Carbonate Formation from CO2 via Oxo versus Oxalate Pathway

Oct 14, 2010 - (h) Denning , R. G.; Green , J. C.; Hutchings , T. E.; Dallera , C.; Tagliaferri , A.; Giarda , K.; Brookes , N. B.; Braicovich , L. J...
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Organometallics 2010, 29, 5504–5510 DOI: 10.1021/om100479r

Carbonate Formation from CO2 via Oxo versus Oxalate Pathway: Theoretical Investigations into the Mechanism of Uranium-Mediated Carbonate Formation† Ludovic Castro,§ Oanh P. Lam,‡ Suzanne C. Bart,‡ Karsten Meyer,*,‡ and Laurent Maron*,§ ‡

Department of Chemistry and Pharmacy, Inorganic Chemistry, University Erlangen-Nuremberg, Egerlandstrasse 1, D-91058 Erlangen, Germany, and §Universit e de Toulouse, INSA, UPS, LPCNO, 135 Avenue de Rangueil, F-31077 Toulouse, France, and CNRS, LPCNO, F-31077 Toulouse, France Received May 17, 2010

We report theoretical investigations of the reaction of [((MeArO)3mes)U] with CO2 in order to support previously reported experimental data. Experimentally, the reaction in toluene leads to the immediate formation of the bridging carbonate complex [{((MeArO)3mes)UIV}2(μ-η2:η2-CO3)] at room temperature. DFT calculations show that the preferred reaction pathway is a three-step mechanism: first, the formation of a dinuclear CO2 complex, followed by concomitant release of CO, forming the corresponding bridging μ-oxo species. The final step involves insertion of a CO2 molecule into a U-O bond, forming the carbonate product. Calculations reveal this three-step process to be thermodynamically favorable and kinetically accessible. An alternate pathway that proceeds through an oxalate dinuclear complex is also explored. Although the oxalate complex is calculated to be the thermodynamic product of the reaction, a high activation barrier prevents its formation.

Introduction While carbon dioxide has been implicated as a greenhouse gas, it also has the potential to play a positive role in industry, as its low cost and abundance make it an ideal chemical feedstock. A challenge often encountered in CO2 transformation is the activation of thermodynamically strong carbonoxygen double bonds. By identifying appropriate ligand environments and electron-rich metal centers, reduction of CO2 is possible, yielding various products such as oxos, oxalates, and carbonates.1 Uranium complexes are ideal for carbon dioxide activation because of the oxophilicity, Lewis acidity, and highly reducing nature of low-valent U(III) congeners. In addition, the highly polarized bonds in mid-valent U(IV) amido and high-valent U(V) imido complexes offer a unique platform to explore the activation of strong bonds. Previous results from our laboratory utilizing tris(aryloxide)-derivatized † Part of the Dietmar Seyferth Festschrift. *To whom correspondence should be addressed. E-mail: karsten. [email protected]; [email protected] (1) (a) Castro-Rodriguez, I.; Meyer, K. J. Am. Chem. Soc. 2005, 127, 11242–11243. (b) Laitar, D. S.; Muller, P.; Sadighi, J. P. J. Am. Chem. Soc. 2005, 127, 17196–17197. (c) Lee, G. R.; Maher, J. M.; Cooper, N. J. J. Am. Chem. Soc. 1987, 109, 2956–2962. (d) Evans, W. J.; Seibel, C. A.; Ziller, J. W. Inorg. Chem. 1998, 37, 770–776. (e) Davies, N. W.; Frey, A. S. P.; Gardiner, M. G.; Wang, J. Chem. Commun. 2006, 4853–4855. (f ) Summerscales, O. T.; Frey, A. S. P.; Cloke, F. G. N.; Hitchcock, P. B. Chem. Commun. 2009, 198–200. (2) (a) Castro-Rodriguez, I.; Olsen, K.; Gantzel, P.; Meyer, K. Chem. Commun. 2002, 2764–2765. (b) Nakai, H.; Castro-Rodriguez, I.; Zhakarov, N. L.; Rheingold, A. L.; Meyer, K. Inorg. Chem. 2004, 43, 855–857. (c) Lam, O. P.; Heinemann, F. W.; Meyer, K. C. R. Chim. 2010, 13, 803–811.

