0 vanadium(V) cation. The third π bond consists of two partial π bonds, involving donation from the filled 2pπAO on each oxygen that lies in the O-V-O plane, into the third empty t2g orbital, on the d0 vanadium(V) cation.41

Addition of H• (e− + H+) to one of the oxygen atoms in 2a–c causes dramatic changes in π bonding. The added electron goes into the V–O π* orbital that is of δ symmetry relative to the remaining V=O bond (as is typical of [LnVIV(O)] complexes18a,41,42,72a). The π bonding between the vanadium center and the hydroxide ligand is quite weak, as indicated by calculations of the calculated changes in V–O bond distance and spin densities as a function of the O-V-O-H dihedral angle (see Supporting Information). Complex 1c is well-described as having a V[equivalent]O triple bond to the oxo oxygen and a V–O single bond to the hydroxyl group. Formation of 1c by transfer of a hydrogen atom to 2c therefore results in the formal loss of a full V-O π bond.

In contrast, reduction of RuIV(O)L(bpy)2+ complexes results in the loss of only half a π bond. Unlike VV dioxo complexes such as 2c, which have no 3d electrons, ruthenium(IV)–mono-oxo compounds have four electrons in the Ru t2g 4d AOs. Two of these electrons occupy Ru–O π* orbitals.41,91d These complexes therefore have only two half π bonds. Although the total Ru–O π bond order is formally one, each of these half π bonds is expected to be weak.93

Addition of H• to the ruthenium oxo group lowers the formal Ru-O π bond order to 0.5. In Ru(OH)F+, the remaining unpaired electron preferentially occupies a 4d AO that lies in the F-Ru-O plane. A scan of the F-Ru-O-H dihedral angle (see Figure S13 and the accompanying text in the Supporting Information) finds that the global minimum occurs at F-Ru-O-H = 99.8°, where the lone pairs in 2pπ AOs on both F and OH can interact with this singly occupied 4d AO. The scan shows two maxima, at F-Ru-O-H = 357.6° and 232.4°, and a secondary minimum, at F-Ru-O-H = 232.4°. The averaged maxima are 6.9 kcal mol−1 higher in energy than the averaged minima, and have 0.34 Å longer Ru-OH bond lengths. These results indicate that, unlike the case in the vanadium complex 1c, there is a significant, albeit weak, π bond to the OH group in Ru(OH)F. Consistent with this conclusion is the finding that, at the equilibrium geometry of Ru(OH)F, 17.0% of the unpaired spin is localized on the hydroxyl group; whereas, in 1c, only 3.1 % of the unpaired spin appears on the hydroxyl group.

In summary, the calculations provide a detailed understanding of the differences in the heights of the barriers to hydrogen atom self-exchange in the vanadium and ruthenium systems. As illustrated in Figure 10, addition of H• to 2 decreases the π bond order for one V–O bond from 1.5 to close to 0, while the other V–O increases its π bond order from 1.5 to 2. In contrast, addition of H• to the RuIV(O)X complexes results in a bond order decrease of only ca. 0.5 of a weak π bond. The larger changes in π bond order upon hydrogen atom addition to 2 than to RuIV(O)X+ are reflected in larger changes in M–O bond lengths, both observed and computed, for the vanadium than the ruthenium complexes. These larger changes in M-O bond lengths on H• addition to 2, compounded by the higher force constants for the V–O bonds, are the origin of the larger inner-sphere reorganization energies for hydrogen self-exchange in the vanadium vs. the ruthenium complexes.

Figure 10
Summary of changes in π bonding upon addition of a hydrogen atom to 2a–c or to [RuIV(O)(L)(bpy)2].

