Edmond Collange

Alberto Flamini

Rinaldo Poli
email: rinaldo.poli@u-bourgogne.fr

,2-tricyanoethanide in Water: A Completely Sorted out Square Scheme

Stopped-Flow and DFT Studies of Proton Transfer and

Isomerization of 5-Amino-3-imino-1,2,6,7-tetracyano-3H-pyrrolizine and Its Related Base 2-(5-Amino-3,4-dicyano-2H-pyrrol-2ylidene)-1,

Introduction

The title pyrrolizine compound HL (see Chart 1) [Fares, 1995 #4905] has attracted our attention because it is a species of prospective use in materials chemistry: it is a fluorescent dye, it can be sublimed to afford thickness-controlled thin films, and it coordinates to a number of transition metal cations. [Bonamico, 1989 #4915;Bonamico, 1991 #4906;Bonamico, 1992 #4907;Bonamico, 1993 #4908;Bonamico, 1993 #4909;Flamini, 2000 #4629] Like other similar conjugated heteropolycyclic molecules and their conjugate anions and complexes, [Preininger, 1996 #5053;Mills, 1995 #5054;Klein, 1994 #5055;Wolfbeis, 1997 #5056;Murkovic, 1997 #5057;Plaschke, 1995 #5058;Donckt, 1994 #5059;Exstrom, 1995 #5060;Daws, 1997 #5061] it shows suitable properties for application in the construction of optochemical sensors. [Fares, 1998 #4910;Panusa, 1996 #4911;Flamini, 1999 #4912] HL, whose bicyclic nature is shown by an Xray crystallographic study, [Fares, 1995 #4905] is prepared by protonation in water of the sodium salt of the anion L' -(Chart 1), which derives from a reductive autocondensation of tetracyanoethylene. [Flamini, 1992 #4913] The "open" form of the anion shown in Chart 1 has been evidenced by an X-ray structure of the tetraphenylarsonium salt. [Fares, 1995 #4914] However, the anion rearranges to a "closed" bicyclic form also upon coordination to metal centers (see ML in Chart 1). <Chart 1 > Prior to the present investigation, no information was available on the existence of a noncoordinated "closed" form of the anion, i.e. L -, and of an "open" form of the acid, HL ' , nor on the acid dissociation constants of HL and HL', nor on the rates and mechanisms of the opening and the reverse closing reactions on either the neutral molecule or the anion. These questions have been addressed in this study.

Experimental Section

Materials. Compounds NaL' . xH2O and HL were synthesized according to the published procedures. [Flamini, 1992 #4913;Fares, 1995 #4905] Standard solutions of NaOH and HCl were obtained from Prolabo Normadoses. Buffers or buffer constituents (acids and conjugated bases) solutions were prepared from usual commercial products: sodium formate, formic acid, sodium acetate, acetic acid, sodium monochloroacetate, monochloroacetic acid, boric acid, hexamethylenetetramine, imidazole, hydroxylamine, histidine, sodium phosphate, sodium hydrogen phosphate, sodium dihydrogen phosphate, and sodium carbonate. Sodium chloride was used to adjust the ionic strength to 0.1 (with the exception of the very acidic or basic solutions).

Instrumentation. IR spectra were measured with a Perkin-Elmer 16F PC FT spectrometer, visible spectra were recorded on Varian Cary 5 and Cary 1 spectrophotometers. The stopped flow kinetic runs were carried out with a Hitech SF-61-DX2 apparatus coupled with a Hitech diodearray UV-visible spectrophotometer.

Kinetic measurements. (a) By stopped flow. Kinetic traces were recorded at 25 °C in the pH range 2-12. The cell path length is 1 cm. The pH values for the buffer solutions were determined with a Tacussel Minisis pHmeter, calibrated at 25 °C with NBS standards. For the very acidic solutions (pH < 1) the acidity function H0 is used in place of pH. [Paul, 1957 #5205] Two types of manipulations were carried out as follows.

Type 1 experiments (from HL). A 2.5 . 10 -3 M HL stock solution was prepared in acetonitrile, since this compound is only sparingly soluble in water. For each stopped-flow kinetics run, the HL solution was prepared by diluting this stock solution 1:50 (v/v) with 0.1 M HCl ([HL] = 5 . 10 -5 M). This operation was carried out no more than 1 min prior to injection into the stopped-flow cell. Under these conditions, the advancement of the conversion reaction prior to injection (k  10 -4 s -1 , see Results and Discussion) is negligible. The above solution was mixed in the stopped-flow cell with an equivalent volume of an aqueous solution containing NaOH (for the kinetic run at pH ≥ 12) or weak base B at appropriate concentration (thereby forming a B/BH + buffer), for the kinetic runs in the 6-12 pH region. The final proportion of organic solvent in the cell (1%) is considered negligible.

Type 2 experiments (from L' -). An approximately 5 . 10 -5 M aqueous solution of NaL' was mixed with an equivalent volume of either HCl solutions or with buffer solutions. The concentration of the starting L' -solution for type 2 experiments cannot be adjusted by weight, because the starting material NaL' . xH2O has a variable amount of water. This does not represent a problem for the determination of the rate constants since all transformations follow first order rate expressions. Knowledge of the exact concentration, however, is necessary to analyze the equilibrium spectra. This knowledge was obtained as follows. Each stock solution of NaL' . xH2O was used to record a spectrum in a pH ca. 7 buffer, where this remains the only species at equilibrium (see Results). This spectrum was compared with that obtained from a type 1 experiment in a pH 6.75 buffer, where a solution of HL at a known concentration transforms quantitatively to L' -.

(b) By mixing and monitoring. At pH ≤ 2, the evolution of the system is very slow (k < ca. 5 . 10 -4 s -1 ). To avoid possible diffusion problems, the kinetics at low pH were monitored by classical spectrophotometry in 1 cm cells at 25°C. The NaL' . xH2O and HL stock solutions were prepared as described above and mixed with aqueous HCl to obtain the desired pH.

(c) Data analysis. All kinetic runs were analyzed with the program Specfit, [Binstead, 2000 #4869] while the fitting of the kinetic model to the observed first order rate constants as a function of pH (see Results) was carried out with the program CurveFit. [Raner, 1992[Raner, #2239] ] Preparation and characterization of HL'. NaL' . xH2O (0.050 g) was dissolved in aqueous hydrochloric acid (0.01 M, 250 mL). The resulting red-violet solution (max = 520 nm) was evaporated to dryness in a vacuum freeze dryer. The residue was thorougly washed with water and dried in air. A microcrystalline dark brown powder was obtained (0.030 g), which was analyzed as HL' . 1. 3H2O. Calcd. C,52.37;H,2.31;N,37.49. Found: C,51.53;H,2.20;N,38.23. IR (KBr pellets 0.75% w/w and nujol mulls): max 2224 (CN), 1654, 1526, 1378 cm -1 (all intense bands). Visible spectrum (THF): max 520, 580 (sh) nm (freshly prepared solution).

