Comparative DFT Analysis of Ligand and Solvent Effects on the

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Organometallics 2010, 29, 5121–5131 DOI: 10.1021/om100326z

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Comparative DFT Analysis of Ligand and Solvent Effects on the Mechanism of H2 Activation in Water Mediated by Half-Sandwich Complexes [Cp0 Ru(PTA)2Cl] (Cp0 = C5H5, C5Me5; PTA = 1,3,5-triaza-7-phosphaadamantane)† G abor Kov acs,† Andrea Rossin,‡ Luca Gonsalvi,‡ Agustı´ Lled os,*,† and ,‡ Maurizio Peruzzini* †

Departament de Quı´mica, Universitat Aut onoma de Barcelona, Bellaterra, Barcelona, Spain, and Consiglio Nazionale delle Ricerche, Istituto di Chimica dei Composti Organometallici (ICCOM-CNR), Via Madonna del Piano 10, 50019 Sesto Fiorentino (Firenze), Italy

‡

Received April 21, 2010

The activation of hydrogen by complexes [Cp0 Ru(PTA)2Cl] (Cp0 = C5H5, C5Me5; PTA = 1,3,5triaza-7-phosphaadamantane) in water was investigated in a comparative DFT study carried out using a discrete þ continuum model based on a four-water-molecule cluster. As a starting point were chosen the η2-H2 dihydrogen complexes [Cp0 Ru(PTA)2(η2-H2)]þ (1H2), which are formed initially upon reaction of the chloride precursors with hydrogen gas. A rationale for the experimental data, showing that the monohydrido complex [CpRu(PTA){PTA(H)}H]þ (4a) and the dihydrido complex [Cp*Ru(PTA)2(H)2]Cl (3b) are the stable products for the two systems, is proposed, together with an in-depth analysis of both ligand and solvent effects in the stability of the different species, leading to more general mechanistic implications for metal-mediated hydrogen activation in water.

Introduction The activation of hydrogen is a topic of great interest for both fundamental studies and practical applications. During the last four decades, extensive research worldwide has been focused on this area,1,2 including studies on transition metal hydrogenation catalyst design and testing, structure and bonding analysis, mechanistic investigations, and theoretical calculations.3,4

The interaction of hydrogen with a transition metal fragment can lead either to dihydrogen M(η2-H2) coordination without H-H bond cleavage or to complete dissociation of the H2 molecule with concomitant dihydrido complex M(H)2 formation. If the dihydrogen ligand initially forming a nonclassical complex is acidic enough, it can be deprotonated by a suitable base, and a monohydride can be formed through heterolytic cleavage of the H-H bond.5,6 Further protonation at the metal center can also give a dihydrido complex.7 The multifaceted reactivity observed is largely dependent on the ancillary ligands, electron-donating groups favoring dihydride formation. The acidity of dihydrogen complexes is also strongly affected by the auxiliary ligands, decreasing as they become more electron-donating. Extensive studies of ligand effects on dihydrogen reactivity have been performed on complexes of general formula [(C5R5)MH2(diphosphine)]þ (M = Fe, Ru, Os; R = H, Me),8 proving that they can exist in either pure dihydrogen or pure trans-dihydride form, depending on the metal, the R substituent on the cyclopentadienyl ligand, and the stereolectronic nature of the diphosphines. Systematic studies of

† Part of the Dietmar Seyferth Festschrift. Dedicated to our colleague and friend Dietmar Seyferth, a pioneer of organometallic chemistry and a giant Editor-in-Chief of Organometallics. *To whom correspondence should be addressed. E-mail: agusti@ klingon.uab.es (A.L.); [email protected] (M.P.). (1) (a) Halpern, J. J. Phys. Chem. 1959, 63, 398–403. (b) Kubas, G. J.; Ryan, R. R.; Swanson, B. I.; Vergamini, P. J.; Wasserman, H. J. J. Am. Chem. Soc. 1984, 106, 451–452. (c) Lee, J. C., Jr.; Rheingold, A. L.; Muller, B.; Pregosin, P. S.; Crabtree, R. H. J. Chem. Soc., Chem. Commun. 1994, 1021– 1022. (d) Morris, R. H.; Schlaf, M. Inorg. Chem. 1994, 33, 1725–1726. (2) (a) Kubas, G. J. Chem. Rev. 2007, 107, 4152–4205. (b) Kubas, G. K. J. Organomet. Chem. 2009, 694, 2648–2653. (3) (a) James. B. R. Homogeneous Hydrogenation; Wiley: New York, 1973. (b) Jessop, G.; Morris, R. H. Coord. Chem. Rev. 1992, 121, 155–284. (c) Heinekey, D. M.; Oldham, W. J., Jr. Chem. Rev. 1993, 93, 913–926. (d) Crabtree, R. H. Angew. Chem., Int. Ed. Engl. 1993, 32, 789–805. (e) Stanislaus, A.; Cooper, B. H. Catal. Rev. 1994, 36, 75–123. (f ) Morris, R. H. Can. J. Chem. 1996, 74, 1907–1915. (g) Torrent, M.; Sola, M.; Frenking, G. Chem. Rev. 2000, 100, 439–493. (h) Heinekey, D. M.; Lledos, A.; Lluch, J. M. Chem. Soc. Rev. 2004, 33, 175–182. (i) Kubas, G. J. Catal. Lett. 2005, 104, 79–101. ( j) The Handbook of Homogeneous Hydrogenation; de Vries, J. G.; Elsevier, C. J., Eds.; Wiley-VCH: Weinheim, Germany, 2006. (4) (a) Recent Advances in Hydride Chemistry; Peruzzini, M.; Poli, R., Eds.; Elsevier SA: Amsterdam, The Netherlands, 2001. (b) Metal Dihydrogen and σ-Bond Complexes; Kubas, G. J., Eds.; Kluwer/Academic Plenum Publishers: New York, 2001.

(5) (a) Maseras, F.; Lled os, A.; Clot, E.; Eisenstein, O. Chem. Rev. 2000, 100, 601–636. (b) Niu, S.; Hall, M. B. Chem. Rev. 2000, 100, 353– 405. (c) Ienco, A.; Calhorda, M. J.; Reinhold, J.; Reineri, F.; Bianchini, C.; Peruzzini, M.; Vizza, F.; Mealli, C. J. Am. Chem. Soc. 2004, 126, 11954– 11965. (6) Kubas, G. J. Adv. Inorg. Chem. 2004, 56, 127–177. (7) Besora, M.; Lled os, A.; Maseras, F. Chem. Soc. Rev. 2009, 38, 957–966. (8) (a) Jia, G.; Lau, C.-P. Coord. Chem. Rev. 1999, 190-192, 83–108. (b) Jimenez-Tenorio, M.; Puerta, M. C.; Valerga, P. Eur. J. Inorg. Chem. 2004, 17–32. (c) Morris, R. H. Coord. Chem. Rev. 2008, 252, 2381–2394.

