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Selectivity in the C−H Activation Reaction of CH3OSO2CH3 with [1,2,4-(Me3C)3C5H2]2CeH or [1,2,4-(Me3C)3C5H2][1,2-(Me3C)2-4(Me2CCH2)C5H2]Ce: To Choose or Not To Choose Evan L. Werkema,† Ludovic Castro,‡ Laurent Maron,*,‡ Odile Eisenstein,*,§ and Richard A. Andersen*,† †

Department of Chemistry and Chemical Sciences Division of Lawrence Berkeley National Laboratory, University of California, Berkeley, California 94720-1460, United States ‡ LPCNO, Université de Toulouse, INSA, UPS, LPCNO, 135 avenue de Rangueil, F-31077 Toulouse, France, and CNRS, LPCNO, F-31077 Toulouse, France § Institut Charles Gerhardt, Université Montpellier 2, CNRS 5253, cc 1501, Place E. Bataillon, F-34095 Montpellier France S Supporting Information *

ABSTRACT: The experimental reaction of [1,2,4(Me3C)3C5H2]2CeH, Cp′2CeH, and CH3OSO2CH3 begins by α-C−H activation of the SCH3 group, forming Cp′2 CeCH2SO2(OCH3), which evolves into Cp′2CeOCH3 with elimination of CH2 (and presumably SO2). Prolonged heating of this mixture (days at 60 °C) forms Cp′2CeOSO2CH3 and CH3OCH3. The metallacycle [1,2,4-(Me3C)3C5H2][1,2(Me3C)2-4-(Me2CCH2)C5H2]Ce, when presented with the choice of C−H bonds in CH3S and CH3O groups, deprotonates both with comparable rates, ultimately forming Cp′2 CeOCH3 and Cp′2CeOSO2CH3 at 20 °C. The experimental studies are illuminated by DFT calculations on the experimental systems, which show that the hydride selects the more acidic CH3S bond, whereas the metallacycle reacts with C−H bonds of both the CH3S and CH3O groups of CH3OSO2CH3. In the metallacycle reaction, the initially formed regioisomers, Cp′2CeCH2SO2(OCH3) and Cp′2CeCH2OSO2CH3, rearrange to the observed products, Cp′2CeOCH3 and Cp′2CeOSO2CH3, respectively. Furthermore, C−H activation at the SCH3 group forms two isomers of Cp′2CeCH2SO2(OCH3) in the reaction of CH3OSO2CH3 with the metallacycle and only one in the reaction with the hydride. The lack of selectivity in the reactions of the metallacycle relative to the hydride is due to the metallacycle’s greater thermodynamic advantage and lower energy barriers, which are linked to the higher bond energy of Ce−H relative to Ce−C in the metallacycle.

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INTRODUCTION The products of the reaction between [1,2,4-(Me 3 C) 3 C5H2]2CeH, abbreviated Cp′2CeH, and CH3X, where X is a halide, are Cp′2CeX and CH4 (eq 1).1,2 When these C-X bond

Cp′2 CeH + CH3X → Cp′2 CeX + CH4

energy pathway rather than a two-step process in which initial C−H-activation is followed by elimination and trapping of the CH2 fragment by H2. When the CO bond splitting reaction was extended to dialkylethers that contain α- or β−C-H bonds, the net reaction with Cp′2CeH is as shown in eq 2.3 Experimental

(1)

Cp′2 CeH + (RCH2)2 O → Cp′2 Ce(OCH2R) + CH3R

splitting reactions were extended to CH3OCH3 (X = OCH3 in eq 1), an intermediate was isolated and shown to be the methoxymethyl derivative Cp′2Ce(η2-CH2OCH3) that slowly forms Cp′2CeOCH3 and CH4.2 Thus, the net C−O bond splitting reaction, and by implication the C-X bond splitting reactions, proceeds by an initial α−C-H bond activation followed by trapping of the methylene fragment by dihydrogen as Cp′2 CeOCH3 is formed. This mechanistic postulate was supported by DFT calculations on the model metallocenes where Cp′ was replaced by C5H5 that showed a direct one-step process in which the methyl group migrates as an electrophile is a higher © 2012 American Chemical Society

(R = Me, Et, n‐Pr)

(2)

and computational studies showed that these reactions are initiated by α- and β-C−H activation and the β-C−H activation step is followed by elimination of RCHCH2, a pathway Received: September 6, 2011 Published: January 20, 2012 870

dx.doi.org/10.1021/om200842t | Organometallics 2012, 31, 870−881

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approximately 1:2:7. It is noteworthy that no resonances due to A are observed and those due to B are observed only when excess CH3OSO2CH3 is present. The reaction of Cp′2CeH with excess CH3OSO2CH3 may be summarized as illustrated in eq 3,

that is lower in energy than CHR (carbene) elimination. Thus, experimental and computational studies showed that the mechanisms of these seemingly straightforward net C−X bond splitting reactions are not so simple. The question, however, arises as to the generality of the twostep pathway in C−X bond splitting reactions. This article explores this question, the thermodynamic and kinetic control of selectivity in intermolecular C−H bond activation reactions, and the associated question of the effect of the electronegativity of X (eq 1) on the selectivity by presenting Cp′2CeH with a choice between two different CH3X groups within the same molecule, CH3OSO2CH3. The set of elementary steps that comprise the net reaction is explored on the model C5H5 derivatives and the experimental 1,2,4-(Me3C)3C5H2 derivatives, illustrating the necessity, in this case, to use the more complete model in order to bring the computed energies into agreement with experimental data. These studies are of general interest to mechanistic organometallic and organic chemists, since deceptively simple electrophilic substitution reactions may not always proceed by a one-step methyl group transfer pathway and the mechanisms of these seemingly simple reactions are important for a more complete understanding of the thermochemical and kinetic basis of C−H bond activation reactions in stoichiometric and catalytic reactions.4