Computational details. All calculations were carried out by density functional theory (DFT) methods with the Gaussian98 package, [Frisch, 1998 #5021] using the B3LYP functional [Becke, 1993[Becke, #2918] ] and the LANL2DZ basis set. The starting geometries for the optimization of HL and L' -were taken from the corresponding X-ray structures, [Fares, 1995 #4905;Fares, 1995 #4914] while those of L -and HL' were adapted from these by removal and addition of a proton. The dielectric effect of the aqueous solvent was included with use of the Polarizable Continuum Medium (PCM) method [Tomasi, 1994 #5022] of Tomasi and coworkers in the Integral Equation Formalism (IEF) version using the cavity built with the United Atom Topological Model. [Barone, 1997 #5023] This routine is included in the Gaussian package. The energies reported are the total solute free energies including the nonelectrostatic contribution due to cavitation, dispersion, and repulsion energies.

Results

(a) Spectroscopic and synthetic studies

Compound HL is stable only in the solid state or in anhydrous organic solvents. Its visible absorption spectrum shows a broad absorption band centered at ca. 580 nm as already described in previous studies. [Fares, 1995 #4905] The addition of water or buffered aqueous solutions at pH > 5 to organic solutions of HL provokes a rapid spectral evolution indicating the disappearance of HL and the formation of L' -, which is characterized by a stronger band centered at ca. 560 nm. The visible spectrum of this species has also been described in previous studies. [Bonamico, 1998 #4916] The addition of acidic solutions, on the other hand, leads to a more complex evolution. In addition, at low pH (< 2) an irreversible decomposition with a pH dependent rate takes place to afford a species which is characterized by an absorption band centered at 490 nm. The kinetics of these isomerization and decomposition processes have been examined (type 1 experiments) and will be presented in a subsequent section.

Compound L' -has been used in the present study as the sodium salt, which is soluble in water and stable in the pH range 6-14. Acidification of this solution below pH 6 produces an instantaneous change (within the time of mixing in the stopped-flow cell) of the visible spectrum, yielding a new band which is centered at 525 nm (at pH < ca. 2). Following this acidification process, the spectrum further changes slowly to yield a final equilibrium situation which involves the formation of HL and which is a function of pH. The kinetics of this equilibration process has been investigated by the type 2 experiments (vide infra).

The speed of the protonation process of L' -is consistent with the new species at 525 nm being the protonated form of L' -, HL', without molecular rearrangement (e.g. to a closed isomer). This transformation is fully reversible: raising the pH above 6 restores the spectrum of L' -. At pH < 2 a subsequent irreversible change occurs, due to the decomposition of HL mentioned above.

However, since the isomerization of HL' to HL becomes slower at lower pH values (see Results section), solutions of HL' at very low pH are quite stable, more so than solutions of its isomer HL at the same pH, giving a preliminary indication that the decomposition of HL' may indeed proceed via an initial isomerization to HL.

The initial spectra (ca. 1 ms) during type 2 experiments as a function of pH are shown in Figure 1. A Specfit analysis of these spectra yields the acidity constant of HL' (pK'a = 3.63±0.01).

<Figure 1>

Unfortunately, it has not been possible to fully characterize compound HL', because of its instability in solution. Even in anhydrous organic solvents, e.g. THF, it isomerizes to HL in a few hours. On the basis of the IR spectrum an open structure for HL' similar to that of L' -, as expected (1 or 2, see Computational Studies in section (d) below), can be proposed. As a matter of fact, the IR band pattern of HL' (2224, 1654, 1526, 1378 cm -1 ), see Figure 2, is very close to that of NaL ' (2220, 1640, 1534, 1348 cm -1 ) and significantly different from that of HL, the latter consisting of a greater number of bands in the same region (2235, 1667, 1598, 1517, 1485, 1378, 1266 cm -1 ). The higher energy portion of the spectrum, where the symmetric N-H stretching band occurs, is not similarly informative for the purposes of this comparison, because both energy and shape of this band are affected by the hydrogen bonds in the solid state. [Fares, 1995 #4914] <Figure 2>

(b) Ring opening/closing kinetics: from HL (type 1)

These experiments were carried out starting from a pure solution of HL (see details in the Experimental section). Since the decomposition of HL is very slow (ca. 10 -4 s -1 ) at pH 1, the starting solutions of HL were prepared by diluting stock solutions of the compound in MeCN with 0.1 M HCl. The sparing solubility of HL in neat water prevented us to carry out our studies under pure aqueous conditions. b1. From pH 5 to pH 13. These experiments involve the quantitative opening reaction of HL, giving L' -. At pH < ca. 8, a direct transformation of HL to L' -without the observation of intermediates takes place. The band centered at 580 nm for HL is replaced by the 560 nm band of L' -by a pH dependent first order rate constant. At pH > 8, however, the initial spectra are characterized by the presence of a new band centered at 660 nm, which subsequently disappears in a first-order decay process to yield L' -. Figure 3 shows the spectral evolution in a typical run at pH 9.20. At high pH values (> 9), the rate constant of this process becomes pH independent.

These transformations are irreversible, because the treatment of L' -under low pH conditions does not afford pure HL but rather an equilibrium mixture of different species whose ratio depends on the pH (see type 2 kinetic runs below).

<Figure 3> 10

The first order rate constants for the isomerization processes (kiobs) at the various pH are collected in Table 1 and are shown graphically in Figure 4. The above observations are in agreement with (i) the existence of form L -(max ca. 660 nm) which forms instantaneously by deprotonation of HL at high pH and (ii) quantitative, pH independent opening of L -to L' -. At pH > 8, the opening reaction is so fast that the first recorded spectra already show the presence of a large amount of L' -(e.g. see Figure 3) in addition to L -and HL. It is to be remarked that the max of the band assigned to L -is close to that observed for all the ML2 derivatives, where the closed form of the L -ligand has been verified by X-ray crystallography. [Bonamico, 1989 #4915;Bonamico, 1991 #4906;Bonamico, 1992 #4907;Bonamico, 1993 #4908;Bonamico, 1993 #4909] <Table 1 and Figure 4> b2. At pH between 2 and 5. Under these pH conditions, the opening reaction is not quantitative, but rather leaves a certain pH-dependent fraction of HL at equilibrium. In addition, the final form of the opened isomer of the ligand also depends on pH, a mixture of L' -and HL' being obtained in this pH range according to Figure 1. The extrapolated equilibrium spectrum is correctly analyzed by Specfit as a mixture of the initial closed form HL and the mix of the opened species HL' and L' -. The observed isomerization rate constants obtained by this method (kiobs in Table 1) correspond to those of the type 2 experiments in the same pH interval (vide infra). The latter, however, were determined more accurately and a more complete equilibration study has therefore been carried out starting from L' -(see section (c)). b3. At lower pH. As mentioned above, compound HL decomposes at low pH values. The decomposition is a pH dependent first order process, the rate constants (kd) at pH < 0.5 being directly obtained from clean first order decay processes that are not affected by the isomerization to the open isomer HL' (see Table 1). Evidently, the isomerization of HL is too slow under these conditions to effectively compete with the decomposition process.