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ligand effects on η2-H2 acidity are also available for [(C5R5)Ru(H)2L2]þ complexes (R = H, Me; L = monodentate or bidentate phosphines).9 In order to analyze the details of ligand effects in hydrogen activation, great potentiality lies in the possibility to compare complexes differing only in the R groups of the cyclopentadienyl ligands: it has been established that the replacement of Cp (R = H) by Cp* (R = Me) decreases both the stability of the dihydrogen form compared with that of the dihydride tautomer and its acidity.9 The results so far discussed were obtained mainly by studies carried out in organic solvents. Studies on the nature and activation of hydrogen have moved their focus in recent years to alternative media, in particular water, which is of interest in replacing current processes with more environmentally friendly protocols carried out in such medium.10 While much is known about the reactivity of dihydrogen complexes in organic phases, very few water-soluble dihydrogen complexes have been identified yet.10-12 Several intermediates containing the η2-H2 moiety in water have been proposed, particularly to explain the H/D exchange in D2O as well as potential intermediates in hydrogenase enzymes.13 However, a clear picture on how water can affect the dihydrogen reactivity and how ligand effects are influenced by the aqueous environment has not been attained yet. Some of us have reported on the chemoselective catalytic reduction of organic substrates in biphasic H2O/n-octane systems using water-soluble Ru(II) complexes such as [Cp0 Ru(PTA)2Cl] (Cp0 = C5H5, Cp, 1a; Cp0 = C5Me5, Cp*, 1b; PTA = 1,3,5-triaza-7-phosphaadamantane).14a Formation of metal hydrides was observed by HPNMR techniques under 30 bar of H2 and temperatures ranging from 45 to 80 °C, allowing for a comparative study of the Cp-Cp* substitution effects on H2 activation in water. In these experiments it was observed that, whereas for 1a the monohydrido complex [CpRu(PTA)2H] (2a) is the stable product in the whole temperature range, in the Cp* system the dihydrido complex [Cp*Ru(PTA)2(H)2]Cl (3b) initially formed at 45 °C is found to convert to the monohydride analogue of 2a, namely, [Cp*Ru(PTA)2H] (2b), only upon heating to 80 °C (Scheme 1).15 The nature of monohydrides 2a,b was confirmed by independent syntheses with NMR and IR characterization. Later, Frost et al. showed that by (9) (a) Jia, G.; Morris, R. H. J. Am. Chem. Soc. 1991, 113, 875–883. (b) Jia, G.; Lough, A. J.; Morris, R. H. Organometallics 1992, 11, 161–171. (10) (a) Szymczak, N. K.; Tyler, D. R. Coord. Chem. Rev. 2008, 252, 212–230. (b) Aqueous-Phase Organometallic Catalysis: Concepts and Applications, 2nd ed.; Cornils, B.; Herrmann, W. A., Eds.; Wiley-VCH: Weinheim, Germany, 2004. (11) (a) Li, Z. W.; Taube, H. J. Am. Chem. Soc. 1991, 113, 8946–8947. (b) Aebischer, N.; Frey, U.; Merbach, A. E. Chem. Commun. 1998, 2303– 2304. (12) (a) Szymczak, N. K.; Zakharov, L. N.; Tyler, D. R. J. Am. Chem. Soc. 2006, 128, 15830–15835. (b) Gilbertson, J. D.; Szymczak, N. K.; Crossland, J. L.; Miller, W. K.; Lyon, D. K.; Foxman, B. M.; Davies, J.; Tyler, D. R. Inorg. Chem. 2007, 46, 1205–1214. (c) Szymczak, N. K.; Braden, D. A.; Crossland, J. L.; Turov, Y.; Zakharov, L. N.; Tyler, D. R. Inorg. Chem. 2009, 48, 2976–2984. (13) (a) Zhao, X.; Georgakaki, I. P.; Miller, M. L.; Mejia-Rodriguez, R.; Chiang, C.-Y.; Darensbourg, M. Y. Inorg. Chem. 2002, 41, 3917– 3928. (b) Heinekey, D. M. J. Organomet. Chem. 2009, 694, 2671–2680. (14) (a) Bola~ no, S.; Gonsalvi, L.; Zanobini, F.; Vizza, F.; Bertolasi, V.; Romerosa, A.; Peruzzini, M. J. Mol. Catal. A: Chem. 2004, 224, 61– 70. For reviews on PTA chemistry, see: (b) Phillips, A. D.; Gonsalvi, L.; Romerosa, A.; Vizza, F.; Peruzzini, M. Coord. Chem. Rev. 2004, 248, 955– 993. (c) Bravo, J.; Bola~ no, S.; Gonsalvi, L.; Peruzzini, M. Coord. Chem. Rev. 2010, 254, 555–607. (15) Akbayeva, D. N.; Gonsalvi, L.; Oberhauser, W.; Peruzzini, M.; Vizza, F.; Br€ uggeller, P.; Romerosa, A.; Sava, G.; Bergamo, A. Chem. Commun. 2003, 264–265.

Kov acs et al. Scheme 1

running similar NMR experiments at lower temperature (5 °C) the stable cationic hydride [CpRu(PTA){PTA(H)}H]þ (4a), containing one N-protonated PTA ligand, formed via a formal intramolecular base-assisted hydrogen activation, with one PTA molecule acting as a base through one of its three nitrogen atoms.16 The monohydride 2a was also synthesized by the reaction of 1a with KOH in THF at low temperature and characterized by single-crystal XRD analysis.16a Recently, we have analyzed by theoretical methods the mechanistic profile of the H-H bond splitting by 1a in aqueous solution.17 The aqueous medium was represented by discrete-continuum modeling using small clusters of three water molecules in the optimizations and introducing the effect of the bulk solvent by a continuum method. The process starts with ligand substitution of chloride at first by H2O, which, in turn, is quickly displaced by H2 in 1a, resulting in the formation of [CpRu(PTA)2(η2-H2)]þ (1aH2). Water-assisted, concerted deprotonation of the H2 ligand in 1aH2 and proton transfer to the PTA ligand gives the protonated complex [CpRu(PTA){PTA(H)}H]þ (4a) upon overcoming a very low energy barrier. The formation of [CpRu(PTA)2(H)2]þ (3a) was characterized by a higher barrier than that found for the previous step, and it is therefore unfavored. Thus, the monohydride 4a appears as the most stable species in the overall process of H-H bond splitting in water. The computational study was in agreement with the experimental results obtained for the Cp complex, where no formation of the dihydride 3a was observed.15,16 In contrast to this behavior, the formation of the dihydrido complex [Cp*Ru(PTA)2(H)2]þ (3b) was observed when 1b was used instead of 1a under the same experimental conditions.15 Water-soluble ruthenium PTA complexes with C5R5 ligands are therefore very well suited to analyze ligands and solvent effects in the H2 activation in water. Hereby we complete our previous theoretical study by comparing the energy profiles and the mechanisms ruling either H-H homolytic or heterolytic splitting in water for [CpRu(PTA)2(η2-H2)]þ (1aH2) and [Cp*Ru(PTA)2(η2-H2)]þ (1bH2), respectively. The possible pathways for the formation of compounds 3 and 4 from 1 (Scheme 2) have been analyzed. The description of the discrete water environment has been improved by enlarging the water cluster from three to four water molecules, creating a computational model closer to the real system. (16) (a) Frost, B. J.; Mebi, C. A. Organometallics 2004, 23, 5317–5323. (b) Mebi, C. A.; Frost, B. J. Organometallics 2005, 24, 2339–2346. (c) Mebi, C. A.; Nair, R. P.; Frost, B. J. Organometallics 2007, 26, 429–438. (17) Rossin, A.; Gonsalvi, L.; Phillips, A. D.; Maresca, O.; Lled os, A.; Peruzzini, M. Organometallics 2007, 26, 3289–3296.