C. The nature of the e/H+ transfer: PCET vs. HAT

Traditionally, any reaction that involved net transfer of H• from one reagent to another (eq 5 above) has been termed hydrogen atom transfer or HAT. Some years ago, Meyer proposed that HAT should be restricted to processes where “the transferring electron and proton come from the same bond.” 1,94 More recently, some of us showed that H• self-exchange between phenoxyl radical and phenol does not resemble HAT, but rather involves proton transfer between two oxygen lone pairs coupled to electron transfer between phenoxyl π orbitals.95 We termed this mechanism proton-coupled electron transfer (PCET) to distinguish it from HAT, and others have found this distinction useful.71 Unfortunately, PCET now has a much broader usage, often referring to any electron transfer process whose overall rate is affected by proton(s). Truhlar and coworkers have recently proposed an alternative distinction between HAT and PCET processes.96

Using the original definition of PCET, the vanadium self-exchange reaction is a PCET process, not HAT. There is concerted transfer of e + H+ in which the electron is not associated with the VO–H bond. This is shown by the nature of the SOMO at the transition structure (Figure 11). The SOMO is V–OH π antibonding and is, at each vanadium, orthogonal to the O–H bond that is being made or broken. This is the clear indication of a PCET pathway.

Analysis of the ruthenium self-exchange reactions is more complicated because there are three half-filled orbitals at the transition structure (Figure S14). One of these SOMOs resembles the vanadium PCET SOMO but another resembles the SOMO at a classical HAT transition structure. This illustrates a difficulty with the PCET/HAT dichotomy. Only in favorable cases can this distinction be made, even in the simplest case of self-exchange reactions. In cross reactions, the situation is often much more complicated. In the reactions of 2a with xanthene and dihydroanthracene, for instance, the mechanism is PCET from the perspective of the vanadium center but HAT from the perspective of the C–H bond. In many cases it can be quite challenging to make a PCET/HAT distinction in the absence of detailed calculations. Therefore we favor using the term HAT in its traditional way, to “refer to what is transferred between reactants in the net sense and not to the mechanism of the event,”97 unless the context makes clear that a detailed orbital picture is implied.

Conclusions

Reactions that interconvert dioxo-vanadium(V) compounds [VVO2(R2bpy)2]BF4 (R = tBu, Me, H) (2a–c) and oxo-hydroxo-vanadium(IV) compounds [VIVO(OH)(R2bpy)2]BF4 (1a–c)18 have been examined. Together, these compounds comprise a system of sterically uncrowded and isolable metal-oxo and metal-hydroxo compounds that differ only by a hydrogen atom (H+ + e). Compound 2a abstracts a hydrogen atom from 2,6-tBu2-4-MeO-C6H2OH (ArOH), TEMPOH, dihydroanthracene, xanthene, and hydroquinone, while 2,4,6-tBu3C6H2O• abstracts H• from 1a. The reaction of 2a with ArOH reaches a measurable equilibrium, indicating that the VO–H bond dissociation free energy (BDFE) in 1a is 70.6 kcal mol−1. The reactions of 2a with hydrogen atom donors are all surprisingly slow, with second order rate constants from 1 × 10−1 M−1 s−1 to 2 × 10−7 M−1 s−1. Hydrogen atom pseudo-self-exchange between VIVO(OH)(R2bpy)2+ and VVO2(R′2bpy)2+ (R, R′ = tBu, Me) is also surprisingly slow, with kse = 1.3 × 10′2 M′1 s′1 at 298 K. This is 6 × 106 slower than the self-exchange rate constant measured for the sterically similar RuIV(O)(py)bpy22+ + RuIII(OH)(py)bpy22+.27e The activation parameters for vanadium self-exchange, ΔH = 15 ± 2 kcal mol−1 and ΔS = −16 ± 5 cal mol−1 K−1 (A = 4 × 109 s− 1), suggest that this is an adiabatic reaction with a relatively large barrier. Using the vanadium self-exchange rate constant, the Marcus cross relation predicts the rate constants of the HAT cross rates within two orders of magnitude in most cases.

DFT calculations of the vanadium self-exchange reaction 1c + 2c, using a PCM solvent model, give a calculated barrier height of ΔH = 12.4 kcal mol−1 above the separated reactants and 17.1 kcal mol−1 above the hydrogen-bonded precursor complex [(bpy)2(O)V–OH (...) O=V(O)–(bpy)2]2+. Passage to the transition structure involves large changes in the V–O bond lengths. These distortions are larger than the corresponding bond length changes in model ruthenium complexes RuIV(O)(X)bpy2+/RuIII(O)(X)bpy2+ (X = F, pyrrolate). The difference between the vanadium and ruthenium complexes is due to a larger change in the net π bond order in the former than in the latter, which can, in turn, be traced to the differing d-orbital occupancies of the V vs. Ru complexes. The larger bond length distortions and higher force constants make the Marcus inner sphere reorganization energy, λi, larger for the vanadium than for the ruthenium reactions, leading to the dramatic 106 difference in reactivity.