(c) Ring opening/closing kinetics: from L' -(type 2)

These experiments have been carried out between pH 0.60 and 4.97 and involve the reversible closing reaction of L' -. At pH > 4.97 the extent of HL formation at equilibrium is not sufficient to allow the extraction of sufficiently precise rate constants. At pH ≥ 2, the equilibration can be handled as a pure 1st order process, yielding the kiobs values reported in Table 1. The knowledge of the spectrum of the pure closed form allows a kinetic data analysis in terms of the forward (closing of HL'/L' -, ki+) and reverse (opening of HL, ki-) components (ki+ + ki-= kiobs) . These values are also collected in Table 1. A typical equilibrium spectrum is shown in

K i = k i+ k i- = [L ] eq [L ' ] eq [HL] + [L ] eq - eq = [HL'] + [L' ] eq - eq
(1)

(d) Computational studies: structure of HL' and L -.

As species HL' and L -(in a free state, i.e. not coordinated to a transition metal) have not been previously described and their instability prevented us from obtaining detailed structural information, we have carried out computational studies on these species. In addition, we have also investigated the well known HL and L' -congeners in order to probe the validity of the computational method and to verify the energetics of the isomerization process on both neutral and ionic forms. The common numbering scheme used for the open and closed isomers is illustrated in Chart 2. The energetic results are summarized in Figure 7 and the relevant optimized geometric parameters are collected in Table 2.

<Chart 2, Figure 7 and Table 2> The structures of forms L' -and HL were optimized starting from the experimental structures as determined by X-ray crystallography. The structure of the closed anionic form L -has been calculated by removing an exo proton from the amino group of HL. In this way, the remaining endo proton may establish a stabilizing intramolecular H-bonding interaction with the lone pair of the imino group on the other 5-member ring. The alternative isomers where the two imino protons are either both endo or both exo were considered less likely. The endo/endo isomer was not calculated. The exo/exo isomer, on the other hand, was calculated for two reasons. First, this conformation is experimentally found in all transition metal complexes of L -. Second, we wanted to verify the stabilizing effect of the intramolecular H bonding interaction and evaluate its importance. The exo/exo isomer was indeed found higher in energy as expected, the difference relative to the endo/exo isomer being 6.4 kcal/mol for the isolated ions.

Concerning the open neutral species HL', this can be generated in principle by addition of a proton to L' -at either the pyrrolyl nitrogen atom to yield the 2-tricyanovinyl-3,4-dicyano-5amino-1H-pyrrole form (1), at the nitrogen atom of the amino substituent to yield a zwitterionic 2-(5-ammonio-3,4-dicyano-2H-pyrrol-2-ylidene)-1,1,2-tricyanoethanide form (2), or at the  carbon atom of the tricyanovinyl group. Discarding the latter possibility which would disrupt the vinyl conjugation with the pyrrole ring, the other two possibilities were both considered and the corresponding geometries optimized. As a result, form 2 turned out to be 34 kcal/mole less stable than form 1, suggesting that the structure of HL' is as shown in 1. It can be noted that, besides the absence of a charge separation, structure 1 also benefits from an extended  delocalization on the amino substituent. An isomeric structure for the anion L' -was also considered, namely a tautomer having an imino group at the C2 position and a proton on N1. However, besides not corresponding to the crystallographically determined structure, a much higher energy (45 kcal/mol) was calculated for this optimized structure.

As shown in Figure 7, the calculations in vacuum (energy values in parentheses) yield a more stable closed isomer for the acid by 11.1 kcal mol -1 and also a more stable closed isomer for the conjugate base by only 0.3 kcal mol -1 . Both results slightly overestimate the relative stability of the closed form, since the observed 4:1 equilibrium between HL and HL' corresponds to a ∆G = 0.8 kcal/mol while the quantitative transformation of L -to L' -(equilibrium ratio < 1:100) implies that the latter is at least 2.7 kcal/mol lower in free energy. No significant entropic contribution is expected in these isomerization processes, thus ∆G ≈ ∆H ≈ ∆E. We considered that the neglect of solvation energy could be partly responsible for this discrepancy, thus we repeated the calculations with the solvent as a polarizable continuum model. [Tomasi, 1994 #5022] Although the situation is improved for the neutral species, it becomes worse for the anionic one. The neglect of explicit hydrogen bonding interactions between the organic molecules and water could account for the remaining discrepancy. Indeed, the closed forms L - and HL are characterized by a greater extent of intramolecular H bonding and consequently should be stabilized to a lesser extent by the intermolecular H bonding with the solvent. On the other hand, it is also possible that the H bonding interactions with water are more favourable than the intramolecular H bonding interactions. Thus, a more representative form for L -in water would be the exo/exo conformer which has indeed a higher energy than the L' -isomer (6.1 kcal/mole from the gas-phase calculation and only 0.1 kcal/mol from the PCM calculation). The explicit neglect of the H bonding interactions with water could account for any residual discrepancy. Nevertheless, the disagreement between theory and experiment remains within the few kcal mol -1 that are usually considered as customary for the density functional methodologies.

An analysis of the geometries in Table 2 confirms the suitability of the chosen computational method. The optimized geometries for the known compounds L' -and HL match pretty closely the experimentally determined structures. Certain calculated distances are closer to experiment for the structures optimized in vacuum, others for the structure optimized in water.

As the experimental structures were obtained in the crystalline state, neither of the calculation can in fact be directly compared with the experiment. The largest discrepancy is observed in the C-N distances of the cyano groups, for which the vacuum and water calculations show values that are close to each other and systematically longer than 1.18 Å, while the corresponding experimental values are 0.04-0.07 Å shorter. Otherwise, the agreement between calculated and experimental values is excellent, the calculated values being in general a bit longer with only few exceptions.

Discussion (a) Acid/base equilibria and isomerization rates

We shall first examine the acid/base and isomerization behaviour of the HL/L' -system without consideration of the acid decomposition process. The kinetic and spectroscopic studies have revealed the existence of the basic form of the closed species, L -, and the acidic form of the open species, HL'. The protonation/deprotonation processes within each isomer are, as expected, extremely fast while the isomerization is rate limiting. The kinetic model is therefore based on the square Scheme 1, where the proton transfer (vertical) processes are supposed to be instantaneously achieved and at equilibrium at all times.

<Scheme 1>

Given expression ( 1), the kinetic model can be treated as a simple first order equilibration between species L and L', according to Scheme 2, justifying the observation of isosbestic points for all kinetic runs (except those involving further decomposition at low pH values, see Results section). The values of ki+ and kiat different pH are available experimentally (Table 1) as illustrated in the Results section. Under the conditions in which ki->> ki+ (pH > 6) a quantitative opening of L is observed.