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Computational Details The methodology employed is very similar to that used in our previous study describing the hydrogen activation by 1a.17 Density functional theory calculations were carried out to identify the structures of the reaction intermediates and transition states of the H-H bond splitting process. All DFT calculations used the program package Gaussian0318 and the B3LYP functional.19 The present work concerns the analysis of the solvent intervention in the H-H bond breaking process. For this purpose a discrete þ continuum modeling of the reaction medium was employed, including solvent effects both by the explicit inclusion in the computational system of discrete water molecules and by expressing the bulk solvent effects through the polarizable continuum model (PCM).20 This water cluster (discrete) þ bulk solvent (continuum) approach has been successfully applied to the theoretical modeling of ruthenium-catalyzed hydrogenations in water21 and more recently to the protonation of rhenium trans-dioxo complexes in aqueous solution.22 (18) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Montgomery, J. A., Jr.; Vreven, T.; Kudin, K. N.; Burant, J. C.; Millam, J. M.; Iyengar, S. S.; Tomasi, J.; Barone, V.; Mennucci, B.; Cossi, M.; Scalmani, G.; Rega, N.; Petersson, G. A.; Nakatsuji, H.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Klene, M.; Li, X.; Knox, J. E.; Hratchian, H. P.; Cross, J. B.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Ayala, P. Y.; Morokuma, K.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Zakrzewski, V. G.; Dapprich, S.; Daniels,A. D.; Strain, M. C.; Farkas, O.; Malick, D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Ortiz, J. V.; Cui, Q.; Baboul, A. G.; Clifford, S.; Cioslowski, J.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi, I.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.; Challacombe, M.; Gill, P. M. W.; Johnson, B.; Chen, W.; Wong, M. W.; Gonzalez, C.; Pople, J. A. Gaussian 03, revision C.02; Gaussian, Inc.: Wallingford, CT, 2004. (19) (a) Lee, C. T.; Yang, W. T.; Parr, R. G. Phys. Rev. B 1988, 37, 785–789. (b) Becke, A. D. J. Chem. Phys. 1993, 98, 5648–5652. (c) Stephens, P.; Devlin, F.; Chabalowski, C.; Frisch, M. J. Phys. Chem. 1994, 98, 11623–11627. (20) (a) Miertus, S.; Scrocco, E.; Tomasi, J. J. Chem. Phys. 1981, 55, 117–129. (b) Miertus, S.; Tomasi, J. J. Chem. Phys. 1982, 65, 239–245. (c) Tomasi, J.; Persico, M. Chem. Rev. 1994, 94, 2027–2094. (d) Cossi, M.; Barone, V.; Cammi, R.; Tomasi, J. Chem. Phys. Lett. 1996, 255, 327–335. (21) (a) Kov acs, G.; Schubert, G.; Jo o, F.; Papai, I. Organometallics 2005, 24, 3059–3065. (b) Kovacs, G.; Ujaque, G.; Lledos, A.; Joo, F. Organometallics 2006, 25, 862–872. (c) Rossin, A.; Kovacs, G.; Ujaque, o, F. Organometallics 2006, 25, 5010–5023. G.; Lledos, A.; Jo (22) Gancheff, J. S.; Kremer, C.; Denis, A.; Giorgi, C.; Bianchi, A. Dalton Trans. 2009, 8257–8268.

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The optimization of minima and transition states was carried out using the water cluster model incorporating four water molecules. In this cluster extension the proton-accepting water molecule is hydrogen-bound to two other molecules, providing a more realistic representation of the experimental aqueous environment, where a complex H-bonding network in the solvent is present. Frequency calculations were performed on all the optimized structures in order to characterize the stationary points (both minima and transition states) and to calculate the gas-phase Gibbs energies at 298 K. For each transition structure the intrinsic reaction coordinate (IRC) routes toward the corresponding minima were calculated.23 If the IRC calculations failed to reach the energy minima on the potential energy surface, geometry optimizations from the final phase of the IRC path were additionally performed. Optimization and characterization of the stationary points were performed with basis set 1 (BS1). In this basis set, the LANL2DZ pseudopotential was employed for the core electrons of the ruthenium center,24 whereas its valence electrons were described by the LANL2DZ basis set.24 The standard 6-31G(d) basis set was used for all the other atoms.25 An extended hydrogen-bonded network involving the four water molecules and the basic centers of the complexes was found in all the optimized structures. To obtain a better energy estimation, single-point calculations were performed on the BS1-optimized geometries using a more extended LANL2DZþf/6-31G(d,p) basis set (BS2).18,25 In basis set 2 polarization functions were added to all the atoms, including a series of f-functions on ruthenium.26 The effect of the bulk medium was considered through the application of the polarizable continuum model as implemented in Gaussian 03 (εwater = 78.39). The Gibbs energies in water (ΔGwater) were obtained adding the Gibbs energies in the gas phase to the contribution of the Gibbs energy of solvation from the continuum model. All the Gibbs energy values (ΔGwater) in the text are obtained from single-point calculations carried out with the extended (BS2) basis set on the BS1-optimized structures. In order to compare the acidity of Cp and Cp* dihydrogen and dihydride species in water, pKawater values have been estimated using the cluster approach.27 In this methodology few explicit water molecules are included in the first solvation shell of the proton and of the protonated and deprotonated species (see below for more details). It has been shown that the use of clusters gives solvation Gibbs energies of charged solutes in excellent agreement with the experimental values obtained from thermodynamic studies of ion-water clusters.27a This explicit-implicit solvation approach has been applied to the calculation of the pKa of several organic acids,28 in both water (23) (a) Fukui, K. Acc. Chem. Res. 1981, 14, 363–368. (b) Gonzalez, C.; Schlegel, H. B. J. Chem. Phys. 1989, 90, 2154–2161. (c) Gonzalez, C.; Schlegel, H. B. J. Phys. Chem. 1990, 94, 5523–5527. (24) (a) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 270–283. (b) Wadt, W. R.; Hay, P. J. J. Chem. Phys. 1985, 82, 284–298. (25) (a) Hehre, W. J.; Ditchfield, R.; Pople, J. A. J. Chem. Phys. 1972, 56, 2257–2261. (b) Hariharan, P. C.; Pople, J. A. Theor. Chim. Acta 1973, 28, 213–222. (c) Francl, M. M.; Pietro, W. J.; Hehre, W. J.; Binkley, J. S.; DeFrees, D. J.; Pople, J. A.; Gordon, M. S. J. Chem. Phys. 1982, 77, 3654– 3665. (d) Rassolov, V. A.; Pople, J. A.; Ratner, M. A.; Windus, T. L. J. Chem. Phys. 1999, 109, 1223–1229. (26) Ehlers, A. W.; Boehme, M.; Dapprich, S.; Gobbi, A.; Hoellwarth, A.; Jonas, V.; Koehler, K. F.; Stegmann, R.; Veldkamp, A.; Frenking, G. Chem. Phys. Lett. 1993, 208, 111–114. (27) (a) Bryantsev, V. S.; Diallo, M. S.; Goddard, W. A., III. J. Phys. Chem. B 2008, 112, 9709–9719. (b) Creati, F.; Coletti, C.; Re, N. Organometallics 2009, 28, 6603–6616. (28) (a) Pliego, J. R., Jr.; Riveros, J. M. J. Phys. Chem. A 2002, 106, 7434–7439. (b) Fu, Y.; Liu, L.; Li, R.-Q.; Liu, R.; Guo, Q.-X. J. Am. Chem. Soc. 2004, 126, 814–822. (c) Kelly, C. P.; Cramer, C. J.; Truhlar, D. G. J. Phys. Chem. A 2006, 110, 2493–2499. (d) Ho, J.; Coote, M. L. J. Chem. Theor. Comput. 2009, 5, 295–306. (e) Ding, F.; Smith, J. M.; Wang, H. J. Org. Chem. 2009, 74, 2679–2691.