We have previously emphasized the importance of driving force in understanding hydrogen atom transfer reactivity. This study emphasizes that differences in intrinsic barriers heights can have a dramatic effect as well.

Experimental Section

Reagent grade chemicals and solvents were obtained from Aldrich, Eastman Organic, Strem, Fisher Scientific or EMD Chemicals and used without further purification unless otherwise noted. CH2Cl2 was dried using a “Grubbs type” Seca Solvent System installed by GlassContour.98 MeCN (Burdick and Jackson low water brand) was stored in an argon-pressurized steel drum plumbed directly into a glovebox. Deuterated solvents were purchased from Cambridge Isotope Laboratories. CD2Cl2 was dried by stirring over CaH2 followed by vacuum transfer and stored in a dark bottle in a glovebox freezer. CD3CN was dried over CaH2, vacuum transferred to P2O5 for 30 min, then CaH2 again, and transferred to a clean flask and stored in a glovebox. TEMPOH29 and tBu3PhO•53 were prepared following literature procedures. tBu3PhOH and tBu2(MeO)PhOH were recrystallized twice from hexanes or ethanol, respectively. Hydroquinone was recrystallized from MeCN.

The vanadium compounds were synthesized under air, dried in vacuo, and stored in a N2 glovebox. All other reactions were performed under N2 using standard glovebox and Schlenk techniques, often in sealable (J-Young) NMR tubes or quartz cuvettes fitted with Teflon stopcock. 1H and 13C NMR spectra were recorded on Bruker Avance or DRX spectrometers (300 and 500 MHz) at ambient temperatures unless otherwise noted and are referenced to a solvent peak. Electrospray ionization mass spectra (ESI/MS) were obtained on a Bruker Esquire-LC ion trap mass spectrometer as solutions in acetonitrile. IR spectra were recorded as KBr pellets or CH2Cl2 solutions in a cell with NaCl windows (as noted) using either a Perkin-Elmer 1720 or a Bruker Vector 33 FT-IR spectrometer and are reported in cm−1. UV-vis spectra were recorded using a Hewlett-Packard 8453 spectrometer and are reported as λ in nm (Σ, M−1 cm−1). Experimental details for the X-ray structure of 2c are given in the Supporting Information.

Syntheses

The preparations of the tBu2bpy derivatives are given here; the similar procedures for the Me2bpy and bpy complexes are given in the Supporting Information.

cis-[VIVO(OH)(4,4′-tBu2bpy)2]BF4 (1a)

In a simpler procedure than the literature method,18b aqueous HBF4 (48 wt %, ~0.3 mL, 2.3 mmol) was added to a solution of V(O)SO4•2(H2O) (260 mg, 1.33 mmol) in 2 mL deionized H2O, to make the pH 1.0 (by pH meter). BaCO3 (0.270 g, 1.36 mmol) was added, evolution of CO2 was observed, and the white precipitate (BaSO4) was removed by filtration through Celite. Upon addition of tBu2bpy (870 mg, 3.24 mmol), the clear blue solution changed to red-brown and a yellow precipitate formed. Stirring for 2 h and filtration gave a light yellow solid, which was washed with Et2O and dried in vacuo to give 752 mg (80%) of 1a. Crystals of 1a were obtained by slow evaporation of a CH2Cl2/toluene (1:1 v/v) solution. 1H NMR (CD3CN): δ 1.365, 1.584 1.672, 2.705 (br, tBu, 9H each). IR (KBr pellet): 969 cm−1 vs v(V=O), 1615 vs v(C=C, C=N), 2961 vs v(C–H). UV-Vis (CH3CN): 753 (26), 521 (37), 301 (24620), 261 (11875). ESI/MS (CH3CN): 622 [1a + 2H - BF4]+, 619 [1a]+.