<Scheme 2>

It can be easily shown (annex I) that kiand ki+ are related to the kinetic and equilibrium parameters of Scheme 1 according to equations 2 and 3, respectively.

k i-= k 0-+ k 1-[H + ]/K a 1 + [H + ]/K a
(2)

k i+ = k 0+ + k 1+ [H + ]/K' a 1 + [H + ]/K' a (3)
The opening rate constant (ki-, equation 2) can be analyzed in the entire pH interval, spanning a rate range of ca. 7 orders of magnitude, although this becomes a quantitative process (ki+ << ki- kiobs) only at pH > ca. 5 (see Figure 4). The shape of the curve passing through all the experimental points indicates that at a high pH, both [H + ]-dependent terms in equation 2 become negligible and the limiting value of kiobs = kiis therefore equal to k0-(the ring opening rate for L -). As the pH is lowered, the function enters a wide pH interval (between ca. 2 and 8) where log kiis linearly dependent on pH with slope 1. This signifies that while the [H + ]/Ka term becomes important in the denominator, the numerator term k1-[H + ]/Ka remains negligible relative to k0-.

Thus, equation 2 can be simplified to equation 4 down to ca. pH 2. This means that, in that region, compound HL isomerizes mainly via its conjugate base L -, following a fast deprotonation pre-equilibrium.

k i-= k 0- 1 + [H + ]/K' a (4) (pH > 2)
The break point between the two linear regions of slopes 1 and 0 affords the value of the pKa, which is 8.61±0.03 by fitting of equation 4 to the experimental data. This fitting also affords log k0-= 2.58±0.03 (k0-= 3.8(3) . 10 2 s -1 ). A curvature leading to a second break point for kiis visible in Figure 4 at low pH. This is the region in which the numerator term k1-[H + ]/Ka of equation 2 starts to become non negligible. A global fitting of equation 2 to the kidata yields log k1-= -5.37±0.12 (k1-= 4.3(1.2) . 10 -6 s -1 ), the rate of ring opening of the acidic form HL.

The closing reaction (ki+, equation 3) can only be analyzed in the pH interval between 0.6 and 5. An S-shaped behaviour similar to that of kiis shown by the data (Figure 4). Thus, the high pH region gives rise to a constant value for ki+ = k0+ = 1.62(7) . 10 -2 s -1 (log k0+ = -1.79±0.02). Down to pH 2 the closing reaction of L' -takes place mainly via closing of the anionic form (k0+) to afford L -, followed by instantaneous protonation. The protonation equilibrium of L' -(K'a) slows the reaction by sequestering starting material into the less reactive acidic form. The break point of the curve corresponds to the pK'a, which is calculated as 3.62±0.02 by fitting of the ki+ experimental data. This value is identical within experimental error to that determined as shown in section (a) of Results and Figure 1, (log K'a = 3.63±0.01), further indicating the internal consistency and correctness of the model. Fitting equation 3 with all the ki+ data affords the value of k1+ = 1.8(2) . 10 -5 s -1 (log k1+ = -4.75±0.06).

The knowledge of all rate and ionization constants allows us to analyze the interconversion between the closed and open forms L' and L in terms of the two possible pathways of the square Scheme 1 (kiobs = k0obs + k1obs, where k0obs and k1obs are given by eq. 5 and 6). As shown in Figure 8, the isomerization takes place preferentially via the anionic form at pH > 0.67.

k 0obs = k 0+ 1 + [H + ]/K' a (5) k 0- 1 + [H + ]/K a + k 1obs = k 1+ K' a /[H + ] + 1 (6) + k 1- K a /[H + ] + 1 <Figure 8>
All the determined rate constants also allow us to calculate equilibrium constants for the isomerization process in the basic, anionic state (log K0 = log(k0+/k0-) = -4.37±0.04) and in the acidic, neutral state (log K1 = log(k1+/k1-) = 0.62±0.13). It is important to notice that the low pH breakpoint value of the ki+ curve (at pH = 0.66±0.09) corresponds within experimental error to the low pH breakpoint value of the kicurve (at pH = 0.66±0.14), as demanded by the characteristics of the square scheme (K'aK0 = KaK1). This is yet another check that our kinetic and equilibrium data are internally consistent. The ratio ki+/ki-= Ki as defined in equation 1, namely the ratio between closed and open isomers in all their forms, is also represented in Compound HL can be isolated as a pure material from acidic water solutions only thanks to its lower solubility relative to HL'. Once HL is redissolved in acidic solutions, it is kinetically rather stable. On the other hand, pure solutions of HL' can only be kinetically generated upon acidification of solutions of L' -. Finally, L -can only be observed as a fleeting intermediate and rapidly isomerizes to the thermodynamically more stable L' -, although it is stabilized by coordination to transition metals as shown previously. [Bonamico, 1989 #4915;Bonamico, 1991 #4906;Bonamico, 1992 #4907;Bonamico, 1993 #4908;Bonamico, 1993 #4909]

(b) Decomposition of HL

The kinetic measurements show that compound HL decomposes under strongly acidic conditions (pH < 2), while HL' decomposes only via its preliminary isomerization to HL. Since the isomerization of HL' to HL (k1+) is extremely slow, solutions of HL' are quite stable under strongly acidic conditions. As shown in Figure 4, the decomposition of HL follows a rate law which is first order in [H + ]. This is consistent with a rapid protonation equilibrium, followed by rate-determining decomposition of the protonated species H2L + (Scheme 3). No deviation from linearity for the log kd vs. pH data is apparent from Figure 4, indicating that the pKa value for the H2L + species is lower than -1.

<Scheme 3>

Conclusions

All the kinetic and thermodynamic parameters concerning the acid-base and isomerization equilibria involving HL and L' -have been determined. They are based on a model whose consistency has been checked on several points. We would like to make reference to other studies where kinetic models similar to that used here have allowed sorting the acid dissociation constants and the rates of reversible chemical equilibria. [Strehlow, 1962 #5002;Chiang, 1988 #4994;Chiang, 1992 #4995;Legouin, 1999 #4999] We are not aware, however, of precedents for the complete determination of all forward and reverse chemical rate constants and acid dissociation constants for a square scheme such as that illustrated in the present study.

Further progress in the chemistry of HL and/or L' -is now possible taking advantage of the information gathered with the present study. On the other hand, the successful applications of these materials ultimately depends on the feasibility of their chemical derivatization. As a matter of fact, after the initial attempts to exploit HL, L' -or ML in sensor devices, as mentioned above,

we are now considering the modification of these species through chemical substitution for more advanced applications such as, for instance, monolayers covalently bound to the surface of crystalline silicon, or LB films. The information gained on the HL/L' -system with the current work is highly valuable for further development along these directions.