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triazacyclononane-based chelating ligands with variable substituents (1,4,7-tris(3-R-5-tert-butyl-2-hydroxybenzyl)1,4,7-triazacyclononane, R = tert-butyl, adamantyl, diamantyl)2 have demonstrated the influence of the chelator’s steric environment and the uranium center’s oxidation state on product formation from reactions with carbon dioxide and other small molecules. Addition of carbon dioxide to the low-valent U(III) complex [((tBuArO)3tacn)U] immediately results in the formation of the bridging oxo species [{((tBuArO)3tacn)U}2(μ-O)] with extrusion of carbon monoxide.3 This product is formed via a two-electron reduction of carbon dioxide by two trivalent uranium centers, generating a chemically inert, linear μ-oxo species bridging two tetravalent uranium centers. Employing the sterically more encumbering adamantyl-derivatized complex effectively prevents formation of dinuclear species, resulting in the production of a charge-separated uranium carbon dioxide species, [((AdArO)3tacn)UIV(CO2•-)], which features a one-electron-reduced CO2 ligand with an end-on oxygen-bound coordination mode.4 The same macrocyclic ligands support the mid-valent uranium(IV) amido complexes [((RArO)3tacn)U(NHMes)] (Mes= 2,4,6-trimethylphenyl, R = t-Bu, Ad), which readily insert carbon dioxide into the U-N bond, forming the corresponding carbamate complexes.5 In contrast, high-valent uranium(V) complexes with bent imido ligands, e.g., [((RArO3)tacn)U(NMes)], (3) Castro-Rodriguez, I.; Meyer, K. J. Am. Chem. Soc. 2005, 127, 11242–11243. (4) Castro-Rodriguez, I.; Nakai, H.; Zakharov, L. N.; Rheingold, A. L.; Meyer, K. Science 2004, 305, 1757–1760. (5) Bart, S. C.; Anthon, C.; Heinemann, F. W.; Bill, E.; Edelstein, N. M.; Meyer, K. J. Am. Chem. Soc. 2008, 130, 12536–12546. r 2010 American Chemical Society

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Figure 1. Molecular structure cores of [{((tBuArO)3mes)U}2(μ-η2:η2-CO3)] (left) and [{((AdArO)3N)U}2(μ-η1:η2-CO3)] (right) displaying two different binding modes of CO32-.