which shows only the metallocenes and where C and D are Cp′2CeOCH3 and Cp′2CeOSO2CH3, respectively, and A, B, and E are unknown. The net reaction, ignoring B and E, involves an H for OSO2CH3 exchange that proceeds by way of intermediate A. Thus, the reaction between Cp′2CeH and CH3OSO2CH3 is qualitatively similar to that with CH3OCH3, which suggests that A is either Cp′2CeCH2SO2(OCH3) or Cp′2CeCH2OSO2CH3. In a separate NMR experiment, addition of CH3OSO2CH3 to Cp′2CeOCH3 (C) does not yield Cp′2CeOSO2CH3 (D) at 20 °C but, after heating at 60 °C for 2 days, resonances due to D begin to form along with resonances due to CH3OCH3 and the C:D ratio is 12:1. Hydrolysis with D2O of an NMR sample containing A shows deuterium incorporation only in the CH3S site of CH3OSO2CH3 in the 1H and 2H NMR spectra, primarily as CH3OSO2CH2D with a smaller amount of the dideuterated species CH3OSO2CHD2. This finding suggests that A is Cp′2CeCH2SO2(OCH3). In addition, hydrolysis (H2O) when only D and B are present and examination of the hydrolysate by GCMS shows the presence of (Me 3 C) 3 C 5 H 3 and (Me3C)2(Me2EtC)C5H3, symbolized by Cp′H and Cp″H, respectively. This latter point suggests that A forms C by elimination of the fragments CH2 and SO2 and the former fragment is trapped by a C−H bond of Cp′H, as observed previously.1,2 The fate of SO2 is unknown. In earlier studies the reaction between the metallacycle and CH3X was helpful in identifying the intermediates, since H2 is not present and therefore cannot act as a trap. The products of the reaction between CH3OSO2CH3 and the metallacycle are less complex, and their rates of formation are different from those observed with Cp′2CeH, though the net reaction yields C and D (eq 4).

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RESULTS Solution Studies with Methyl Methanesulfonate, CH3OSO2CH3. Mixing Cp′2CeH with a slight excess of methyl methanesulfonate, CH3OSO2CH3, in C6D6 in an NMR tube at 20 °C results in the disappearance of the resonances in the 1H NMR spectrum due to Cp′2CeH and the appearance of three sets of Me3C resonances within 20 min. The new resonances, labeled A−C, are in an approximate area ratio of 2:1:2, respectively. The resonances assigned to C are due to the known compound Cp′2CeOCH3,2,5 but those labeled A and B are unknown; their chemical shifts are given in the Experimental Section. After 1 day at 20 °C, the resonances due to A disappear and the ratio of B to C is 1:2; the ratio does not change over 4 days at 20 °C. Heating for 3 days at 60 °C results in the appearance of a new set of Me3C resonances, labeled D, and the ratio B:C:D is approximately 1:5:1. Further heating for 16 days increases resonances due to D at the expense of C, and the ratio B:C:D after 5 days is 1:2.5:8.5. The resonances labeled D are due to a new compound, Cp′2CeOSO2CH3, prepared and characterized as outlined in the Experimental Section. When the solution is heated for an additional 12 days, the resonances due to C disappear and a new set appears, labeled E, such that the ratio B:D:E is 2.5:24:1; the identity of E is unknown. The reaction between Cp′2CeH and CH3OSO2CH3 proceeds differently when the hydride is in excess, although the outcome of the net reaction is ultimately the same as when CH3OSO2CH3 is in excess. Thus, when Cp′2CeH is in excess, resonances due to C form along with those due to an unknown, F, after about 20 min at 20 °C. The resonances due to Cp′2CeH, C, and F are in the approximate ratio 3:1:1, respectively. This ratio does not change on standing for 1 day at 20 °C or on heating for 1 day at 60 °C. No resonances due to either A or B are observed. Addition of a slight excess of CH3OSO2CH3 results in formation of resonances due to B and an augmentation of the resonances due to C at the expense of F; the ratio B:C is 1:3, and this ratio is constant over 7 days at 60 °C. Adding more CH3OSO2CH3 and heating to 60 °C results in the formation of D; after 16 days the ratio B:C:D is

Addition of excess CH3OSO2CH3 to the metallacycle in an NMR tube in C6D12 at 20 °C results in formation of resonances due to A and C within 20 min. After 40 min the ratio of A to C is 2:1; after 2.5 h resonances due to D appear and the ratio A:C:D is 4:4:1. After 1 day only resonances due to C and D in a ratio of 5:1 are present. Thus, A, C, and D form at 20 °C, which is in contrast to the hydride reaction, where D forms from C at 60 °C. As in the reaction of Cp′2CeH, hydrolysis with D2O when resonances due to A are present in the 1H NMR spectrum and examination of the 2H NMR spectrum show deuterium in the CH 3 S site, identifying A as 871


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with (C5H5)2CeH ([Ce]H), referred to as the small model, and [1,2,4-(Me3C)3C5H2]2CeH ([Ce]′H), referred to as the f ull model. The experimental results show that the hydride and the metallacycle follow different reactivity patterns and a computational study of the reactions of CH3OSO2CH3 with two different models of the metallacycle was initiated. In the small model, symbolized by {Ce}CH2, the metallacycle is modeled as (C5H5)(C5H4C(Me2)CH2)Ce, while the full model, symbolized by {Ce}′CH2, is [1,2,4-(Me3C)3C5H2][1,2-(Me3C)2-4(Me2CCH2)C5H2]Ce. The computational study focuses on the first elementary step between the methyl methanesulfonate and the hydride or the metallacycle. Following formation of an adduct, the first step consists of the activation of a C−H bond of either the CH3S or CH3O group of CH3OSO2CH3. It is also conceivable that O− Me or S−Me bonds are cleaved, i.e. direct methyl transfer, and these reactions are also explored. The reactions that follow these initial steps, with the exception of the step that yields CH3OCH3 from the methoxide complex and CH3OSO2CH3, are not part of the computational study. The computational study using the small model agrees qualitatively with the experimental results for the hydride, but difficulties arise when the calculations for the metallacycle are compared with the experimental results. The presence of the CMe3 substituents on the cyclopentadienyl groups results in subtle differences in the relative energies of the various pathways, and the experimental results can only be reproduced with the full model. For this reason, the results are presented in detail with the full model, which has the advantage that comparisons can be made between the calculated and experimental structures. Comparison of the results between the small and full models provides a better understanding of how the substituents on the cyclopentadienyl rings manipulate the shape of the Gibbs energy profiles and, therefore, a detailed understanding of the observed selectivity that is not properly reproduced with the small model. All of the Gibbs energy profiles are presented at a temperature of 298 K (see Computational Details), and the activation barrier is defined as the difference in Gibbs energy between a transition state and the reactants to which it is directly connected. (2). Reactants. Experimental information on the structure of methyl methanesulfonate is lacking, but DFT and MP2 calculations show a preference for the anti conformation (CH3−O−S−CH3 torsion angle of 180°)8 with the gauche conformation lying only 1−4 kcal mol−1 higher (Scheme 1).