Annex I

Given relation (1) and the pKa relations,

[L] = [L -] (7) [L'] = [L' -]' (8)
where

 = {1+[H + ]/Ka} and ' = {1+[H + ]/K'a} For the equilibration of Scheme 1, v = {k0+[L' -]+k1+[HL']}-{k0-[L -]+k1-[HL]} = {k0++k1+[H + ]/K'a}[L' -]-{k0-+k1-[H + ]/Ka}[L -] = '[L' -]-[L -] (9)
where  = k0-+k1-[H + ]/Ka and ' = k0++k1+[H + ]/K'a By using ( 7) and ( 8)

, v = ('/')[L']-(/)[L]
Comparison with Scheme 2 gives ki+ = '/' and ki-= /, i.e. equations ( 2) and (3). ref. [Fares, 1995 #4905].

Captions for Figures

  Figure 5.

  Figure 4 as a function of pH. It is worth noting that the ki+ and kivalues obtained from the two different analyses of type 2 experiments show continuity as a function of pH. The values of kiobs obtained from type 1 and type 2 experiments, including those calculated as (ki+ + ki-) at pH < 2, are equally placed on a continuum curve. The kiobs in the pH > 5 region (where the opening reaction is quantitative) are nicely aligned with the kidata of the pH < 5 region. Finally, the values of kd from the two different kinetic analyses of type 1 and 2experiments are again aligned on a continuum curve (straight line). Thus, all data from all types of experiments are internally consistent. Figure4also shows the ratio of opening and closing rate constants, Ki = ki+/ki-, representing the ratio between the global concentration of the closed and opened isomers in all of their forms as a function of pH, see equation 1. The anionic form of the closed species, L -, is never present at equilibrium in any appreciable concentration.

  Figure 4. The theoretical curve fitting these experimental points is calculated from all acidity and rate constants in Scheme 1. According to our model, Ki becomes the pH-independent isomerization equilibrium constant for the acidic forms, K1, at low pH whereas it converges to the pH-independent isomerization equilibrium constant of the basic forms, K0, at high pH. These two horizontal asymptotes are shown by the two dashed lines in Figure 4. The individual fractional distribution at equilibrium (cf. annex II) of the observable species [HL], [HL'] and [L' -] as a function of pH is shown in Figure 9. <Figure 9> From Figure 9 it can be observed that [L' -] will be essentially the only species present in solution at pH > 6, whereas the maximum equilibrium concentration of [HL] (79 %) is attained at pH ≤ 2. The species distribution of Figure 9 is valid, of course, only in the absence of HL decomposition.

  Figure 1. Spectra of solutions obtained from HL and buffers at different pH at zero time.

Figure 2 .

 2 Figure 2. IR spectra in Nujol mull of compounds HL', HL . 1/2NAPH (NAPH = 1-

Figure 3 .

 3 Figure 3. Representative spectra of the time evolution in a type 1 experiment at pH 9.20. Time

Figure 4 .

 4 Figure 4. Plot of log k from Table1(rate constants in s -1 and equilibrium constants) vs pH or

Figure 5 .

 5 Figure 5. Initial and equilibrium spectra for a type 2 run at pH 3.37. The spectrum of pure HL

Figure 6 .

 6 Figure 6. Representative spectra for a type 2 run at pH 1.30. A: starting spectrum; B:

Figure 7 .

 7 Figure 7. Energetics for the isomerization process of HL/HL' and L -/L' -according to the

Figure 8 .

 8 Figure 8. Fractional distribution of the isomerization pathways between open and closed isomers

Figure 9 .

 9 Figure 9. Equilibrium fractional distribution of species HL, HL' and L' -as a function of pH.

  Figure 1

  Table1). At lower pH, the isomerization rate is too slow for significant rate data to be determined. The results of a typical analysis (pH 1.30) are shown in Figure6. It should be remarked that the spectrum of the final decomposition product (species C) is computed by Specfit and corresponds within experimental error to that obtained from the direct decomposition of HL from the type 1 experiments.

<Figure 6>

All values of kiobs, ki+, ki-, and kd from type 1 and type 2 experiments in Table

1

are graphically shown in

Table 2 .

 2 Relevant B3LYP/LANL2DZ optimized a structural parameters for L' -, L -, HL', and HL, and comparison with experiment Values on the left and right refer to optimization in vacuum and water medium, respectively. b From ref.[Fares, 1995 #4914].

	(distances in Å; angles in degrees)
	Parameter	L' -		L -	HL'	HL
		Optim.	Exper. b	Optim.	Optim.	Optim.	Exper. c
	N1-C2	1.345/1.355	1.341(6)	1.423/1.449	1.368/1.370	1.375/1.379	1.400(6)
	N1-C5	1.395/1.393	1.389(6)	1.401/1.410	1.416/1.408	1.402/1.402	1.364(7)
	N1-C8	-	-	1.409/1.443	-	1.418/1.425	1.412(6)
	C2-N2	1.379/1.363	1.349(7)	1.301/1.298	1.361/1.342	1.349/1.339	1.305(7)
	C2-C3	1.449/1.455	1.420(7)	1.482/1.469	1.422/1.434	1.429/1.436	1.393(7)
	C3-C4	1.426/1.418	1.395(7)	1.418/1.409	1.429/1.420	1.447/1.441	1.429(7)
	C3-C9	1.413/1.410	1.415(8)	1.414/1.409	1.417/1.412	1.416/1.410	1.434(7)
	C4-C5	1.445/1.445	1.414(7)	1.424/1.414	1.420/1.424	1.405/1.400	1.373(7)
	C4-C10	1.420/1.416	1.446(7)	1.425/1.419	1.424/1.421	1.422/1.420	1.432(7)
	C5-C6	1.427/1.419	1.401(7)	1.413/1.409	1.435/1.423	1.437/1.429	1.431(7)
	C6-C7	1.411/1.414	1.400(7)	1.432/1.414	1.405/1.407	1.406/1.401	1.364(9)
	C6-C11	1.444/1.441	1.420(7)	1.423/1.418	1.442/1.442	1.425/1.424	1.429(9)
	C7-C8	1.432/1.425	1.429(7)	1.479/1.472	1.429/1.423	1.489/1.491	1.470(8)
	C7-C12	1.433/1.426	1.427(8)	1.410/1.406	1.431/1.427	1.415/1.411	1.404(7)
	C8-N8	1.186/1.186	1.141(7)	1.313/1.303	1.186/1.185	1.296/1.291	1.268(8)
	C9-N9	1.189/1.187	1.138(9)	1.187/1.187	1.184/1.184	1.185/1.185	1.115(7)
	C10-N10	1.186/1.185	1.128(8)	1.183/1.181	1.182/1.182	1.182/1.182	1.120(8)
	C11-N11	1.182/1.181	1.155(8)	1.184/1.182	1.181/1.180	1.181/1.181	1.129(10)
	C12-N12	1.187/1.186	1.149(8)	1.189/1.188	1.183/1.184	1.185/1.184	1.143(8)

a c From

Acknowledgements. RP and EC are grateful to the Conseil Régional de Bourgogne for funding that has allowed the purchase of the stopped-flow apparatus.