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Figure 1. Relevant structures in the H-H homolytic splitting pathway of 1bH2. H atoms in the Cp* and PTA ligands have been omitted for clarity. In parentheses relative Gibbs energies of the structures (kcal mol-1). and nonaqueous solvents, giving results in good agreement with the experimental values, and it has been recently applied also to the estimation of pKa of transition metal complexes.27

Results and Discussion In the previously published paper17 we demonstrated that the dihydrogen complex 1aH2 is the key intermediate from which the H-H activation starts. Hence, it appears reasonable to start the calculations of the H-H bond splitting in these species from the dihydrogen complexes [Cp0 Ru(PTA)2(η2-H2)]þ (Cp0 = Cp, 1aH2; Cp0 = Cp*, 1bH2) surrounded by a cluster of four water molecules. 1. From the Dihydrogen to a Dihydrido Complex: The Oxidative Addition Process. Homolytic H-H splitting was found to be unfavorable compared with the heterolytic splitting in 1aH2.17 However, the reaction of H2 with 1b gives the dihydride, so it can be anticipated that the more basic properties of the permethylated ring may favor oxidative addition.4b,5a For these reasons, we first considered the homolytic pathway for 1b. The optimized structure of [Cp*Ru(PTA)2(η2-H2)]þ (1bH2) solvated by four water molecules is shown in Figure 1. The dihydrogen ligand has a calculated H-H distance of 0.845 A˚. It is coordinated

asymmetrically, with Ru-H bond lengths of 1.768 and 1.827 A˚, as a consequence of the extensive solvation by the water cluster. In contrast, in the free cation the η2-H2 ligand is symmetrically coordinated to the metal (both Ru-H distances are 1.771 A˚). The formation of a hydrogen bond between the H2 ligand and a water molecule is responsible for this asymmetry. The longer bond length corresponds to the H atom that lies closer to the water cluster [d(O-H) = 2.253 A˚], indicative of an interaction between the coordinated H2 and the solvent. Experimental evidence for the participation of coordinated H2 ligands in intermolecular hydrogen bonding with H-bond acceptors has been demonstrated recently.12a One of the water molecules in the cluster is hydrogen bonded to one PTA nitrogen atom [d(N-H) = 1.884 A˚]. The extra water molecule, with respect to the previous model,17 was placed next to the one that would accept the proton by assisting the deprotonation of the H2 ligand. H-H oxidative addition leads to a dihydride. In agreement with the lack of experimental detection of cis-[Cp0 Ru(PTA)2(H)2]þ and as found for the Cp-based system,17 a cisdihydride species was not located as a minimum for the Cp* case either. The only stable dihydride was found to be the trans-dihydrido complex 3b. In this structure the water chain bridging the N and H(H2) centers in 1bH2 is no longer

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present and the cluster is situated relatively far [d(O-H) = 2.989 A˚] from the hydride ligand. This is not surprising, if we consider the anionic character of the hydride ligand. The solvated trans-dihydride 3b is 1.6 kcal mol-1 more stable than its dihydrogen tautomer. Given the trans-hydride disposition in 3b, the dihydrogen f dihydride interconversion also entails an important stereochemical change, passing from a three-legged piano stool in 1bH2 to a transoid square-based four-legged piano stool in 3b. The transition state for the rearrangement [TSb(oxad)], depicted in Figure 1, has a dihydride nature (Ru-H bond distances of 1.594 and 1.600 A˚), with no interaction between the two hydrogen atoms (distance H-H of 2.488 A˚). The geometry around the ruthenium atom in this transition state is similar to that reported for the dihydrogen f trans-dihydride isomerization in related half-sandwich complexes.29 The calculated Gibbs energy for the activation leading to the formation of the dihydride through this pathway is 32.1 kcal mol-1. As a consequence of such a high barrier, the H-H homolytic splitting does not seem to be a feasible mechanism even for the more basic Cp* ligand. During the hydrogen migration, the water cluster undergoes a remarkable rearrangement. Thus, in the transition state the Ru-H 3 3 3 Owater hydrogen bond has completely disappeared, as shown by the long H 3 3 3 Owater separation of 3.193 A˚ (Figure 1). In agreement with this decreased interaction with the solvent in the transition state, the ΔGq in water is higher than that calculated in the gas phase (the latter being equal to 26.4 kcal mol-1). For comparative purposes we have recalculated the homolytic pathway for 1aH2 with an expanded cluster of four water molecules. The optimized geometries of all the species along this path are drawn in Figure S1. The Gibbs activation energy (31.2 kcal mol-1) is similar to that found for the Cp* complex and also similar to that calculated with a threewater-molecule cluster (30.3 kcal mol-1).17 Hence, we can state that the presence of an additional water molecule in the model does not affect the barrier of the oxidative addition significantly. Furthermore, these data suggest that the water cluster is not directly involved in this activation pathway. 2. From the Dihydrogen Complex to a Monohydride: The Heterolytic H-H Bond Cleavage. It has been proposed that an external base can accelerate the heterolytic H2 splitting by mediating a proton transfer between the coordinated H2 and the ancillary ligand.30 Recent reports stress the role of water molecules as efficient proton-transfer catalysts for the [1,2]H shift31a and for the assisted heterolytic splitting of H2.31b Starting from 1bH2, the H-H heterolytic splitting should give the monohydride [Cp*Ru(PTA)2H] (2b) and a solvated oxonium ion (reaction 1).