Cp′2CeCH2SO2(OCH3). Hydrolysis (H2O) and analysis by GCMS when only resonances due to C and D are present in the 1H NMR spectrum show the presence of I and II at m/z 312 and 326 amu, respectively, along with Cp′H. The [2 + 4] cycloaddition of sulfene with cyclopentadiene or a substituted cyclopentadiene is one of their characteristic reactions, which suggests that, if sulfene is formed, trapping by Cp′H and Cp″H will yield I and II.6 The cycloadduct between sulfene and C5HMe5 is an isolated compound, suggesting the use of C5HMe5 as a trap in the reaction shown in eq 4. This conjecture is correct, since using C5HMe5 as a solvent for the reaction in eq 4 followed by hydrolysis and analysis by GCMS show that a cycloadduct with m/z 214 amu is obtained. This trapping experiment suggests that I can arise by Cp′ trapping of sulfene when A converts to C. The cycloadduct II can then arise from Cp″, which may be formed by CH2 elimination from the hypothetical intermediate Cp′2CeCH2OSO2CH3. The formation of C and D at 20 °C after 1 day implies that C−H activation of both CH3S and CH3O groups is possible with the metallacycle. This conjecture is supported by a labeling experiment with CD3OSO2CH3, since hydrolysis (H2O) and GCMS analysis of the hydrolysate shows some deuterium incorporation in Cp′H in the case of the metallacycle but not the hydride.

In earlier studies, the reactivity patterns of the hydride and metallacycle were similar when only one type of C−H bond was present. In the present case, two different types of C−H bonds are present and the reaction patterns show some similarities and some differences, summarized as follows. (i) In the reaction of the hydride, D only forms when the reaction mixture is heated to 60 °C, but in the reaction of the metallacycle, D forms at 20 °C. In an independent reaction, isolated Cp′2CeOCH3 and CH3OSO2CH3 do not react at 20 °C, but on heating to 60 °C, D and CH3OCH3 form. Thus, in eq 4, D is not formed from C but is formed in an independent path. (ii) In the reaction of the hydride, hydrolysis yields only Cp′H and Cp″H, showing that CH2 (and by implication SO2) is eliminated as A goes to Cp′2CeOCH3 (C). In the reaction of the metallacycle, Cp′H, Cp″H, and the CH2SO2 cycloadducts I and II are formed, implying that CH2SO2 and CH2 are eliminated in independent paths. (iii) In the reaction of the metallacycle, neither B nor E is formed. It is clear that, although the net reactions illustrated in eqs 3 and 4 are similar, the hydride and metallacycle do not follow identical pathways. These similarities and differences are delineated by DFT calculations described below.

Scheme 1. The Two Conformers of the Methyl Methanesulfonate

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COMPUTATIONAL STUDIES (1). Methods and Models. In previous studies, the Cp′2CeH complex was modeled by (C5H5)2CeH. This simplified model gave results that are consistent with the available experimental information.1−3,7 It is now possible to compute very large models with reasonable computational effort, and this article reports a study of the reactions of CH3OSO2CH3

Similar results obtained with the B3PW91 functional and the basis set given in the Computational Details show that the gauche conformation is 1.4 kcal mol−1 higher in energy than the anti conformation. It should be noted that related sulfonates 872


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such as CF3SO2OCH3,9 CF3SO2OCF3,10 and CH3SO2SCH311 prefer a gauche conformation, as shown by gas electron diffraction and calculations. In CH3OSO2CH3, the calculated O−CH3 distance of 1.437 Å is significantly longer than that of 1.404 Å calculated for CH3OCH3 (the experimental O−C distance in CH3OCH3 is 1.416(5) Å), while the S−C(Me) distance of 1.798 Å is shorter than that of 1.826 Å calculated in CH3SCH3 (the experimental S−C distance is 1.819(5) Å12). These bond distances suggest that the O−CH3 bond is weaker in CH3OSO2CH3 than in dimethyl ether, while the S−CH3 bond energy is comparable to or perhaps greater than in CH3SCH3. However, the thermochemistry of sulfonic acids and esters is plagued with many experimental problems and the ΔHf value of CH3OSO2CH3 is not known with certainty.13 Two views of the calculated structure of [Ce]′H are shown in Figure 1. The structural parameters are in excellent agreement

is slightly shorter than the Ce−CH2 bond in the benzyl complex (2.577 and 2.584 Å for the two isomers found in the solid-state structure15). The ring centroid−Ce−ring centroid angle of 163° shows that cyclometalation of one of the CMe3 groups opens the angle by 8° relative to that in the hydride. The small models for the hydride and the metallacycle have similar shapes relative to the full models but significantly different bond lengths and angles. For instance, in [Ce]H, the Ce−C bonds are slightly shorter (2.807 Å) than in [Ce]′H (2.843 Å) and, more importantly, the ring centroid−Ce−ring centroid angle is much smaller (133° vs 155°). In the metallacycle models, the difference between the ring centroid−Ce− ring centroid angles is even larger: 135° in {Ce}CH2 and 163° in {Ce}′CH2. (3). CH Activation Profiles. Differences between the Hydride and the Metallacycle. (a). Gibbs Energy Profiles. As reported in earlier articles, in which small models of the cerium complexes were used, the mechanism of the reaction between the hydride or the metallacycle and CH3X begins by coordination followed by C−H activation. The C−H activation is best viewed as a proton transfer reaction, which proceeds by way of a metathesis transition state. The pathways are similar when the full model is used. The expected adducts, transition states, associated intermediates, and products for the C−H activation of the CH3S and CH3O groups in the reaction with [Ce]′H and {Ce}′CH2 are shown in Figures 3 and 4. All transition states are described by the group that is activated and the oxygen that is coordinated to the metal. Note that, in all cases but one, the oxygen that is interacting with the metal at the transition state is the same one that is used to coordinate CH3OSO2CH3 to the metal complex. The associated full Gibbs energy profiles for all these reactions are given in the Supporting Information. The Gibbs energy for the adducts, transition states, and intermediates for C−H activation of the CH3S and CH3O groups by the hydride and the metallacycle are shown in Figure 5 using the notation as in Figures 3 and 4. The reactions are initiated by the coordination of CH3OSO2CH3 to [Ce]′H by either a sulfonyl oxygen or the ether oxygen (blue bars in Figure 5). In all cases, the coordination is endoergic but coordination by way of the sulfonyl oxygen is less so. For instance, the adducts associated with the CH3S/O and CH3S/OCH3 transition states have Gibbs energies of 4 and 18 kcal mol−1 relative to separate reactants, respectively. These positive energies of coordination, in which the loss of entropy is not compensated by the coordination enthalpy, are presumably largely due to steric effects. The sulfonyl oxygen binds slightly better to the metallacycle, {Ce}′CH2, than to the hydride, since the change of Gibbs energy for coordination is approximately 0. The factors that make the metallacycle a slightly better Lewis acid are not easy to identify, but distortion of the environment around the metal center associated with the conformation of the rings in the metallacycle may be largely responsible. The Gibbs energies of the transition state relative to the separate reactants, shown as red bars in Figure 5, are an appropriate computational measure of the rate of the reactions, since the adducts have either the same or higher energy than the separate reactants. The preferred reaction of the hydride and the metallacycle is C−H activation of the CH3S group following coordination by the sulfonyl oxygen (reaction CH3S/O in Figure 5). The C−H activation of the CH3O group has significantly higher transition state energies, especially with the