Annex II

The total concentration of all species at equilibrium is CT = [L]eq+[L ']eq = [L -]eq+'[L' -]eq (10) From equation ( 9) (v = 0 at equilibrium):

From the combination of eqs. ( 10) and ( 11) and use of the acidity constants: [Paul, 1957 #5205] f Too slow to be determined accurately.


Erwan Le Grognec

Rinaldo Poli
email: rinaldo.poli@u-bourgogne.fr

Philippe Richard

Rosa Llusar

Santiago Uriel

Interaction of Half-Sandwich Alkylmolybdenum(III) Complexes with B(C6F5)3. The X-ray Structure of [CpMo( 4 -C4H6)(-Cl)(-CH2)(O)MoCp][CH3B(C6F5)3]

    Interaction of half-sandwich alkylmolybdenum(III) complexes with B(C6F5)3. The X-ray structure of [CpMo(η4-C4H6)(µ-Cl)(µ-CH2)(O)MoCp][CH3B(C6F5)3]

Introduction

Polymerization of -olefins is one of the most important industrial processes. The combination of group 4 metallocenes and methylaluminoxane (MAO) has been actively investigated in view of its homogeneous process and the possibility to control the stereoregularity of the polymerization though the control of the ligand architecture. [1][2][3][4][5] A few studies have also been carried out with isoelectronic group 5 metal complexes such as CpNb( 4 -diene)R2 6,7 and Tp'Nb( 2 -alkyne)R2 (Tp' = substituted hydrotris(pyrazolyl)borate). 8 The prevalent concept of these works is the analogy (structure and electron count) to the architecture of the group 4 metallocenes, where the 14 electron metallocene alkyl cation [Cp2MR] + is the proposed active species. 1,3,4 By analogy with the work done on the group 5 metals, we considered the possibility that the half-sandwich 17-electrons molybdenum(III) complexes available in our laboratory 9,10 could also catalyze ethylene polymerization. Our initial experimental and theoretical studies of these complexes in ethylene polymerization catalysis have been recently reported. 11 In combination with methylaluminoxane, very low activities has been observed to yield polymers of high linearity and very high molecular weights. The theoretical studies have identified the low polarity of the Mo-CH3 bonds, as compared with that of the Nb-CH3 bonds in the analogous Nb(III) derivatives, as the cause of the low activity because the activation process is made less favorable. Other strong organo-Lewis acids such as B(C6F5)3 have also been used to generate the active species in situ. 12 Under the assumption that the active specie is the [CpMo( 4 -diene)R] + complex , we were interested in probing the reactivity of the dimethyl complexes CpMo( 4 diene)Me2 and Cp* analogues with the methyl abstracting agent B(C6F5)3 in order to evaluate the possibility to isolate the putative active species. The results of these investigations are reported in this contribution.

Results and Discussion

The addition of one equivalent of B(C6F5)3 to solutions of CpMo( 4 -diene)Me2 (diene = 1,3-butadiene or 2,3-dimethyl-1,3-butadiene) in toluene at -20°C leads to a rapid color change from green to orange. Analogous color changes were also observed upon treating the same Mo complexes with MAO. 11 In each case, the product appears stable at room temperature in hydrocarbon solvents and under an inert atmosphere, at least for a few days. Unfortunately, all our crystallization attempts (see Experimental) have yielded oily products. Therefore, the only information available to us on the nature of the reaction products comes from solution spectroscopic studies. Monitoring the reactions by EPR at -20°C shows that, in each case, the characteristic pseudo-binomial 11-line pattern of the dimethyl Mo(III) complex, 10 due to the apparent equivalence of 10 hydrogen atoms of the molecule, disappears and is replaced by a new, broader signal at g = 1.995. Both starting and final EPR spectra are essentially independent on the nature of the diene ligand, since the nature of the substitution of the diene ligand does not influence the g value nor the multiplicity of the spectrum. Figure 1 shows the example of the butadiene system. The new signal is slightly broader than that of the starting complex and is characterized by a different coupling pattern.

<Figure 1>

Several adducts of B(C6F5)3 with transition metal complexes have been characterized spectroscopically and, in a few cases, also by X-ray crystallography. [13][14][15] The borane reagent unfailingly attacks an alkyl group whenever at least one is present, to generate a M

•••R•••B(C6F5)3
moiety where the M-R bonding interaction is greatly weakened. In the presence of Lewis bases L (i.e., THF, pyridine, phosphines), sometimes even very weak ones such as organic halides, the alkyl is typically abstracted resulting in charge separation, to generate [M(L)] + [R-B(C6F5)3] - species. [15][16][17][18] However, boranes have also been shown to give rise to electrophilic additions to the terminal carbon atom of a coordinated diene to yield complexes containing a betaine allylic ligand, e.g. M[ 3 -C3H4-CH2B(C6F5)3] starting from a M( 4 -C4H6) complex. [19][20][21][22][23][24] Therefore, the question of the site of attack (methyl or diene ligand) for complexes CpMo( 4 -C4H6)Me2 is not a trivial one a priori. A good spectroscopic handle on the nature of the interaction is provided by 11 B and 19 F NMR, but this is unfortunately prevented in our case by the paramagnetism. The inability to observe a 19 F NMR signal suggests, in fact, that the borane molecule is covalently linked to the paramagnetic Mo(III) complex.