undergo multiple-bond metathesis upon exposure to carbon dioxide to generate uranium(V) terminal oxo species, [((RArO)3tacn)U(O)], and the corresponding free isocyanate, MesNCO.5 Replacing the triazacyclononane anchor with either a mesitylene ring or a single nitrogen atom has a drastic effect on the reactivity of the electron-rich uranium(III) species with carbon dioxide. Exposure of [((tBuArO)3mes)U] and [((AdArO)3N)U] to one atmosphere of carbon dioxide results in the formation of the bridging uranium carbonate species [{((tBuArO)3mes)U}2(μ-η2:η2-CO3)] and [{((AdArO)3N)U}2(μ-η1:η2-CO3)] (Figure 1).6 Similar reactivity was observed by Cloke et al., who reported formation of carbonate by a U(III) metallocene upon exposure to CO2, producing [{(η8-C8H6{SiiPr3-1,4}2)(η5-C5Me4H)U}2(μ-η1:η2-CO3)].7 Gardiner et al. have demonstrated similar carbon dioxide reduction with samarium.8a The authors report the formation of a porphyrinogen-based dinuclear carbonate-bridged Sm(III)/Sm(III) complex, for which the authors propose an unstable and unobservable metal oxalate intermediate. Gardiner’s proposal is reasonable, considering Evans’ stabilization and isolation of samarium oxalate complex [{(C5Me5)2Sm}2(μ-η2:η2-O2CCO2)] in high yield from reductive coupling of carbon dioxide by the divalent samarocene [(C5Me5)2Sm(THF)2].8b Although a samarium carbon dioxide radical anion is proposed as the intermediate species in solution, no evidence for the formation of carbonate species is reported on the pathway to the oxalate species or as a product from the decomposition of [{(C5Me5)2Sm}2(μ-η2:η2-O2CCO2)]. Herein, we present a comprehensive computational analysis elucidating the mechanism of carbonate formation by the trivalent uranium coordination complex [((tBuArO)3mes)U]. Using the binding modes presented in Figure 1 as (6) Lam, O. P.; Bart, S. C.; Kameo, H.; Heinemann, F. W.; Meyer, K. Chem. Commun. 2010, 46, 3137–3139. (7) Summerscales, O. T.; Frey, A. S. P.; Cloke, F. G. N.; Hitchcock, P. B. Chem. Commun. 2009, 198–200. (8) (a) Davies, N. W.; Frey, A. S. P.; Gardiner, M. G.; Wang, J. Chem. Commun. 2006, 4853–4855. (b) Evans, W. J.; Seibel, C. A.; Ziller, J. W. Inorg. Chem. 1998, 37 (4), 770–776. (9) (a) Vallet, V.; Schimmelpfennig, B.; Maron, L.; Teichteil, C.; Leininger, T.; Gropen, O.; Grenthe, I.; Wahlgren, U. Chem. Phys. 1999, 244, 185–193. (b) Hayton, T. W.; Boncella, J. M.; Scott, B. L.; Palmer, P. D.; Batista, E. R.; Hay, P. J. Science 2005, 310, 1941–1943. (c) Iche-Tarrat, N.; Barros, N.; Marsden, C. J.; Maron, L. Chem.;Eur. J. 2008, 14, 2093–2099. (d) Ismail, N.; Heully, J. L.; Saue, T.; Daudey, J. P.; Marsden, C. J. Chem. Phys. Lett. 1999, 300, 296–302. (e) Pepper, M.; Bursten, B. E. Chem. Rev. 1991, 91, 719–741. (f ) Li, J.; Bursten, B. E.; Liang, B. Y.; Andrews, L. Science 2002, 295, 2242–2245. (g) Gagliardi, L.; Roos, B. O. Nature 2005, 433, 848–851. (h) Denning, R. G.; Green, J. C.; Hutchings, T. E.; Dallera, C.; Tagliaferri, A.; Giarda, K.; Brookes, N. B.; Braicovich, L. J. Chem. Phys. 2002, 117, 8008–8020. (i) Groenwold, G. S.; Gianotto, A. K.; Cossel, K. C.; Van Stipdonk, M. J.; Moore, D. T.; Polfer, N.; Oomens, J.; De Jong, W. A.; Vissher, L. J. Am. Chem. Soc. 2006, 128, 4802–4813. ( j) Yahia, A.; Arnold, P.; Love, J. B.; Maron, L. Chem. Commun. 2009, 17, 2402–2404. (k) O' Grady, E.; Kaltsoyannis, N. Dalton Trans. 2002, 6, 1233–1239.

a structural basis, we have performed computational analyses to determine the potential mechanisms for formation of [{((tBuArO)3mes)U}2(μ-η2:η2-CO3)]. We utilized both RECPs since the size of the large systems involved does not allow the use of a very small core ECP. For these studies, the 5f-in-core ECP is more appropriate, while the very small core pseudopotential was used for the calculation of oxidation reactions.

Computational Details Computational studies are well established and can effectively describe the structure and reactivity of uranium-containing molecules.9 However, both relativistic effects (solved by using relativistic effective core potentials, RECPs) and electronic correlation effects (mainly due to the presence of 5f unpaired electrons in the ground state) must be considered, which increase the complexity of the calculations.9a,10 Despite the fact that density functional theory (DFT) leads to excellent representations of geometries and vibrational frequencies, electronic correlation effects previously did not allow the use of this method for these systems.9a,10b,11 More recently, studies have shown that DFT methods in combination with very small core RECPs (explicit treatment of the 5f electrons) can be applied to investigate the reactivity of uranium coordination complexes.12 Moreover, new f-in-core relativistic effective core potentials (ECPs) derived by Moritz et al. have overcome the problem of dealing with open-shell systems.13 Moritz et al. have shown that these ECPs lead to reasonable geometrical parameters and dissociation energies. In recent studies, Castro et al. have generalized these results to reaction energies and activation barriers of reactions involving organouranium complexes.12b,14 Uranium atoms were treated with two different effective core potentials. First, the 5f-in-core ECP adapted to the uranium þ4 oxidation state was used in combination with its adapted basis set for uranium atoms in order to account for reactions involving dinuclear complexes.13 Second, the very small core ECP was used in combination with its adapted basis set to study mononuclear complexes and changes of oxidation state.15 Carbon, oxygen, and hydrogen atoms have been described with a 6-31G(d,p) double-ζ basis set.16 Calculations were carried out at the (10) (a) Pyykk€ o, P. Inorg. Chim. Acta 1987, 139, 243–245. (b) Vallet, V.; Maron, L.; Schimmelfennig, B.; Leininger, T.; Teichteil, C.; Gropen, O.; Grenthe, I.; Wahlgren, U. J. Phys. Chem. A 1999, 103, 9285–9289. (c) Dolg, M.; Fulde, P.; Stoll, H.; Preuss, H.; Pitzer, R. M.; Chang, A. Chem. Phys. 1995, 195, 71–82. (11) Clavaguera-Sarrio, C.; Hoyau, S.; Ismail, N.; Marsden, C. J. J. Chem. Phys. A 2003, 107, 4515–4525. (12) (a) Barros, N.; Maynau, D.; Maron, L.; Eisenstein, O.; Zi, G. F.; Andersen, R. A. Organometallics 2007, 26, 5059–5065. (b) Castro, L.; Yahia, A.; Maron, L. ChemPhysChem 2010, 11, 990–994. (13) Moritz, A.; Cao, X.; Dolg, M. Theor. Chem. Acc. 2007, 118, 845– 854. (14) Castro, L.; Yahia, A.; Maron, L. Dalton Trans. 2010, 39, 6682. (15) (a) Kuchle, W.; Dolg, M.; Stoll, H.; Preuss, H. J. Chem. Phys. 1994, 100, 7535–7543. (b) Cao, X. Y.; Dolg, M.; Stoll, H. J. Chem. Phys. 2003, 118, 487–497. (16) Hehre, W. J.; Ditchfield, R.; Pople, J. A. J. Chem. Phys. 1972, 56, 2257–2261.