Figure 1. Two views of the DFT optimized structure of Cp′2CeH, [Ce]′H. Ce, C, and H are shown in pink, gray, and white, respectively.

with the experimental parameters.14 The calculated average Ce−C distance of 2.843 Å is close to the experimental value of 2.81(2) Å. The ring centroid−Ce−ring centroid angle of 155° is identical with the experimental value, and the conformation of the two substituted cyclopentadienyl rings is similar to that observed in the solid-state structure. The Ce−H bond distance of 2.124 Å is longer than the distance of 1.90(5) Å found in the solid-state structure. However, it is important to note that the hydride position in the X-ray structure is not well determined, even though it was located and refined isotropically, and the Ce−H distance is most likely artificially short. The calculated structure of the metallacycle, two views of which are shown in Figure 2, cannot be compared to an

Figure 2. Two views of the DFT optimized structure of the metallacycle {Ce}′CH2. Ce, C, and H are shown in pink, gray, and white, respectively.

experimental structure, because none is available. The calculated structure may, however, be compared with the solid-state structure of the two benzyl derivatives Cp′2Ce(CH2C6H5) and Cp′2Ce(CH2-4-MeC6H4). The calculated average Ce−C(Cp′) distance of 2.843 Å in the metallacycle is similar to that in the hydride. The Ce−CH2 distance of 2.536 Å 873


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Figure 3. Adducts, transition states, intermediates, and products for the C−H activation of the CH3S group, with transition state labels indicating the activated and the coordinated group. For example, CH3S/O indicates activation of CH3S by way of coordination through the oxygen of the sulfonyl group.

difference in reactivity of the hydride and metallacycle, will be amplified in the Discussion. The computational results presented so far have been obtained using the full model. The reader might wonder why, since earlier studies in this series used the small model. The reason is that, in the present case, the computational results using the small model do not account for the experimental results (Figure 6). For the small model, the coordination of CH3OSO2CH3 is exoergic for the hydride and the metallacycle, which is not the case in the full model. Furthermore, the hydride exhibits a slightly larger Gibbs coordination energy than the metallacycle in the small model, which is opposite to the result found with the full model. This is perhaps related to the different steric environments, since the small model of the metallacycle contains a CMe2CH2 group and is intrinsically more bulky than C5H5 in Cp2CeH. This unbalanced situation is not present in the calculations with the full model, where the number of carbons is the same for the hydride and the metallacycle. The C−H activation calculated in the small model is more exoergic with the metallacycle than with the hydride by about 11 kcal mol−1, a result that is similar to that obtained with the full model. However, the Gibbs activation barriers, calculated

hydride. The energy of the transition state is largely determined by the acidity of the C−H group, since the reaction is a proton transfer. The preferred reaction of the CH3S group stems from its greater acidity compared to that of the CH3O group.16 The transition state energies are higher for the hydride than for the metallacycle by 4 kcal mol−1 for C−H activation of CH3S and by 5 kcal mol−1 for that of CH3O. This difference is essentially due to the less favorable coordination energy of CH3OSO2CH3 in the hydride, which makes the metallacycle slightly more reactive (kinetically) than the hydride. The metallacycle also has a thermodynamic advantage over the hydride in all of these reactions, since the intermediates (green bars) are all at lower energies. A thermodynamic cycle, written in term of Gibbs energies and enthalpies, illuminates this energy difference (Scheme 2). In this cycle, step 1 is the hydrogenation of the metallacycle, for which the calculations give a Gibbs energy of −12.4 kcal mol−1 and an enthalpy of reaction of −22.3 kcal mol−1. Step 2 is the C−H activation reaction with the metallacycle, and step 3 is the C−H activation with the hydride. This cycle shows that when the metallacycle and the hydride give the same products, the Gibbs energy and enthalpy of the reaction differ by 12.4 and 22.3 kcal mol−1, respectively, in favor of the metallacycle. The concept, which is key to understanding the 874


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Figure 4. Adducts, transition states, intermediates, and products for the C−H activation of the CH 3O group with transition state labels as in Figure 3.

Scheme 2. Thermodynamic Cycle, Using Gibbs Energies (Upper and Black) and Enthalpies (Lower and Purple) for the C−H Activations by the Metallacycle and the Hydride

Figure 5. Gibbs energies, in kcal mol−1, for the adducts (blue), transition states (red), and intermediates (green) for the C−H activation of CH3OSO2CH3 with the full model for the hydride and the metallacycle. See Figures 3 and 4 for notation.

results; these results are described in detail in the Supporting Information. (b). Geometries of Intermediates, Transition States, And Products. Only the geometries of representative transition states and intermediates are discussed for the full models. The extrema that are not discussed here are shown in the Supporting Information, as are all of the Gibbs energy profiles.

from the adducts in the small model, are similar for the hydride and the metallacycle, which conflicts with the experimental 875


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Figure 8. DFT optimized structures of (a) A and (b) the isomer of A, with selected distances in Å. Ce, O, S, C, and H are shown in pink, red, yellow, gray, and white, respectively.

Figure 6. Gibbs energies, in kcal mol−1, for the adducts (blue), transition states (red), and intermediates (green) for the C−H activation of CH3OSO2CH3 with the small model for the hydride and the metallacycle. See Figures 3 and 4 for notation.

that is an isomer of A (Figure 8b). This four-membered ring has an elongated S−CH2 bond as in A and also an elongated intracyclic S−O bond. Therefore, this isomer can yield CH2 and SO2 or the sulfene, CH2SO2, without rearrangement on forming the methoxide [Ce]′OCH3. This isomer of A can account for the formation of the sulfene. The transition state for C−H activation of the CH3O group with the ether oxygen coordinated to cerium has a structure similar to that found in the C−H activation of CH3OCH3 (Figure 9a).2 This transition state leads to an intermediate with

The transition states for the C−H activation at the CH3S group by the hydride and the metallacycle, shown in parts a and b of Figure 7, respectively, highlight the similarities between

Figure 7. Transition states for the C−H activation of the CH3S group of CH3OSO2CH3 by (a) [Ce]′H and (b) {Ce}′CH2, with selected distances in Å and angles in deg. Ce, O, S, C, and H are shown in pink, red, yellow, gray, and white, respectively.