The digital fitting of the EPR experimental spectrum provides useful information on the site of attack. The fitting was carried out only for the butadiene system, for which a slightly improved spectral resolution allowed a more accurate analysis. The fitting gives the best agreement for an A3M3R2X2 spin system (see Experimental section). This is interpreted as resulting from two sets of three equivalent H atoms from the two chemically inequivalent methyl groups, plus two sets of two pseudo-equivalent H atoms from the two chemically distinct diene terminal (syn and anti) positions. Previous EPR studies of paramagnetic diene complexes have systematically shown coupling to the terminal syn and anti H atoms and no observable coupling to the internal H atoms. 9,10,[25][26][27] Thus, the observed spectral modifications suggest that an interaction has occurred between the Lewis acid B(C6F5)3 and one methyl group. The nonequivalence of the two methyl groups would then arise from the fact that one is bonded only to the Mo atom whereas the other one is bonded to the Mo and B atoms. This asymmetry should make the two opposite termini of the diene ligand also nonequivalent, giving rise to a RSXY system rather than a R2X2 system for the diene H atoms. However, these atoms are far removed from the source of the asymmetry, yielding pseudo-equivalence, at least within the resolution limits of our spectrometer. Indeed, the quality of the fit significantly increases upon breaking the symmetry from 64 to 334 (the correlation increases from 0.990 to 0.992 and the error decreases from 1.743 . 10 -2 to 1.342 . 10 -2 ). An almost insignificant improvement takes place upon further breaking the symmetry to 3322, because of the accidental equivalence of aH to syn and anti diene protons. Previous studies have also highlighted the near-degeneracy of the hyperfine couplings to these two types of atoms. 10,26,27 Further lowering the symmetry to a 331111 spin system does not further improve the quality of the fit. The alternative interpretation of a borane attack at one of the diene termini seems less likely for two reasons. First, the diene terminal H spin system would be expected to show greater asymmetry. Second, and more importantly, the resulting betaine allyl ligand would show coupling to the allylic terminal H atoms (thus 3 rather than 4 H atoms), according to the available precedents for similar allylmolybdenum(III) complexes. 27,28 Attempts to simulate the EPR spectrum on the basis of a 33111 model, however, systematically gave unacceptable fits. It is to be noted that the aH values for the two inequivalent methyl groups (6.79 and 4.42 G) are not too different from each other, suggesting that the borane-bonded methyl group is still solidly attached to the Mo center. This is consistent with the observed reactivity (or lack thereof), vide infra. For comparison, the aH value for the methyl H atoms in dimethylmolybdenym(III) complexes are in the 5.3-6.3 G range, depending on the nature of the diene and cyclopentadienyl ligands. 10,29 The increased broadness of the spectrum relative to that of the starting dimethyl complex may be related to the onset of a dynamic exchange of the two methyl groups between the borane-bonded and nonbonded positions. A similar phenomenon has been shown to broaden the 1 H NMR resonances for diamagnetic compounds having the 13,30,31 The EPR spectrum recorded at different temperatures (both lower and higher) has a poorer resolution and no useful information could be obtained from its analysis.

M•••R•••B(C6F5)3 moiety.
Additional indication for the occurrence of a methyl rather than diene attack comes from previous computational investigations, showing that the Mulliken charge on the methyl ligands is more negative than that on the diene terminal position. 11 This is true also for the corresponding Nb system, for which the formation of the [CpNb( 4 -diene)Me] + [Me-MAO] -active site in ethylene polymerization catalysis has been proposed. 7,32 Finally, the borane attack at the methyl group for compound CpMo( 4 -C4H6)Me2 is indicated by the characterization of the title compound, which results from a decomposition process (vide infra).

When the addition of CpMo( 4 -C4H6)Me2 to the B(C6F5)3 solution was followed by the addition of a slight excess of PMe3, no spectral modifications were observed by EPR. This surprising observation shows that the phosphine molecule is incapable of displacing the methyl group from either the boron atom [giving back the starting dimethyl Mo(III) complex and Me3PB(C6F5)3] or the molybdenum atom. The latter process would afford [CpMo( 4 -C4H6)-(PMe3)Me][MeB(C6F5)3] by substitution of the labile [CH3B(C6F5)3] -ligand with PMe3 and should dramatically affect the shape of the EPR spectrum (notably a strong coupling to the P nucleus would be expected). A reaction of this type is undertaken, for instance, by the boratabenzene chromium complex (C5H5BMe)Cr(Me){MeB(C6F5)3} upon addition of PMe3. 18 The resistance of the putative CpMo( 4 -C4H6)Me{MeB(C6F5)3} product against displacement of the borate anion is in agreement with the strength and low polarity of the Mo-CH3 bonds in the half sandwich Mo(III) precursors, as indicated theoretically by DFT calculations 11 and experimentally by protonation studies. 10 Dissolution of the putative CpMo( 4 -C4H6)Me{MeB(C6F5)3} product in dichloromethane and diffusion of pentane, with participation of adventitious oxygen or water, afforded crystals of an oxygen and chlorine-containing decomposition product which were suitable for a full characterization, including an X-ray analysis. While the chlorine atom is obviously originating from the CH2Cl2 solvent, a likely source of the oxygen atom is residual water in the Celite used for the filtration. While oven-dried Celite was used and the product isolation was reproduced, omission of the filtration procedure only resulted in the formation of oily products. The compound corresponds to a salt of the dinuclear [CpMo( 4 -C4H6)(µ-Cl)(µ-CH2)(O)MoCp] + cation and the [B(C6F5)CH3] -anion. An ORTEP view of the two ions is shown in Figure 2 and selected metric parameters are given in Table 1. <Figure 2 and Table 1> Cation and anion are well separated from each other in the unit cell. The minimum contact is 2.368 Å Å between atoms H4 and H60b. Very weak interactions are observed between two symmetry-related cations (O1 … H2, 2.478 Å; O1 … H2-C2, 146.1º), which are therefore pairwise arranged in the unit cell as shown in Figure 3. The anion shows a regular tetrahedral coordination around the boron atom. The B-CH3 distance of 1.658(6) Å is quite typical for this anion. 33,34 The cation is composed of the two fragments CpMo( 4 -C4H6) and CpMo(O) bridged by a chlorine atom and a CH2 group. The hydrogen atoms of the bridging methylene group were not directly located by the X-ray analysis, but their presence is unambiguously shown by the NMR characterization (vide infra). The central Mo2(-Cl)(-CH2) moiety is planar and the two Mo fragments are oriented in such a way as to place the two Cp rings on the opposite sides of the bridging plane (anti). The two metal atoms are sufficiently close to each other [2.8801( 5 However, we should not assign, from this observation alone, a formal oxidation state IV to the Mo(1) center (and consequently an oxidation state IV also to the Mo(2) center). The alternative assignment as a III,V compound would also be consistent with the observed trends if we suppose that the presence of the adjacent Mo(V) center increases, via the bridging ligands and the direct metal-metal bond, the effective positive charge on the Mo(III) center. The Mo-X (bridging atom) distances to Mo(1) and Mo(2), however, are rather similar Cp*MoCl2(O), 35 1.668(7) Å for [CpMoOCl]2(-O), 36 or 1.676(5) Å for [Cp*MoOCl]2(-O), 37 than with analogous distances in cyclopentadienyl complexes of Mo(IV), such as 1.721(2) Å for (C5H4Me)2Mo(O). 38 As mentioned above, the structure of this decomposition product is indirect evidence that the borane reagent has attacked the methyl group rather than the diene ligand in the dimethylmolybdenum(III) precursor, because one methyl group is found to bind the boron atom in the anion, while the putative residual [CpMo( 4 -C4H6)(CH3)] + becomes incorporated, after deprotonation, into the cationic unit, the methyl ligand becoming the bridging methylene group. The high isolated yield (90%) excludes the fact that this complex is a by-product. The formation of the dinuclear cation can be ascribed to the instability of the putative 15-electron intermediate deriving from the methyl abstraction by B(C6F5)3. An analogous instability has also been observed for related group 5 complexes, e.g. [Cp*Ta( 4 -C4H6)Me] + and [Tp'Nb( 2 -PhCCMe)Me] + (Tp' = hydridotris(pyrazolyl)borate), 7,8,39 which, like our proposed Mo(III) intermediate, could not be isolated.