Instead of reaction 1, a concerted water-chain-assisted deprotonation of the H2 ligand and proton transfer to the PTA ligand led directly to the N-protonated monohydride (4a).17 With (H2O)4 we have characterized the process of reaction 1. The transition state [TSb(1f2)] corresponding to the deprotonation of complex 1bH2 is shown in Figure 2. It can be seen that the transferred proton lies between the oxygen of the adjacent water molecules and the hydride ligand [d(O-H) = 1.237 A˚ and d(H-H) = 1.018 A˚]. This transition state leads to the formation of 2b (lying 6.2 kcal/mol over 1bH2) interacting with a H3Oþ cation stabilized by the surrounding water molecules. In 2b the oxonium ion is pointing to the hydride ligand of the complex. The short O-H and H-H distances (1.074 and 1.219 A˚, respectively) indicate the presence of a strong dihydrogen bond. These dihydrogen-bonded structures have also been located in the (reverse) transition metal hydride protonation reaction.7,32 Deprotonation of the dihydrogen ligand by water is very easy, taking place with a Gibbs activation energy of only 5.9 kcal mol-1. In the previous study with an (H2O)3 cluster, we were not able to find structures such as 2b.17 The proton was transferred in a concerted mechanism from the coordinated η2-H2 molecule to a N atom of PTA through the chain of hydrogen bonds. To analyze whether the discrepancy is related either to the different nature of the complex (Cp vs Cp*) or to the description of the solvent [(H2O)3 vs (H2O)4 cluster], we have studied the deprotonation of 1aH2 assisted by (H2O)4. A minimum for 2a very similar to 2b has been located, consisting of the [CpRu(PTA)2H] monohydride interacting with a solvated oxonium ion (see Figure S2). As seen for the Cp* complex, the proton transfer takes place with a low barrier (6.0 kcal mol-1). The improvement of the solvent discrete modeling by extending the water cluster allows for stabilization of (H3O)þ. The neutral monohydride intermediate, in which one proton of the nonclassical ligand has “jumped” to the solvent, is located as a minimum. However, the main feature of the H2 deprotonation reaction remains unchanged: 2a and 2b are very shallow minima33 that evolve toward more stable geometries through protonation by a oxonium molecule from the solvent. 3. Protonation of the Monohydride. The proton transfer from the dihydrogen ligand to a solvent water molecule affords a neutral monohydride and a solvated oxonium cation. The proton in [(H3O)(H2O)3]þ can be transferred back to the basic centers of the ruthenium complex. The formation of a solvated H3Oþ as a product of the dihydrogen deprotonation leads to the consideration of two possible further pathways for the neutral complex: N-protonation at PTA (reaction 2) and metal protonation (reaction 3).

½CpRuðPTAÞ2 ðη2 -H2 Þþ 3 ðH2 OÞn

f ½CpRufPTAðHÞgðPTAÞHþ 3 ðH2 OÞn

f ½CpRuðPTAÞ2 H 3 ½H3 Oþ ðn - 1Þsolv

ð1Þ

However, in our previous study of 1aH2 deprotonation in the presence of (H2O)3 we were not able to locate such a species. (29) Baya, M.; Maresca, O.; Poli, R.; Coppel, Y.; Maseras, F.; Lled os, A.; Belkova, N. V.; Dub, P. A.; Epstein, L. M.; Shubina, E. S. Inorg. Chem. 2006, 45, 10248–10262. (30) Hutschka, F.; Dedieu, A. J. Chem. Soc., Dalton Trans. 1997, 1899–1902. (31) (a) Liang, Y.; Zhou, H.; Yu, Z.-X. J. Am. Chem. Soc. 2009, 131, 17783–17785. (b) Friederich, A.; Drees, M.; der G€unne, J. S. A.; Schneider, S. J. Am. Chem. Soc. 2009, 131, 17552–17553.

½CpRuðPTAÞ2 H 3 ½H3 Oþ solv

ð2Þ

½CpRuðPTAÞ2 H 3 ½H3 Oþ solv

f ½CpRuðPTAÞ2 ðHÞ2 þ 3 ðH2 OÞn

ð3Þ

Protonation at the PTA Ligand. The protonation of the PTA ligand by [(H3O)(H2O)3]þ was first analyzed. It was not (32) Dihydrogen Bonds: Principles, Experiments and Applications; Bakhmutov, V. I., Ed.; John Wiley & Sons, Inc.: Hoboken, NJ, 2008. (33) 2b is a local minimum only on the gas-phase potential energy surface. It is placed above TSb(1f2) when including entropic and solvent corrections.

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Figure 2. Relevant structures in the deprotonation of 1bH2 by water. H atoms in the ligands have been omitted for clarity. In parentheses relative Gibbs energies of the structures (kcal mol-1).

possible to obtain a minimum equivalent to 2b where the proton of the oxonium was pointing to the nitrogen of PTA. All the optimization attempts always ended up in structure 4b, with a proton transferred from the oxonium to one nitrogen atom of a PTA ligand (Figure 3). In our fourwater-molecule cluster, the proton is transferred without energy barrier to the PTA ligand, provided that it is placed close enough to one N atom of PTA. Thus, the barrier for the N-protonation from 2b is related only to the proton migration in water. Although this phenomenon cannot be properly tackled with our limited solvent model, it is well known to be a very fast process;34 hence it would not be the rate-determining step of the overall process. N-Protonation gives 4b as depicted in Figure 3 with a chain of three hydrogen-bonded water molecules connecting the hydride and the N-protonated hydrogen. The water chain starts with a dihydrogen bond at the hydride site and ends with a normal hydrogen bond between N-H of the quaternized PTA nitrogen and one oxygen atom. The fourth water molecule is solvating the water molecule involved in the stronger hydrogen bond. A species analogous to 4b was spectroscopically identified as a product of the reaction of 2a with one equivalent of HBF4,16b which was found to be the most stable species in the overall process of H-H bond splitting by 1aH2 with a cluster of three water molecules.17 The present study confirms the (34) (a) Agmon, A. Chem. Phys. Lett. 1995, 244, 456–462. (b) Day, T. J. F.; Schmitt, U. W.; Voth, G. A. J. Am. Chem. Soc. 2000, 122, 12027– 12028. (c) Lapid, H.; Agmon, N.; Petersen, M. K.; Voth, G. A. J. Chem. Phys. 2005, 122, 014506. (d) Han, J.; Zhou, X.; Liu, H. J. Power Sources 2006, 161, 1420–1427.