these two reactions. In both cases, the C−H activation is a proton transfer, as shown by the almost linear alignment of the three active atoms in the transition state. The C−H activation of the CH3S group by either the hydride or the metallacycle results in an intermediate, A (Figure 8a), that contains a four-membered ring. The Ce−C distance of 2.751 Å is longer than that of 2.536 Å calculated in the metallacycle {Ce}′CH2. The methoxy group is pointing away from the cerium complex, as this orientation avoids the bulky CMe3 groups. The significant elongation of the S−CH2 distance from 1.445 Å in the adduct of CH3OSO2CH3 to 1.714 Å in A shows that the bond is significantly activated, and the intracyclic S−O bond is only elongated by 0.04 Å relative to that in the adduct. The loss of CH2 and of SO2 from A has not been studied by computations. The C−H bond activation of the CH3S group using the OMe group to form the adduct yields a four-membered ring

Figure 9. C−H activation at the CH3O group coordinated through the ether oxygen using the full model, with selected distances in Å: (a) transition state; (b) intermediate. Ce, O, S, C, and H are shown in pink, red, yellow, gray, and white, respectively.

a three-membered ring shown in Figure 9b, which can be compared to the experimentally observed structure of Cp′2Ce(η2-CH2OCH3). The Ce−C distance of 2.568 Å is longer than the equivalent distance of 2.488(4) Å in Cp′ 2 Ce(η 2 CH2OCH3). The C−O distance of 1.458 Å is close to that of 1.466 Å in Cp′2Ce(η2-CH2OCH3), while the Ce−O bond distance of 2.594 Å is significantly longer than the value of 2.406 Å in Cp′2Ce(η2-CH2OCH3), which may be due to the electron-withdrawing influence of the SO2CH3 group. The threemembered ring yields Cp′2CeOSO2Me (D) by elimination 876


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Figure 10. Methyl transfer reactions with the hydride, metallacycle, and methoxide.

The preferred methyl transfer occurs by cleaving the O−Me bond rather than the S−Me bond (see the Supporting Information). All of the reactions have relatively high activation barriers, although they are all strongly exoergic because they form Cp′2CeOSO2CH3 with Ce−O bonds. The methoxide reaction is the least exoergic, because the reactant already has a Ce−O bond. The transition states for the direct methyl transfer reactions are relatively close in energy to that of the C−H activation of OCH3 for both the hydride and the metallacycle. In the case of the hydride, the energy of the methyl transfer transition state is marginally lower than that of the C−H activation (29 kcal mol−1 vs 30 and 31.3 kcal mol−1). In the case of the metallacycle, the methyl transfer has an activation barrier 3.6 kcal mol−1 higher than that of the more favorable C−H activation of the CH3O group. The direct methyl transfer might compete with the activation of the CH3O group but not with that of CH3S. The high activation barrier in the case of the methoxide is in agreement with experiment, since prolonged heating at 60 °C is required for the transformation of C to D. These trends are also obtained in the small model; these data are documented in the Supporting Information. The geometry of the transition state for the reaction of the methoxide with CH3OSO2CH3 is shown in Figure 12a. The reaction starts with an energetically demanding coordination of CH3OSO2CH3 by way of the sulfonyl oxygen, since the resulting adduct has a Gibbs energy of 11.7 kcal mol−1. The product of the reaction is a κ2-coordinated OSO2CH3. The methyl transfer takes place by way of an SN2 transition state. Despite the relatively short distances between the bridging methyl and the two oxygens, the methyl transfer is a highenergy process.

of CH2, which has not been studied computationally. Thus, the activation at the C−H bond of the CH3O group yields D directly. The transition states for C−H activation of the CH3O group with the sulfonyl oxygen coordinated to Ce for both the metallacycle and hydride are shown in the Supporting Information. For the metallacycle, this transition state leads to an intermediate that could also form D by elimination of CH2. (4). Methyl Transfer Reactions. The calculated Gibbs energies in the full model for the methyl transfer from CH3OSO2CH3 to the hydride, the metallacycle, and the methoxide for the reactions shown in Figure 10 are illustrated in Figure 11. The Gibbs energy profiles are shown in the Supporting Information.

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DISCUSSION The experimental and computational studies described above show that the reaction pathways chosen when Cp′2CeH is mixed with CH3OSO2CH3 follow the pattern observed in CH3X (X = halide, methoxide) reactions.1,2 The mechanism in

Figure 11. Gibbs energies, in kcal mol−1, for the methyl transfer with the full model for the hydride, metallacycle, and methoxide. 877


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calculated to be exoergic by 12.4 kcal mol−1 and exothermic by 22.3 kcal mol−1, the metallacycle has a thermodynamic advantage relative to the hydride in reactions resulting in the same product.

The value of the exothermicity of the hydrogenation reaction is the difference between the bond dissociation enthalpies of the Ce−C and the Ce−H bonds. In the thermochemical cycle shown in Scheme 3, the calculated enthalpy for the hydrogenation Scheme 3. Thermochemical Cycle, ΔH, for Hydrogenation of the Metallacycle

Figure 12. OCH3/OSO2CH3 exchange by way of methyl transfer, with selected distances in Å: (a) transition state; (b) product. Ce, O, S, C, and H are shown in pink, red, yellow, gray, and white, respectively.

all of these reactions is a two-step process that involves a proton transfer, by way of a metathesis transition state, forming a carbenoid intermediate, followed by trapping of the methylene fragment by either H2 or the C−H bonds in the CMe3 groups in the Cp′ rings. The Gibbs energy barrier of the proton transfer is lower than that of a direct methyl transfer, due to the energy required to form a five-coordinate carbon in the transition state of the latter pathway. Methyl methanesulfonate was chosen, as it provides methyl groups that differ in their hydrocarbon acidities; the C−H bonds of CH3S are likely to be more acidic than those of CH3O, since the gas-phase deprotonation enthalpies of H−CH2SMe and H−CH2OMe are 393 and 407 kcal mol−1, respectively.16c The experimental results and computed Gibbs energy profiles show that Cp′2CeH selects the C−H bonds of the CH3S group, a result that is consistent with the view that the reaction is a proton transfer. In earlier articles, the reactions between CH3X and either the hydride or metallacycle follow similar pathways to the same products. The only difference observed is that the reaction of the metallacycle does not evolve H2 and therefore it is not present as a trap. This difference is exploited to obtain information about intermediates from deuterium labeling and monitoring by NMR spectroscopy. The reactions between the metallacycle and CH3OSO2CH3 do not follow this pattern since, although the net reactions are the same, the metallacycle reacts with C−H bonds of both the CH3S and CH3O groups. Thus, the reaction of the metallacycle and CH3OSO2CH3 generates Cp′2CeCH2SO2(OCH3), called A, which is inferred by labeling studies and NMR spectra; Cp′2CeCH2OSO2CH3 is not observed, but it is required, since Cp′2CeOSO2CH3 is formed at 20 °C. In the hydride reaction, A and Cp′2CeOCH3 form at 20 °C, since the methoxide is stable to CH3OSO2CH3 at 20 °C and only slowly forms Cp′2CeOSO2CH3 and CH3OCH3 at 60 °C. Thus, the relative rates and, accordingly, the mechanisms of these two reactions that both yield Cp′2CeOSO2CH3, are different. Cp′2CeH reacts only with the C−H bonds of the CH3S group and is therefore more selective than the metallacycle. The origin of the selectivity is traced to the relative values of the Gibbs energy of the two starting materials, which are linked by the hydrogenation/dehydrogenation reaction shown in eq 5. Since the hydrogenation of the metallacycle is