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Figure 2. Experimental (left) and modeled (right) chelating ligands. DFT level of theory using the hybrid functional B3PW91.17 Geometry optimizations were performed without any symmetry restrictions, and the nature of the extrema (minima and transition states) was verified with analytical frequency calculations. Gibbs free energies were obtained at T = 298.15 K within the harmonic approximation. IRC calculations were performed to confirm the connections of the optimized transition states. DFT calculations were carried out with the Gaussian03 suite program.18 The electronic density (at the DFT level) has been analyzed using the natural bond orbital (NBO) technique.19 Calculations have been realized in the gas phase since the experimental solvent is toluene, which is a nonpolar and nonprotic solvent and can thus be neglected. Hydrogen atoms have been omitted for clarity in all the 3D structures presented in this paper, and the (MeArO)3mes3- ligand is denoted as L in the schematic representations.

Results and Discussions The experimental tripodal ligand 1,3,5-trimethyl-2,4,6tris((2,4-di-tert-butylphenol)methyl)benzene, (tBuArOH)3mes, was modeled by the sterically less encumbering methyl derivative 1,3,5-trimethyl-2,4,6-tris((2,4-dimethylphenol)methyl)benzene, (MeArOH)3mes (Figure 2), in order to simplify the calculations. Validation of the DFT method was carried out by comparing the molecular structures of the experimental uranium(III) reactant,20 [((tBuArO)3mes)U], with the calculated [((MeArO)3mes)U], using a very small core pseudopotential (Figure 3). The optimized calculated structure for [((MeArO)3mes)U] is a quasi-C3-symmetrical complex with the uranium center coordinated in a distorted trigonal fashion to the three aryloxide oxygen atoms. The average U-O bond distance of 2.156 A˚ is close to the experimentally determined bond distance of 2.164 A˚. The aryloxide rings are positioned perpendicular to both the arene and the plane formed by the three oxygen atoms. The uranium ion is situated 0.444 A˚ below this plane, pointing toward the mesitylene anchor (0.464 A˚ in the experimental structure). The average C-C bond distance within the mesitylene is computed to be 1.422 A˚ (1.419 A˚ in the experimental structure). The planar nature (17) (a) Becke, A. D. J. Chem. Phys. 1993, 98, 5648–5652. (b) Burke, K.; Perdew, J. P.; Yang, W. Electronic Density Functional Theory: Recent Progress and New Directions; Wiley & Sons: New York, 1998. (18) See SI for complete Gaussian03 reference. (19) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Chem. Rev. 1988, 88, 899–926. (20) Bart, S. C.; Heinemann, F. W.; Anton, C.; Hauser, C.; Meyer, K. Inorg. Chem. 2009, 48, 9419–9426.