viability of this intermediate in the Cp*Ru-containing system and shows that it can also be obtained from the dihydrogen complex via a stepwise solvent-mediated proton transfer. The study of the N-protonation of 2a by [(H 3 O)(H2O)3]þ gives very similar results, leading to 4a without any energy barrier along the reaction trajectory (Figure S3). Remarkably, this product is much more stable in the Cp than in the Cp* system [ΔG(4a-1aH2) = -4.9 kcal mol-1; ΔG(4b-1bH2) = -1.4 kcal mol-1]. Protonation at the Metal. Proton delivery to the ruthenium center by H3Οþ would lead to a dihydrido complex. This protonation can occur only from the opposite side with respect to the hydride ligand. As in the case of the N-protonation, proton transfer can therefore occur only if the proton initially transferred from the dihydrogen to an adjacent water molecule jumps to a water molecule situated close to the metal center. As discussed above, proton mobility cannot be described properly with our model; however, it is well known to be a very fast, almost barrierless process.34 We have studied the protonation at the metal center optimizing one structure with the oxonium ion pointing toward the metal. Unlike the N-protonation, this structure is a local minimum (2b0 , Figure 3). Structure 2b0 is similar to 2b, differing only in the arrangement of the solvated [H 3O]þ: the proton to be transferred to the complex is located near the ruthenium center in 2b0 and near the hydride in 2b. The Ru-Hwater distance in 2b0 (2.283 A˚) agrees with the presence of a unconventional Hδþ 3 3 3 Ru hydrogen bond, where the metal acts as a

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Figure 3. Relevant structures in the protonation of 2b by an oxonium ion. H atoms in the Cp* and PTA ligands have been omitted for clarity. In parentheses relative Gibbs energies of the structures (kcal mol-1).

proton acceptor.35 Such structures have been found as initial intermediates in protonation processes at transition metal centers.7 Starting from 2b0 , the proton is transferred to the metal center through transition state TSb(2f3). In this early transition state, the proton lies in between ruthenium and oxygen, with the Ru-H distance at 2.076 A˚, whereas the O-H distance is 1.156 A˚. The Gibbs energy of activation for this proton transfer is very low (3.7 kcal mol-1). The product of metal protonation is the same as that formed by homolytic splitting, i.e., dihydride 3b discussed above. Protonation at the metal center of 2a by [(H3O)(H2O)3]þ follows a very similar pattern; the corresponding structures 2a0 and TSa(2f3) have been located (Figure S3).36 Interestingly, these data show that in the case of 2a metal protonation is even easier: TSa(2f3) is placed 3.4 kcal mol-1 above 2a, whereas TSb(2f3) lies 9.2 kcal mol-1 above 2b. 4. Mechanism of Hydrogen Activation in Water. In Figure 4 the complete Gibbs energy profile for the H-H activation in water by both 1aH2 and 1bH2 is shown. The ΔGwater of all the species [calculated as Gwater(x) - Gwater(1aH2) or (35) (a) Epstein, L. M.; Shubina, E. S. Coord. Chem. Rev. 2002, 231, 165–181. (b) Belkova, N. V.; Shubina, E. S.; Epstein, L. M. Acc. Chem. Res. 2005, 38, 624–631. (36) 2a0 is a local minimum only in the gas-phase potential energy surface. In the Gibbs energy profile in water it is placed above TSa(2f3).

Gwater(x) - Gwater(1bH2) for the generic Cp- or Cp*-containing compound x, respectively] are given in Table 1. The homolytic pathway has Gibbs activation energies of about 30 kcal mol-1 for both complexes. From this profile it can be inferred that neither for 1aH2 nor for 1bH2 does the H-H bond breaking proceed via a homolytic pathway. In both cases the oxidative addition has a much higher barrier than the deprotonation-protonation steps. Thus, the protonation mechanism ought to be of H-H heterolytic bond splitting type for both Cp and Cp* systems. The water solvent takes an active part, initially acting as an external base by deprotonating 1aH2 and 1bH2 to form 2a and 2b. This proton transfer generates an oxonium ion in the water medium. The very fast proton migration in water can efficiently relocate H3Οþ close to the basic centers of the molecule (either a nitrogen atom of PTA or the metal center), to allow further protonation. Further protonation of 2a and 2b either at the nitrogen or at the ruthenium center leads to the cationic species [Cp0 Ru(H)(PTA){PTA(H)}]þ (4) or [Cp0 Ru(H)2(PTA)2]þ (3), respectively (Scheme 2). All the barriers for deprotonation and protonation are low, albeit lower for the Cpcontaining complexes. In both systems, protonation at PTA is easier than that at ruthenium. The highest point in the ΔGwater profiles is found in the metal-protonation step. Although the barrier for the protonation at the metal is notably higher for 2b than for 2a, it is low enough to be

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Figure 4. Gibbs energy profiles for the H2 activation in water by 1aH2 and 1bH2. Table 1. Relative Energies and Gibbs Energies in Water (kcal mol-1) of All the Species Involved in the H2 Activation by Cp- and Cp*-Ruthenium-PTA Complexesa structure

ΔEwater CpRu system

ΔGwater CpRu system

ΔEwater Cp*Ru system

ΔGwater Cp*Ru system

1H2 TS(1f2) 2 4 20 TS(2f3) 3 TS(oxad)

0.0 4.8 0.5 -8.7 1.4 5.3 -3.3 28.4

0.0 6.0 2.1 -4.9 6.1 5.5 -2.3 31.2

0.0 4.4 1.9 -5.6 5.4 9.5 -3.9 29.4

0.0 5.9 6.2 -1.4 11.7 15.4 -1.6 32.1

a The dihydrogen species was taken as zero reference point for the evaluation of ΔEwater and ΔGwater.