of the metallacycle (step 1) is 22.3 kcal mol−1. The homolytic cleavage of the Ce−C bond of the metallacycle yields a pair of radicals {Ce}′•C• (step 2), and the homolytic cleavage of the CeH bond yields a pair of radicals [Ce]′•H• (reverse of step 4). The hydrogenation of the carbon radical by H2 to form a C−H bond and a hydrogen radical is approximately thermoneutral (step 3). The consequence of the cycle is that there is a difference of 22.3 kcal mol−1 in the Ce−H and Ce−C bond dissociation enthalpies, the Ce−H bond being the stronger. There is no experimental information on the Ce−C bond dissociation enthalpy in the metallacycle. However, the available experimental thermochemical data show that metallocene lanthanide− hydride bonds have slightly higher bond dissociation enthalpies than the corresponding lanthanide−alkyl bonds by about 2−8 kcal mol−1.17 A thermochemical cycle for the hydrogenolysis of Cp′2CeCH3 is shown in Scheme 4, using the calculated value Scheme 4. Thermochemical Cycle, ΔH, for Hydrogenation of the Methyl Complex

of 17 kcal mol−1 for the enthalpy change in step 1: viz., the bond dissociation enthalpy of Ce−H is greater than that of 878


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Ce-CH3 by 17 kcal mol−1. The values in Schemes 3 and 4 show that the bond dissociation enthalpies (BDE) decrease in the order Ce−H > Ce-CH3 > Ce−C (metallacycle). A comparison between calculated and experimental values is not possible, since the latter are unavailable. However, it is established that BDE’s for M−H are higher than for M−CH3 for early d and f transition metal metallocenes; in (C5Me5)2MH2 and (C5Me5)2 M(CH3)2 (M = Zr, Hf) the difference is 10 kcal mol−1.17a More importantly, hydrogenolysis of the metallacycle (C5Me5)(C5Me4CH2CH2CH2)Sc shows that the BDE of Sc−H is 18 kcal mol−1 higher than that of Sc−C(metallacycle).17b DFT calculations show that the reaction between (C5H5)2ScCH3 and H2 is exothermic by 10 kcal mol−1.17c Calculations using the methodology described above for the hydrogenolysis of the Ce− CH3 bond in [Ce]-CH3 (small model) shows that the reaction is exothermic by 13 kcal mol−1; the exothermicity increases to 17 kcal mol−1 in the full model (Scheme 4) and 22 kcal mol−1 in the full model of the metallacycle (Scheme 3). While it is not possible to validate these values quantitatively, they are clearly qualitatively aligned with the experimental evidence. This, in turn, validates the relative energy profiles calculated for the reaction of CH3OSO2CH3 with the hydride and metallacycle, showing that the hydride is more selective than the metallacycle. The results of the calculations are in agreement with the experimental results only when the full model is used. However, the small model does correctly reproduce several features of the reactions, and the results obtained at this simpler level of approximation are not without value qualitatively. As is the case with the calculations on the full model, the small model (i) gives lower energy barriers for the C−H activation of the CH3S group compared to the CH3O group, (ii) predicts that the reactions of the metallacycle will be more thermodynamically favorable relative to the hydride, (iii) shows that the methyl transfer is a high-energy process, and (iv) gives the same general shapes to all extrema. The approximations inherent in the small model, however, do not capture important characteristics of the experimental molecules. The small model retains its utility for preliminary exploration of reaction pathways but may be misleading, unless the difference in the energy profiles is large, as in earlier publications on the reactions of CH3X (X = halide, alkoxide with Cp′2CeH).1,2 When the difference in the energy profiles is small and the question of selectivity is being addressed, as in the present case, the more computationally demanding full model is essential.

weaker Ce−C bond in comparison to the Ce−H bond. Since the less reactive hydride does not possess this advantage, it selects the more acidic C−H bond in CH3S rather than CH3O, while the more reactive metallacycle does not discriminate between them.