cis-[VVO2(4,4′-tBu2bpy)2]BF4 (2a)

tBu2bpy (2.4 g, 8.96 mmol) was added to a solution of NaVO3 (0.546 g, 4.48 mmol) in 75 mL aqueous HBF4 (pH = 1). Stirring overnight yielded a pale yellow precipitate, which was collected by filtration. The precipitate was dissolved in CH2Cl2 and the solution was washed three times with aqueous NaHCO3, then with water. The volume was reduced to ~20 mL and 200 mL pentane added dropwise to give a fine yellow precipitate that was collected by filtration and dried in vacuo (2.812 g, 89% yield). Anal. Calcd for 2a•½CH2Cl2 (CH2C12 was observed in these samples by NMR), C36.5H49BC1F4N4O2V: C, 58.53; H, 6.59; N, 7.48. Found: C, 58.7; H, 6.61; N, 7.48. 1H NMR (CD2C12205 K): tBupy A: δ1.28 (s, tBu, 18H); 7.34 (4H, 6-bpy and 5-bpy); 8.13 (s, 2H, 3-bpy); tBupy B: 1.46 (s, tBu, 18H); 7.73 (d, 5.5 Hz, 2H, 5′-bpy H); 8.29 (s, 2H, 3′-bpy H); 8.91 (d, 5.5 Hz, 2H, 6′-bpy H). 13C{1H} NMR (CD2C12 205 K): δ29.65, 29.71 (C(CH3)3); 36.39, 36.03 (CMe3); 119.72, 119.85, 124.84 (2C), 146.48, 149.35, 153.40, 156.06, 156.06, 164.63, 168.21 (bpy). At 298 K in CD2C12 or CD3CN, one broad lBu peak is observed at 1.4 ppm and three very broad bpy peaks are observed at 7.5, 8.3 and 8.9 ppm. ESI/MS (CH3CN): m/z 619 [2a - BF4]+, 87 [BF4] IR (CH2C12 soln): 927, 906 s v(V=0), 1618, 1546 vs v(C=C, C=N), 2973 s ν(C-H). UV-Vis: (CH3CN) 305 (20,000).

Pseudo Self-Exchange Experiments

In a typical experiment, a septum-sealed NMR tube in a N2 glovebox was charged with 2a (500 µL, 17 mM in CD3CN) and (Me3Si)2O (~lµL) as an internal standard. A gastight syringe was loaded with 300 µL of a solution of lb (8 mM in CD3CN). The syringe and NMR tube were then removed from the glovebox and the NMR tube frozen in liquid nitrogen. The contents of the syringe were then added to the NMR tube, which was thawed immediately before inserting into the NMR which was pre-equilibrated at the desired temperature. Spectra were recorded until equilibrium was reached. Reactions starting with solutions of la and 2b were also performed.

The reaction of 2a + ArOH

was monitored by optical spectroscopy (350 – 900 nm). Concentrations at equilibrium were calculated using the known spectra of la and ArO• in MeCN and assuming mass balance according to equation 4. The kinetics were analyzed using SPECFIT52 with a second-order approach to equilibrium model; the vanadium spectra and K4 were fixed. The calculated spectra of ArO• were in good agreement with experiment. Initial concentrations were varied from 1.1 – 23 mM (2a) and 3 – 300 mM (ArOH).39The kinetics of 2a + xanthene were measured using 5 to 20 equivalents of xanthene, and modeled as a second order process in SPECFIT with the spectrum of la fixed.

Solutions of 2,6-tBu2-4-(MeO)C6H2O• (ArO•)

were generated via oxidation of ArOH with NaOH and K3[FeIII(CN)6] in a biphasic benzene/water mixture.53 The resulting 0.2 M solution in C6H6 contained ~ 0.1 to 0.2 M phenol impurities, including ArOH. Extinction coefficients were measured by titrating solutions of ArO• with known amounts of TEMPOH. UV-Vis: 535 (230), 405 (4000), 389 (2800), 336 (5800). The λmax values are consistent with those reported by Modarelli et al. in pentane, 99 but the molar absorptivities measured here are larger.