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of the mesitylene ring allows for an η6 interaction with the uranium, resulting in an average uranium-carbon distance of 2.72 A˚ and a uranium-arenecentroid distance of 2.33 A˚, both in excellent agreement with the experimentally observed distances (2.73 and 2.33 A˚, respectively). It is noteworthy that NBO analysis shows the presence of three single unpaired alpha electrons in the 5f orbitals of uranium, confirming the experimentally determined þ3 oxidation state of this U ion. Experimentally, the reaction of the purple [((tBuArO)3mes)UIII] (1) with CO2 (1 atm) in toluene at room temperature leads to the formation of the light yellow carbonate bridging complex [{((tBuArO)3mes)UIV}2(μ-η2:η2-CO3)].6 Since the fast rate of reaction precludes the isolation of any potential intermediates, the mechanism of this carbonate product formation was probed using DFT calculations. Gaining insight into the mechanism is essential to developing new stoichiometric and potential catalytic cycles for carbonate formation mediated by uranium centers. Based on the arene-anchored tripodal uranium(III) model complex, [((MeArO)3mes)U], three possible pathways were investigated for the reaction with CO2: a concerted mononuclear pathway and an “oxo pathway”, both leading to the experimentally observed carbonate product, and the oxalate formation pathway, leading to an oxalate product. 1. Concerted Mononuclear Pathway. The first pathway (Figure 4) involves the preliminary formation of two different mononuclear U(IV) charge-separated complexes [((MeArO)3mes)UIV(CO2•-)], 2 and 3, with radical anionic CO2 ligands, followed by subsequent coupling of these products to form the bridging dinuclear species 4. Simplified 3D views of the optimized structures are provided in Figure 4 (full geometries are available in the Supporting Information (SI)). Both species 2 and 3 feature carbon dioxide bound in respective η2-OCO and η2-OCO coordination modes. The respective C-O1 and C-O2 (see Figure 5 for atom numbering) bond distances for 2 are nearly equivalent, measuring 1.26 and 1.20 A˚, both elongated from those in free CO2 (1.17 A˚). The O-C-O angle in 2, measuring 143°, possesses significant deviation from linearity. For 3, the C-O bond distance of 1.25 A˚ and the O-C-O angle of 127° are close to the values calculated for a CO2•- radical anion (1.25 A˚ and 134°). Moreover, NBO analysis reveals one unpaired electron lying on an sp orbital of the CO2 carbon for both 2 and 3, confirming the radical nature of the CO2•ligand. Furthermore, two unpaired electrons are observed on the uranium 5f orbitals, consistent with a þ4 oxidation state. These observations support that the coordination of CO2 induces a one-electron oxidation of the trivalent uranium center with concomitant reductive activation of CO2. The formation of 2 and 3 is endergonic, with values of 13.0 and 13.2 kcal/mol, respectively; this results in simultaneous formation of both isomers (the corresponding energy on the energy profile is the largest energy necessary to form both isomers). Subsequently, 2 and 3 can interact with each other through the sterically accessible oxygen atom of 2 and the carbon atom of 3, generating a new bridging complex, 4. The driving force for the formation of 4 is the high energy gain from stabilization of the electron-deficient uranium(IV) centers, making this dimerization exergonic by 27.4 kcal/mol. Even though this pathway would lead to the final bridging carbonate complex 6 (which will be discussed later) via the intermediate 4 with an accessible relative activation barrier of 22.3 kcal/mol, it is not favored compared to the

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Figure 3. Molecular structure of [(tBuArO)3mes)U] shown with 30% probability ellipsoids (top) and calculated structure (bottom). Hydrogen atoms and solvent molecules have been omitted for clarity.

Figure 4. Calculated Gibbs free energy profile of the concerted mononuclear pathway. For clarity, only the three oxygens of the macrocyclic ligands are shown. Full geometries are available in the SI.

alternative oxo pathway. Thus, the transition state 5 is not discussed here. 2. Oxo Pathway. Carbonate formation has been experimentally verified to occur via a bridging oxo intermediate6 and, hence, was explored as a viable pathway. Previously

reported reactivity studies of trivalent uranium species of the tacn system with CO2 have led to the formation of an unreactive bridging oxo complex, in which the μ-oxo ligand is sterically well protected by two, interlinked tert-butylderivatized tris(aryloxide) tacn ligands.2a,3 However, due to

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Figure 5. Schematic representation of intermediates 2, 3, and 4.