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EXPERIMENTAL SECTION

General Considerations. All manipulations were performed under an inert atmosphere using standard Schlenk and drybox techniques. All solvents were dried and distilled from sodium or sodium benzophenone ketyl. Methyl methanesulfonate, CH3OSO2CH3, and trimethylsilyl methanesulfonate, (CH3)3SiOSO2CH3, were obtained commercially and purified by distillation followed by vacuum transfer. Methyl-d3 methanesulfonate, CD3OSO2CH3, was prepared from commercial methanol-d3 and methylsulfonic anhydride and purified as above.18 NMR spectra were recorded on Bruker AV-300, AV-400, and AV600 spectrometers at 20 °C in the solvent specified. J. Young NMR tubes were used for all NMR tube experiments. Electron impact mass spectrometry and elemental analyses were performed by the microanalytical facility at the University of California, Berkeley. The abbreviation Cp′ is used for the 1,2,4-tri-tert-butylcyclopentadienyl ligand, and the metallacycle refers to the complex [1,2,4-(Me3C)C5H2][1,2-(Me3C)2-4-(Me2CCH2)C5H2]Ce. Unless otherwise specified, samples for GC-MS were prepared by adding a drop of nitrogenpurged H2O, agitating, and allowing the samples to stand closed for 10 min. The samples were then dried over magnesium sulfate, filtered, and diluted 10-fold with pentane. A 1 μL sample was injected into a HP6890 GC system with a J&W DB-XLB universal nonpolar column, attached to an HP5973 Mass Selective Detector. Cp′2CeOSO2CH3 (D). Cp′2CeCl5 (0.20 g, 0.31 mmol) was suspended in toluene (20 mL), (CH3)3SiOSO2CH3 (150 μL, 0.97 mmol) was added via syringe, and the yellow solution was stirred at room temperature. After 1 day, the solution color had not changed. An aliquot was taken to dryness, and its 1H NMR spectrum in toluene-d8 contained five paramagnetic resonances in a 2:3:18:18:18 area ratio and no resonances due to Cp′2CeCl. The solution was taken to dryness, and the yellow solid residue was dissolved in pentane (10 mL) and filtered. The solution was reduced in volume until precipitation occurred, warmed to dissolve the precipitate, and cooled to −15 °C, yielding clusters of the yellow crystals. Yield: 0.11 g (0.16 mmol, 52%). Mp: 258−260 °C. 1H NMR (C7D8, 300 MHz): δ 11.67 (2H, ν1/2 = 96 Hz), 3.53 (3H, ν1/2 = 35 Hz), 1.44 (18H, ν1/2 = 60 Hz), −2.61 (18H, ν1/2 = 60 Hz), −12.19 (18H, ν1/2 = 60 Hz). MS (M)+ m/z (calcd, found): 701 (100, 100), 702 (41, 40), 703 (25, 25), 704 (8, 9). Anal. Calcd for C35H61CeO3S: C, 59.88; H, 8.73. Found: C, 59.50; H, 9.03. NMR Tube Reaction of Excess CH3OSO2CH3 and Cp′2CeH in Benzene-d6 or Cyclohexane-d12. Cp′2CeH14 was dissolved in benzene-d6 in an NMR tube, and an excess of CH3OSO2CH3 was added. The purple solution immediately turned orange, and after 2 min, resonances due to Cp′2CeH had disappeared from the 1H NMR spectrum, while resonances due to Cp′2CeOCH3 (C)2,5 and two other species, A and B, had appeared. Species A had three CMe3 resonances in a 1:1:1 area ratio (1H NMR (C6D6) δ 1.20 (ν1/2 = 100 Hz), −1.57 (ν1/2 = 70 Hz), −11.54 (ν1/2 = 100 Hz)), while B had two CMe3 resonances in a 2:1 area ratio (1H NMR (C6D6) δ −0.65 (ν1/2 = 150 Hz), −11.84 (ν1/2 = 70 Hz)). The ratio of A, B, and C was 2:1:2. After 1 day, the resonances due to A had disappeared from the 1H NMR spectrum, and the ratio of B to C was 1:2. A small amount of authentic Cp′2CeOCH3 was added to confirm the identity of the resonances of C, changing the ratio of B and C to 1:3. After 4 days, the ratio was unchanged, and the sample was heated to 60 °C. After 2 days, resonances due to Cp′2CeOSO2CH3 (D) had appeared; the ratio of B, C, and D was 1:5:1. After 9 days, the ratio was 1:6:2.5. More CH3OSO2CH3 was added, and the sample was heated further at 60 °C. After 5 days, the ratio was 1:2.5:8.5. After 12 days, resonances due to C had disappeared from the spectrum, and a pair of new CMe3 resonances in a 2:1 area ratio, E, had appeared (1H NMR (C6D6) δ −2.98 (ν1/2 = 20 Hz), −8.55 (ν1/2 = 20 Hz)). The ratio of B, D, and E

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CONCLUSIONS The experimental and computational study shows that when CH3OSO2CH3 is allowed to react with either the hydride or the metallacycle, the ultimate product Cp′2CeOSO2CH3 (D) is the same but the pathways followed are different. In the case of the hydride, the initial intermediate formed at 20 °C, Cp′2CeCH2SO2(OCH3) (A), evolves into Cp′2CeOCH3 (C), which then slowly forms Cp′2CeOSO2CH3 (D) and CH3OCH3 at 60 °C in the presence of CH3OSO2CH3. In contrast, the metallacycle forms Cp′ 2 CeCH 2 SO 2 (OCH 3 ) (A) and Cp′2CeCH2OSO2CH3, which rearrange to Cp′2CeOCH3 (C) and Cp′2CeOSO2CH3 (D) at 20 °C, respectively. The DFT calculations with the full model of the cerium metallocene compounds show that the selectivity of the hydride for the C−H bonds in the CH3S group and the lack of selectivity displayed by the metallacycle is traced to a significant thermodynamic and a slight kinetic advantage, which is ultimately linked to the 879


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NMR Tube Reaction of Excess CH3OSO2CH3 and Cp′2CeH or the Metallacycle, Quenched with D2O. Cp′2CeH was dissolved in cyclohexane in an NMR tube, and an excess of CH3OSO2CH3 was added. The purple solution immediately turned orange, and after 10 min, the sample was taken to dryness and dissolved in benzene-d6. A drop of D2O was added under nitrogen flush and the sample was agitated, allowed to stand for 10 min, and filtered through anhydrous sodium sulfate. The resonance in the 1 H NMR spectrum corresponding to the S-bound methyl group consisted of a sharp singlet (−SCH3, δ 1.99), a 1:1:1 pattern (−SCH2D, δ 1.98, JHD = 2 Hz), and a 1:3:5:3:1 pattern (−SCHD2, δ 1.97, JHD = 2 Hz). The ratio of the three signals was approximately 52:8:1. The resonance corresponding to the O-bound methyl group consisted of a sharp singlet (δ 3.07). The 2H NMR spectrum contained a resonance corresponding to the S-bound methyl group (−SCH2D, t, δ 1.95, JHD = 2 Hz) but no resonance corresponding to the O-bound methyl group. Identical spectra were obtained when the metallacycle was substituted for Cp′2CeH in this procedure. NMR Tube Reaction of CH3OSO2CH3 and the Metallacycle in Neat Cp*H. Cp′2Ce(CH2C6H5) was dissolved in cyclohexane and heated at 60 °C for 12 h, yielding a solution of the metallacycle. The sample was taken to dryness and dissolved in Cp*H. The sample was heated at 60 °C for 12 h, and the 1H NMR spectrum still contained resonances due to the metallacycle. An excess of CH3OSO2CH3 was added, and the dark purple solution immediately turned orange. The sample was heated to 60 °C for 12 h, and the only paramagnetic resonances in the 1H NMR spectrum were those of Cp′2CeOCH3 (C). Integration relative to the Cp*H resonance indicated that approximately 70% of the metallacycle starting material had been converted to C. GC-MS analysis of the hydrolysate showed a component in addition to Cp′H and Cp*H with (M)+ m/z 214 (Cp*H + O2SCH2). Computational Details. The Stuttgart−Dresden−Bonn relativistic large effective core potential (RECP) was used to represent the inner shells of Ce.19 The associated basis set19 augmented by an f polarization function (α = 1.000) was used to represent the valence orbitals.20 Sulfur has also been represented by an RECP,21 with the associated basis set augmented by a d polarization Gaussian function (α = 0.421). The atoms C, O, and H were represented by an allelectron 6-31G(d,p) basis set.22 Calculations were carried out at the DFT(B3PW91) level23 with Gaussian 03.24 The nature of the extrema (minimum or transition state) was established with analytical frequency calculations, and the intrinsic reaction coordinate (IRC) was followed to confirm that the transition states connect to reactants and products. The zero-point energy (ZPE) and entropic contribution have been estimated within the harmonic potential approximation. The Gibbs free energy, G, was calculated at T = 298.15 K and 1 atm. The difference of enthalpies for Schemes 3 and 4 were assimilated to difference of potential energies. It was verified that the difference enthalpies give the same values.