overcome under the applied experimental conditions. Therefore, kinetic arguments cannot account for the evidence that, despite a common mechanism being at work in both the Cp and Cp* systems, different products are obtained by H2 activation (Scheme 1). The Gibbs energy profile shows that this behavior has a thermodynamic explanation. The relative stability of the dihydrogen, dihydride, and monohydridePTA(H) isomers is the key factor. 5. Ligand and Solvent Effects in the Stability of the Different Species. Three minima were found in the ΔGwater profile: the initial dihydrogen complex 1H2, the monohydridePTAH (4), and the dihydride (3) shown in Scheme 2. In the Cp system 4a is clearly the most stable product, placed 4.9 and 2.6 kcal mol-1 below 1aH2 and 3a, respectively. In the Cp* system 3b is the most stable product, being 1.6 kcal mol-1 more stable than 1bH2 but only 0.2 kcal mol-1 below 4b. In agreement with the experimental results, calculations predict that the H2 activation products should be 4a and 3b in the Cp and Cp* systems, respectively. Moreover, in the Cp* system it should be possible to transform 3b into 4b simply by

raising the temperature, which is the observed experimental behavior. We have further analyzed the origin of this behavior, with the aim of separating ligand and solvent effects. The relative stability of the three tautomers for the Cp and Cp* systems in media of different polarity is drawn in Figure 5. The intrinsic stability of the three species can be assessed with gas-phase calculations. Gas-phase optimization without considering the water cluster gives complex 3 as the most stable species for both systems. Thus, ligand effects favor the formation of a Ru(IV) dihydride such as 3, with PTA phosphines behaving as strong donors. The stability of 3 with respect to 1H2 is enhanced in the Cp* case, as could be expected from the higher electron-donating capability of the Cp* ligand compared with that of Cp. Moreover, complex 4 is considerably less stable than 1H2, although it is more stable for 4a than for 4b. The lower stability of 4b can be attributed to the lower acidity of the dihydrogen ligand, due to the presence of the more basic Cp* ligand. However, without solvent effects, H2 activation in both systems would lead to 3. To evaluate how the solvent can modulate this intrinsic stability, we have calculated the relative Gibbs energies of the three species on changing the solvent from n-heptane (ε = 1.92) to water (ε = 78.49, Figure 5). The ΔGsolvent[3-1] changes very slightly. The species more affected by the medium polarity is 4, which is considerably stabilized when the solvent polarity increases. In solvents with higher dielectric constant than acetone (ε = 20.71) 4a becomes more stable than 1aH2. Complex 4b undergoes a similar stabilization. However, as it was initially (gas phase, ε = 1) much above the dihydrogen than the equivalent Cp complex, it never becomes more stable than 1bH2. If water solvent is described only as a continuum polarizable medium (ε = 78.49), complex 3 is the most stable product for both Cp and Cp* complexes. The bigger stability change of 4 is found when specific interactions with water molecules are considered by means of a discrete þ continuum model. In such a case, the presence of a stronger hydrogen bond between

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Figure 5. Relative stability of the three tautomers for the Cp and Cp* systems in media of different polarity. The Gibbs energy of the dihydrogen complex in each medium is taken as zero reference point.

PTA(H)þ and one water molecule further stabilizes this structure. We have also calculated the Gibbs energy of the transition state of the oxidative addition in the same set of solvents. The ΔGq for the oxidative addition pathway steadily increases with the dielectric constant of the media, from 26.4 kcal mol-1 in the gas phase to 32.1 kcal mol-1 in water (cluster þ continuum model). ΔGq values of 27.8, 28.6, 29.2, and 29.2 kcal mol-1 are obtained in n-heptane, dichloromethane, DMSO, and water (only continuum), respectively. 6. pKa Calculations. The pKa values of coordinated dihydrogen have proved to be very useful to understand reactions that involve dihydrogen homolytic and heterolytic splitting.9a Considerable efforts have been made to create an acidity scale to better understand the acid/base behavior of cationic and neutral metal hydrides and dihydrogen complexes.37 Studies on the complexes [Cp0 RuL2(H2)]þ (R = H, Me; L = PR3) have revealed that dihydrogen acidity is extremely sensitive to the nature of L.9 The nature of the substituents in the cyclopentadienyl ring also affects the pKa, which increases of about two units when Cp is replaced by the more electron-donating Cp*. The estimated pKa of the dihydrogen complexes [Cp0 Ru(dppm)(η2-H2)]þ in THF are 7.4 and 9.2 for R = H and R = Me, respectively.9a,37 The corresponding values for the dihydrido complexes [(C5R5)Ru(PPh3)2(H)2]þ are 8.0 (Cp) and 10.9 (Cp*).37 The dihydrogen complexes are usually slightly more acidic than the corresponding dihydrido tautomers. pKa values of 7.0 and 7.4 have been estimated for the dihydrogen and dihydrido tautomers of [CpRu(dppe)(H2)]þ, respectively.37 The pKa of dihydrogen and dihydrido complexes has been measured experimentally in low-polar nonaqueous solvents, such as THF. To the best of our knowledge, pKa experimental data related to dihydrogen and dihydrido complexes in water are not available. In the framework of the cluster approach, the thermodynamic cycle depicted in Scheme 3 can be used to obtain a good estimate of the Gibbs energy of deprotonation in water (ΔG0sol,deprot) of both dihydrogen and dihydrido complexes. (37) Abdur-Rashid, K.; Fong, T. P.; Greaves, B.; Gusev, D. G.; Hinman, J. G.; Landau, S. E.; Lough, A. J.; Morris, R. H. J. Am. Chem. Soc. 2000, 122, 9155–9171.

From these values, the determination of pKawater is straightforward. In Scheme 3, ΔG0g,deprot represents the Gibbs energy of deprotonation in the gas phase and ΔG*solv(X) is the standard Gibbs energy of solvation of species X. The computed values are collected in Table 2. A value of -13.8 kcal mol-1 has been obtained for the aqueous solvation Gibbs energy of the (H2O)4 cluster [ΔG*solv(H2O)4], using the thermodynamic cycle for the formation of water clusters reported by Goddard et al.27a,38,39 This value is very close to the literature one (-14.1 kcal mol-1). The calculated aqueous solvation Gibbs energy of the Hþ(H2O)4 clustered ion (ΔG*solv{(H2O)4Hþ} = -66.1 kcal mol-1) is similar to that previously reported (-69.6 kcal mol-1).27a,40b ΔG0f* is the conversion factor from an ideal gas standard state of 1 atm to a standard state of 1 M (ΔG0f* = 1.9 kcal mol-1 at 298.15 K),40 and RT ln([H2O]/4) is the correction term (1.5 kcal mol-1) related with the Gibbs energy change of 1 mol of (H2O)4 from 55.34/4 M liquid state41 to 1 M. The Gibbs energy of deprotonation in water is computed from

ΔG0 soln, deprot ¼ ΔG0 g, deprot þ ΔGsolv ð½Ru-HÞ þ ΔGsolv ððH2 OÞ4 Hþ Þ - ΔGsolv ð½Ru-H2 þ Þ 