Figure 6. Calculated Gibbs free energy profile of the oxo pathway. For clarity, only the three oxygens of the macrocyclic ligands are shown in the 3D structures. Full geometries are available in the SI.

an increased flexibility in the mesityl-anchored ligand, a more exposed and, thus, reactive bridging oxo ligand can be supported.20 Therefore, an alternative pathway (Figure 6) via a reactive bridging oxo intermediate was evaluated. This pathway involves preliminary formation of the bridging CO2 species 7, followed by subsequent release of CO to form the intermediate bridging oxo [{((MeArO)3mes)U}2(μ-O)] (9) through transition state 8. Finally, in a low-energy pathway, a free CO2 molecule can insert into 9, thereby forming the bridging carbonate complex 6 through transition state 10. The Gibbs free energy value for the formation of complex 7 is lower by 12.1 kcal/mol than the value for the formation of 4. This is expected since the formation of complex 4 involves two CO2 molecules, whereas the formation of 7 requires only one CO2,21 thus making formation of 7 exergonic by 26.3 kcal/mol with respect to the reactants. The energetic stabilization gained by formation of 7 is on the same order as that of 4 with respect to 2 and 3. The C-O1 and C-O2 (see Figure 7 for atom numbering) bond lengths of 7 measure 1.33 and 1.27 A˚, respectively, and the O-C-O angle in the carbon dioxide ligand is 116°, consistent with a doubly reduced carbon dioxide ligand in a singlet spin state, for which the typical C-O bond distances are 1.34 A˚ and the O-C-O angle is 115°. Thus, complex 7 can be described as a CO22- ligand stabilized by two [(RArO)3mes)U(IV)] cationic fragments. Transition state 8 is an intermediate of the bridging CO2 complex 7 and the bridging oxo complex 9, formed through the release of CO. The corresponding relative activation barrier is equal to 12.5 kcal/mol, characteristic of a kinetically accessible reaction. The C-O1 bond of 8 (see Figure 7 for atom numbering) is elongated and almost cleaved with respect to 7 (1.61 vs 1.33 A˚). The C-O2 bond is shortened (21) Watson, L.; Eisenstein, O. J. Chem. Educ. 2002, 79, 1269–1277.

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Figure 7. Schematic representation of intermediate 7 (left) and transition state 8 (right).

(1.19 vs 1.27 A˚) and is close to that of free CO (1.14 A˚). The O-C-O angle does not vary (116°), whereas the U-O-U angle opens from 164° to 206° in order to develop an sp orbital of the oxygen on CO. Moreover, the uranium atom on the right (see Figure 6) lies significantly below the plane formed by the other uranium and the CO2 moiety (0.60 A˚). From a second-order NBO analysis, it is noteworthy that the oxygen atom of the released CO is back-donating (69 kcal/mol) into the antibonding orbital of the C-O1 bond. Finally, the CO molecule is released and the intermediate bridging oxo complex 9 is formed. Its formation is exergonic by 12.8 kcal/mol with respect to the bridging CO2 complex 7 and, thus, is thermodynamically favorable. Additionally, the bent U-O-U moiety in 9 exposes the orbitals of the oxygen atom and, thus, allows further reactivity such as the insertion of a free CO2 molecule to form the bridging carbonate complex 6. The insertion of CO2 has also been experimentally demonstrated for the structurally characterized μ-oxo complex in a related system, [{(AdArO)3N)U}2(μ-O)], which also features a bent U-O-U moiety.6 The relative activation barrier to transition state 10 in the insertion of CO2 into complex 9 is 20.6 kcal/mol, characteristic of a kinetically accessible reaction. The bending of both the U-O-U and O-C-O moieties allows a lone pair on the oxo ligand to point toward an empty sp2 orbital of the carbon of the CO2 molecule. An oxygen atom of the free CO2 interacts with uranium to form the U-O bond present in the carbonate bridging complex (2.60 A˚), while the corresponding C-O bond is slightly activated (1.20 vs 1.17 A˚ in free CO2); the other C-O bond is unaltered. The experimentally and computationally optimized structures of the bridging carbonate complex 6 are presented in Figure 8. The calculations indicate its formation is highly exergonic (-54.9 kcal/mol with respect to the reactants) and, thus, very favorable. The two uranium centers lie below and above the plane of the carbonate unit, with displacements of U1 at 0.529 A˚ above the plane and U2 at 0.445 A˚ below the plane (experimentally 0.476 and 0.421 A˚, respectively) (Figure 9). The U-CO3-U core features two short U-O bonds (2.366 and 2.365 A˚, calcd; 2.333(4) and 2.332(3) A˚, exptl) and two long U-O bonds (2.556 and 2.547 A˚, calcd; 2.659(4) and 2.603(4) A˚, exptl). The carbonate unit exhibits two shorter C-O distances (1.277 and 1.276 A˚, calcd; 1.285(6) and 1.279(7) A˚, exptl) and one longer C-O distance of 1.321 A˚ (1.305(6) A˚ experimental). The difference between the computationally optimized and experimentally determined structural values may be due to the truncated model that has been applied for the computational study. 3. Oxalate Formation Pathway: A Viable Alternative. Finally, an alternative pathway (Figure 10) has been evaluated, which leads to a bridging oxalate intermediate, [{((MeArO)3mes)U}2(μ-η2:η2-C2O4)] (12), from the preliminary bridging CO2 species 7 and a free CO2 molecule through transition state 11. This reaction is of particular interest, as Gardiner states it may be a viable reaction pathway in the formation of a Sm(III)/Sm(III) bridging carbonate complex and could be for uranium as well.8