was 2.5:24:1. GCMS analysis of the hydrolysate showed components in addition to Cp′H with (M)+ m/z 248 (Cp′H + CH2). The experiment was repeated using cyclohexane-d12 solvent, which did not affect the course of the reaction. The experiment was repeated using CD3OSO2CH3, and GC-MS analysis of the hydrolysate showed no appreciable perturbation of the isotope distribution in the molecular ion of Cp′H. MS (M)+ m/z (calcd, found): 234 (100, 100), 235 (19, 20), 236 (2, 2). NMR Tube Reaction of CH3OSO2CH3 and Excess Cp′2CeH in Cyclohexane-d12. Cp′2CeH was dissolved in cyclohexane-d12 in an NMR tube, and a small amount of CH3OSO2CH3 was added. The purple sample became redder, and after 2 min, the 1H NMR spectrum contained resonances due to Cp′2CeH, Cp′2CeOCH3 (C), and a new species, F, with CMe3 resonances in a 2:1 area ratio (1H NMR (C6D6) δ −2.1 (ν1/2 = 850 Hz), −10.84 (ν1/2 = 100 Hz)). The ratio of Cp′2CeH, C, and F was approximately 3:1:1. The ratio did not change upon standing at 19 °C for 1 day or at 60 °C for an additional 1 day. A slight excess of CH3OSO2CH3 was added, and the solution immediately became red. After 2 min, resonances due to Cp′2CeH had disappeared from the 1H NMR spectrum, and resonances due to B appeared, overlapping those of F. After 5 days at 19 °C, the resonances due to F had disappeared from the spectrum, and the ratio of B and C was 1:3. After 7 days at 60 °C, the ratio had not changed. Another drop of CH3OSO2CH3 was added. After 10 min, the ratio of the paramagnetic species in the 1H NMR spectrum had not changed. The sample was heated at 60 °C for 3 days, and resonances due to Cp′2CeOSO2CH3 (D) appeared in the 1H NMR spectrum; the ratio of B, C, and D was 1:4:2. After 8 days, the ratio was 1:3:3. After 16 days, the ratio was 1:2:7. NMR Tube Reaction of CH3OSO2CH3 and Cp′2CeOCH3 (C) in Benzene-d6. Cp′2CeOCH3 (C) was dissolved in benzene-d6 in an NMR tube, an excess of CH3OSO2CH3 was added, and the sample was allowed to stand at 19 °C. After 2 days, the 1H NMR spectrum contained only resonances due to Cp′2CeOCH3. The sample was heated to 60 °C, and after 2 days, resonances due to Cp′2CeOSO2CH3 (D) and CH3OCH3 had appeared in the spectrum; the ratio of C to D was 12:1. After 5 days at 60 °C, resonances due to E had appeared; the ratio of C, D, and E was 17:17:1. After 12 days, the ratio was 1:6.5:1. After 23 days, the ratio was 1:14.5:5. NMR Tube Reaction of CH3OSO2CH3 and the Metallacycle in Cyclohexane-d12, with and without D2. Cp′2Ce(CH2C6H5)14 was dissolved in cyclohexane-d12 and heated at 60 °C for 12 h, yielding a solution of the metallacycle. An excess of CH3OSO2CH3 was added, and the dark purple solution immediately turned orange. After 5 min at 19 °C, resonances due to the metallacycle had disappeared from the 1 H NMR spectrum, while resonances due to A and Cp′2CeOCH3 (C) had appeared in a 6:1 area ratio. After 40 min, the ratio was 2:1. After 2.5 h, resonances due to Cp′2CeOSO2CH3 (D) had appeared; the ratio of A, C, and D was 4:4:1. After 1 day, the resonances due to A had disappeared from the 1H NMR spectrum, and the ratio of C and D was 5.5:1. Integration relative to the solvent residual proton peak showed that approximately half of the Cp′2Ce(CH2C6H5) starting material had been converted to C. The sample was heated to 60 °C, and after 2 days, the ratio C and D was 1:1. GCMS analysis of the hydrolysate showed components in addition to Cp′H with (M)+ m/z 248 (Cp′H + CH2), 312 (Cp′H + O2SCH2), and 326 (Cp′H + O2SCH2 + CH2). The experiment was repeated, but 10 min after CH3OSO2CH3 was added and the solution turned orange, the sample was immersed in a liquid nitrogen/isopropyl alcohol bath and the headspace was evacuated and replaced with D2. The sample was warmed to 19 °C, and after 10 min, the 1H NMR spectrum contained only resonances due to A and Cp′2CeOCH3 (C) in a 3:1 area ratio. After 1 day, the spectrum contained resonances due to C and D in a 6:1 area ratio. The 2H NMR spectrum did not contain resonances indicating D incorporation into CH3OSO2CH3. The experiment was repeated using CD3OSO2CH3, and GC-MS analysis of the hydrolysate showed a shift in the isotope distribution in the molecular ion of Cp′H toward higher isotopologues. MS (M)+ m/z (calcd, found): 234 (100, 100), 235 (19, 31), 236 (2, 4).

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ASSOCIATED CONTENT

S Supporting Information *

Figures, tables, and text giving full Gibbs energy profiles corresponding to all paths studied, coordinates, energy, and Gibbs energy of all calculated structures, graphical representations of structures not shown in the text, and complete data concerning the results with the small model. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (L.M.); Odile. [email protected] (O.E.); [email protected] (R.A.A.).

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ACKNOWLEDGMENTS This work was supported by the Director, Office of Science, Office of Basic Energy Sciences (OBES), of the U.S. Department of 880


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(21) Bergner, A.; Dolg, M.; Küchle, W.; Stoll, H.; Preuß, H. Mol. Phys. 1993, 80, 1431. (22) Hariharan, P. C.; Pople, J. A. Theor. Chim. Acta 1973, 28, 213. (23) Perdew, J. J. P.; Wang, Y. Phys. Rev. B 1992, 45, 13244. Becke, A. D. J. Chem. Phys. 1993, 98, 5648. Burke, K.; Perdew, J. P.; Yang, W. In Electronic Density Functional Theory: Recent Progress and New Directions; Dobson, J. F., Vignale, G., Das, M. P., Eds.; Plenum: New York, 1998. (24) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Montgomery, Jr., J. A.; 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.; Bakken, V.; 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.

Energy (DOE) under Contract No. DE-AC02-05CH11231. L.M. and O.E. thank the CNRS and le Ministère de l’Enseignement Supérieur et de la Recherche for funding. L.C. thanks the Computer Centers, CCRT of the CEA, CINES, and CALMIP, for a generous donation of computational time. L.M. is a junior member of the Institut Universitaire de France.

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REFERENCES

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