Supplementary Material

1_si_001

Supporting Information Available: Syntheses and characterization of compounds 1b–c and 2b–c; crystallographic data for 2c and Figure S1Figure S14; geometries and energies of all the calculated stationary points; discussion of the bonding in 1c and Ru(OH)F(bpy)2+ as a function of O-M-O-H torsion angle; and the complete list of authors for ref. 61 (for Gaussian03). This material is available free of charge via the Internet at http://pubs.acs.org.

2_si_002

Acknowledgement

We thank the U.S. National Institutes of Health (GM50422) and the University of Washington for financial support to JMM, and the National Science Foundation and Robert A. Welch Foundations for support to WTB. We also thank Dr. Martin Sadilek for assistance with mass spectrometry.

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81. See, for instance, discussions in references 46b, 79, 95, and 96.
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83. fxy82 is usually close to 1.
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89. (a) We also performed calculations with the MPW1K functional.89b The energies for all of the steps in Figure 8 computed with MPW1K/6–31+G(d,p) were similar to those computed with B3LYP/6–31G(d), except for the hydrogen transfer step (c→d). The MPW1K calculations gave values about 10 kcal mol−1 higher for the vanadium complex and 7–8 kcal mol−1 higher for the ruthenium complexes. The barrier heights computed with MPW1K were higher than B3LYP values by 12.8 kcal mol−1 for VO2/V(O)OH and by 11.6 and 10.4 kcal mol−1 for RuO(pyrr)/Ru(OH)(pyrr) and RuO(F)/Ru(OH)(F) respectively. Probably through a fortuitous cancellation of errors, the B3LYP barrier heights are in much better agreement with experiment than are the MPW1K barrier heights. Lynch BJ, Fast PL, Harris M, Truhlar DG. J. Phys. Chem. A. 2000;104:4811.
90. The V-OH bond of 1c is ~0.2 Å shorter than a typical RuIII-OH bond [a structure of RuIIIOH(bpy)2(py)2+ has not been reported]: (a) Man ML, Zhu J, Ng SM, Zhou Z, Yin C, Lin Z, Lau CP. Organometallics. 2004;23:6214–6220. (b) Jitsukawa K, Oka Y, Yamaguchi S, Masuda H. Inorg. Chem. 2004;43:8119–8129. [PubMed] (c) Takahashi Y, Hikichi S, Moro-oka Y, Akita M. Polyhedron. 2004;23:225–234. (d) Nagao H, Aoyagi K, Yukawa Y, Howell FS, Mukaida M, Kakihana H. Bull. Chem. Soc. Jpn. 1987;60:3247–3245. (e) Gibson DH, Pariya C, Mashuta MS. Organometallics. 2004;23:2510–2513. (f) Gemel C, Mereiter K, Schmid R, Kirchner K. Organometallics. 1997;16:5601–5603. (g) Schneider R, Weybermüller T, Wieghardt K. Inorg. Chem. 1993;32:4925–4934. (h) Anderson C, Beauchamp AL. Inorg. Chem. 1995;34:6065–6073.
91. V=O bonds62 are ca. 0.16 Å shorter than corresponding RuIV=O bonds [a structure of RuIVO(bpy)2(py)2+ has not been reported]: (a) Che C, Lai T, Wong K. Inorg. Chem. 1987;26:2289–2299. and refs. therein. (b) Che C, Tang W, Lee W, Wong W, Lai T. J. Chem. Soc. Dalton. Trans. 1989:2011–2016. (c) Wong K, Che C, Yip W, Wang R, Mak TCW. J. Chem. Soc. Dalton. Trans. 1992:1417–1421.
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93. In a three-electron π bond, the electron that occupies the antibonding orbital more than cancels the bonding contributed by one of the electrons in the bonding orbital. See, for example, Jorgensen WL, Borden WT. J. Am. Chem. Soc. 1973;95:6649.
94. Binstead RA, McGuire ME, Dovletoglou A, Seok WK, Roecker LE, Meyer TJ. J. Am. Chem. Soc. 1992;114:173–186. (a) See also reference 1.
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97. Binstead RA, Meyer TJ. J. Am. Chem. Soc. 1987;109:3287–3297. The emphasis is in the original. This definition of HAT precedes Meyer’s different definition (quoted above) in references 1 and 94.
99. Modarelli DA, Rossitto FC, Lahti PM. Tetrahedron Lett. 1989;30:4473–4476.