- ΔGsolv ðH2 O4 Þ - - ΔG0 f - RT ln½H2 O=4Þ ðeq 1Þ and then

pK a water ¼ ΔG0 soln, deprot =2:303RT

ðeq 2Þ

(38) The gas-phase Gibbs energy of complexation (ΔG0g,bind) calculated by Dunn et al. using the CBS-APNO level of theory was used to calculate ΔG*solv(H2O)4.39 For n = 4, ΔG0g,bind = 2.4 kcal mol-1 at T = 298.15 K. A value of -6.3 kcal mol-1 was employed for the Gibbs energy of self-solvation of water [ΔG*solv(H2O)] at T = 298.15 K.27a (39) Dunn, M. E.; Pokon, E. K.; Shields, G. C. J. Am. Chem. Soc. 2004, 126, 2647–2653. (40) (a) Kelly, C. P.; Cramer, C. J.; Truhlar, D. G. J. Chem. Theory Comput. 2005, 1, 1133–1152. (b) Kelly, C. P.; Cramer, C. J.; Truhlar, D. G. J. Phys. Chem. B 2006, 110, 16066–16081. (41) The concentration of H2O and (H2O)n in liquid water are 55.34 M and 55.34/n M, respectively.27a

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Scheme 3. Thermodynamic Cycle Used to Calculate pKawater Values

Table 2. Calculated Gibbs Energies of Deprotonation and Solvation Used to Estimate pKawater Following the Thermodynamic Cycle Shown in Scheme 3 complex

ΔG0g,deprot

ΔG*solv([Ru]-H2þ)

ΔG*solv([Ru]-H)

ΔG0sol,deprot

pKawater

1aH2 1bH2 3a 3b

20.2 25.4 21.7 30.3

-31.2 -17.0 -31.6 -17.7

2.8 14.0 2.8 14.0

-1.5 0.7 0.4 6.3

-1.1 0.5 0.3 4.7

The calculated values of pKawater are gathered in Table 2. A striking result is that the calculated pKa values in water are about seven pKa units lower than those estimated in THF from experimental values for related half-sandwich ruthenium complexes.9,37 The substantial pKa decrease has also been observed experimentally for cationic organic acids when moving from solvents such as acetonitrile and THF to water. For instance, the pKa of pyridinium is 12.5 in MeCN, 8.2 in THF, and 5.2 in water. A similar trend is found in PhNH3þ (10.6, 8.0 and 4.6) and Et3NHþ (18.8, 13.7, and 10.6).28d From the pKawater values in Table 2 it can be inferred that both dihydrogen complexes behave as strong acids in water; consequently, they are completely deprotonated. This unexpected result points out that the existence of cationic dihydrogen complexes in water solution could be doubtful due to the strong acidity of such species. The low acidity of 3b in comparison with that of all the other species is also remarkable. From the pKawater values in Table 2 both dihydrogen complexes and the Cp-dihydride (3a) are strong acids. Thus, when these species are formed in water, they should convert into the corresponding monohydrides. On the contrary, the Cp*-dihydride (3b) is a weak acid and should be only partially deprotonated.

Concluding Remarks Despite the fact that [CpRu(PTA)2Cl] and [Cp*Ru(PTA)2Cl] generate different products when they interact with H2 in aqueous solution, the computational study has shown that this reaction can proceed via a common mechanism. Even in the presence of the more basic Cp* ligand the H-H heterolytic splitting is favored over the oxidative addition. This result casts doubts on the feasibility of a homolytic H2 activation in water; caution must be used when commonly accepted reaction mechanisms are translated from nonpolar solvents to water. In the reaction of the water-soluble Vaska-type complexes trans-Ir(CO)(Cl)L2 (L = aryl or alkyl phosphine) with H2 it was reported that the rate constant increased by a factor of 10 (from 1.3 to 12) when changing the solvent from DMSO (ε = 46.7) to water.42a It is known that polar solvents accelerate the heterolytic activation of H2.43 The significant increase of the reaction rate in water might therefore be

related to a change in the reaction mechanism; a critical reinvestigation of H2 splitting by water-soluble Vaska’s42 and Wilkinson’s44 analogues is suggested. The solvent also plays a key role in the H2 activation performed by [Cp0 Ru(PTA)2Cl]. The overall process consists of a sequence of fast deprotonation-protonation steps: deprotonation of the H2 ligand by the solvent, generating a oxonium ion, followed by protonation of the basic centers of the complex (either a nitrogen atom of PTA or ruthenium). The relative stability of both intermediates and products is governed by their interaction with the solvent. Thus, the monohydride-PTA(H)þ product 4a is strongly stabilized by a polar medium. However, if these explicit interactions are not taken into account, this product is less stable than the corresponding dihydride. The existence of a strong hydrogen bond between PTA(H)þ and the solvent makes [Cp0 Ru(PTA){PTA(H)}(H)]þ (4) the most stable product for the Cp system and as stable as the dihydride [Cp*Ru(PTA)2(H)2]þ for the Cp* complex. A theoretical estimation of the pKa in water of the dihydrogen and dihydride species suggests that these compounds can be much more acidic in water than in organic solvents such as acetonitrile or tetrahydrofuran. From the calculated pKawater values, it arises that both 1aH2 and 1bH2 behave as strong acids in water, and thus they should be completely deprotonated. The same behavior is expected for 3a, while the decreased acidity of 3b makes it stable enough to survive in aqueous solution. In conclusion, this work provides further compelling evidence that the subtle “tuning” of the reaction course takes place in aqueous media, showing the active role of water solvent in deciding the reaction mechanisms and the energetically most favorable products. Further experimental and theoretical studies are nonetheless required to improve the “degree of control” of the resulting H2 activation mechanism in water. (42) (a) Paterniti, D. P.; Roman, P. J.; Atwood, J. D. Organometallics 1997, 16, 3371–3376. (b) Paterniti, D. P.; Atwood, J. D. Polyhedron 1998, 17, 1177–1181. (43) Brothers, P. J. Prog. Inorg. Chem. 1981, 28, 1–61. (44) Jo o, F.; Kovacs, J.; Benyei, A.; Nadasdi, L.; Laurenczy, G. Chem.;Eur. J. 2001, 7, 193–199.

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Acknowledgment. We gratefully acknowledge the European Commission for funding this work through the AQUACHEM Research Training Network (contract no. MRTN-CT-2003-503864). Support from the Spanish MICINN (Projects CTQ2008-06866-CO2-01 and Consolider-Ingenio 2010 CSD2007-00006 and Juan de la Cierva contract to G.K.) is also acknowledged. Ente Cassa di Risparmio di Firenze (ECRF) is gratefully acknowledged for funding this research through the

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project Firenze Hydrolab (http://www.iccom.cnr.it/ hydrolab) and for sponsoring a postdoctoral grant to A.R. Supporting Information Available: Optimized geometries of intermediates and transition states for the reaction of the Cp complex (1aH2). Cartesian coordinates and absolute Gibbs energies of all the calculated species. This material is available free of charge via the Internet at http://pubs.acs.org.