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Figure 8. Molecular structure of [{((tBuArO)3mes)UIV}2(μ-η2:η2-CO3)] shown with 30% probability ellipsoids. Hydrogen atoms and selected solvent molecules have been omitted for clarity (top). Optimized structural representation of the carbonate bridging complex 6 (bottom).

Figure 9. Core structure of calculated 6 depicting the displacements of the two uranium centers above and below the plane of the CO32- unit (black).

The formation of bridging oxalate 12 (Figure 10) is highly exergonic (70.4 kcal/mol with respect to the reactants) and is more favorable than that of the bridging carbonate complex 6. The structure of 12 corresponds to a planar side-on coordinated C2O42- ion bridged between two uranium centers. The transition-state complex 11 features a μ-CO2 complex with a free CO2 molecule approaching the U-CO2-U plane. Compared to complex 7, where the CO2 is coordinated κ2:η2 to the two uranium centers, 11 has an η1:η2 coordination mode, prearranging one uranium center for the coordination of the incoming CO2, which then undergoes activation, causing C-O bond lengthening to 1.17 A˚ and slight bending to 173°. The relative activation barrier of this reaction (through transition state 11) is 22.6 kcal/mol, which is 10 kcal/mol higher than for the formation of the bridging

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Figure 10. Calculated Gibbs free energy profile of the oxalate formation pathway. For clarity, only the three oxygens of the macrocyclic ligands are represented in the 3D. Full geometries are available in the SI.

oxo complex 9. Thus, this pathway is less kinetically favored than the carbonate formation one, allowing observation of the bridging carbonate complex 6 even if formation of the bridging oxalate complex 12 is thermodynamically favored. Since the formation of the kinetic product (carbonate complex 6) is irreversible, the thermodynamically favored product (oxalate complex 12) cannot be formed.

Conclusion The theoretical studies presented here propose a mechanism for the formation of [{((MeArO)3mes)U}2(μ-η2:η2-CO3)] from the reaction of [(MeArO)3mes)U] with CO2. The threestep “oxo pathway” is thermodynamically favorable and kinetically accessible, proceeding through a bent bridging oxo U(IV/IV) intermediate. The μ-oxo complex allows for CO2 insertion to form the carbonate complex due to increased flexibility of the mesityl-anchored ligand system, which overall decreases the steric protection of the complex. Perhaps by further altering the steric protection of the

complex through ligand design, various degrees of reactivity for the μ-oxo moiety can be discovered. The dinuclear oxalate intermediate 12 is calculated to be the thermodynamic product but cannot be formed due to a high activation barrier. Variations of the ligand system may bring the activation energy down and favor the formation of an oxalate instead of the carbonate complex.

Acknowledgment. The authors acknowledge the University of Erlangen-N€ urnberg, the Deutsche Forschungsgemeinschaft (DFG) through the Sonderforschungsbereich SFB 583, and ME1754/2-1 (K.M.), the Institut Universitaire de France, CNRS, and UPS for generous financial support, and CALMIP (CNRS, Toulouse, France) and CINES (CNRS, Montpellier, France) for calculation facilities (L.M.). S.C.B. thanks the Alexander von Humboldt Foundation for a postdoctoral fellowship. Supporting Information Available: Coordinates of all stationary point structures in xyz format. This material is available free of charge via the Internet at http://pubs.acs.org.