<Scheme 1>

Characterization of the title compound is completed by 1 H, 11 B and 19 F NMR spectroscopy and by mass spectrometry. The 1 H NMR spectrum (CDCl3) shows the presence of two different isomers in solution (in a 70:30 ratio by integration of the cyclopentadienyl signals).

These two isomers are interpreted as having a different relative position of the CpMo(=O) and CpMo( 4 -C4H6) moieties across the central Mo2( -Cl)(-CH2) moiety, i.e syn or anti (see Scheme 2) A similar behavior has been observed for complexes [CpMo(-X)( 4 -C4H6)]2 (X=Cl, Br). 27 All protons show distinct signals for each isomer, including the six hydrogen atoms of the diene ligand. The two inequivalent methylene protons give rise to two distinct doublets (JHH = 10.5 Hz) for the major isomer, while a singlet is observed for the minor one. The 11 B chemical shift and the difference of 2.5 ppm between the meta-F and para-F resonances in the 19 F NMR are in good agreement with the formation of the borate ion and with a significant charge separation in the salt. 16,40 The mass spectrum shows peaks in agreement with the given formulas for the cation and anion in the positive and negative chemical ionization modes, respectively.

<Scheme 2>

Conclusion

The borane reagent B(C6F5)3 attacks the electron-rich dimethyl complexes Cp*Mo( 4diene)Me2 at the methyl position, rather than at the diene ligand, to form relatively stable adducts that resist displacement of the [CH3B(C6F5)3] -anion by Lewis bases. The position of borane attack is suggested by the simulation of the EPR data and indirectly by the structural characterization of the title product. The observed reactivity pattern is consistent with previously available Mulliken charges from DFT computations.

The present study has confirmed the great resistance of the Mo(III)-CH3 bond toward heterolytic cleavage, previously shown by the fact that the bond can only be cleaved by acids stronger than H3O + . 10 Yet, compounds of type Cp*Mo( 4 -diene)Me2 has been shown capable to polymerize ethylene in combination with MAO, with a relatively rapid turnover for the active site. The activation step, however, is difficult. 11 The results shown here will guide us toward the development of modified systems with an improved catalytic activity.

Experimental section

General procedures. All reactions were carried out in a Jacomex glove box or by the use of standard Schlenk techniques under an argon atmosphere. The solvents were dried by conventional methods (THF, Et2O, toluene and pentane from sodium benzophenone ketyl and CH2Cl2 from P4O10) and distilled under argon prior to use. EPR measurements were carried out at the X-band microwave frequency on a Bruker ESP300 spectrometer. The spectrometer frequency was calibrated with DPPH (g= 2.0037). NMR spectra were recorded using a Bruker AC-200 or DRX-500. 1 H NMR spectra were referenced to the residual solvent protons of the deuterated solvent. 19 F (282.2 MHz) and 11 B (96.2 MHz) NMR spectra were referenced externally to CFCl3 and BF3 . Et2O, respectively. Cyclic voltammograms were recorded with an EG&G 362 potentiostat connected to a Macintosh computer through MacLab hardware/software. Electrospray mass spectra were recorded using a Micromass Quattro LC instrument. Isotope patterns were recorded for cation and anion and compared with theoretical patterns using the MassLynx 3.5 program. In both cases there was good agreement between the experimental and calculated isotopic mass distributions. The elemental analyses were carried out by the analytical service of the Laboratoire de Synthèse et d'Electrosynthèse Organométallique with a Fisons EA 1108 apparatus. Compound CpMo(diene)Cl2 were prepared as described in the literature. 10 B(C6F5)3 was prepared as described 41 and was sublimed in vacuo before use. 2. The structure was solved by direct methods and refined by the fullmatrix method based on F 2 using the SHELXTL software package. [START_REF] Sheldrick | SHELXTL, version 5.1[END_REF] The non-hydrogen atoms of structure was refined anisotropically; the positions of all hydrogen atoms were generated geometrically, assigned isotropic thermal parameters and allowed to ride on their respective parent carbon atoms. The chlorine and methylene bridge are disordered and were refined with a occupations factors of 0.62 for Cl1 and C2A and 0.38 for Cl2 and C1A. <Table 2>

Reaction of

C (40)

  ) Å] to suggest the presence of a direct metal-metal interaction. The acute Mo-X-Mo angles at the bridging Cl and C(2A) atoms [72.00(4) and 86.62(3)°, respectively] reinforce this proposition. <Figure 3> Relative to the previously reported structure of CpMo( 4 -C4H4Me2-2,3)(CH3)2, the Mo(1)-CNT(1) distance (Mo-Cp centroid) is slightly shorter [1.985(4) vs. 2.019(5) Å], while the Mo(1)-CNT(3) distance (Mo-diene centroid) is slightly longer [1.964(4) vs. 1.925(3) Å], and the Mo(1)-C(2A) distance is considerably shorter relative to the two Mo-CH3 distances [2.090(10) Å vs. an average of 2.225(3) Å]. A shortening of all distances to the (formally) anionic ligands and a lengthening of those to the neutral ones are expected when an oxidation state increase occurs.

  [2.497(2) vs. 2.401(3) Å for Cl, 2.090(10) vs. 2.108(15) Å for C(2A)], in better agreement with a IV,IV formulation. While the two Mo-Cl distances are significantly different with that to Mo(2) being shorter, the two Mo-C(2A) distances are identical by the 3 criterion. The Mo(2)-CNT(2) distance is longer than the Mo(1)-CNT(1) distance. Finally, the Mo(2)-O distance [1.673(3) Å] compares better with those of terminal Mo=O bonds in other cyclopentadienyl complexes of Mo(V), such as 1.684(4) Å for

  The [CpMo(O)Cl] + fragment that completes the structure of the cation must derive from the decomposition of a second molecule of [CpMo( 4 -C4H6)(CH3)][CH3B(C6F5)3] (see Scheme 1).

  CpMo( 4 -C4H6)(CH3)2 with B(C6F5)3. A solution of CpMo( 4 -C4H6)(CH3)2 (84 mg, 0,34 mmol) in toluene (10 mL) was added to a solution of B(C6F5)3 (175 mg, 0,34 mmol) in the same solvent (5 mL) at -20°C, resulting in an immediate color change from green to orange. The solution was stirred for 15 min at room temperature. An aliquot of the solution was examined by EPR, indicating the formation of a new compound (see Results and Figure 1): g = 1.995 (multiplet), aMo = 38.91 G, aH = 6.79 G (3H), aH = 4.42 (3H), aH = 4.55 G (2H), aH = 4.55 G (2H). Cooling the filtered toluene solution to -20°C or -80°C gave either no precipitation or an oil, depending on concentration. The same occurred upon diffusion of pentane into a concentrated (and unfiltered) dichloromethane solution.Reaction of CpMo( 4 -C4H4Me2-2,3)(CH3)2 with B(C6F5)3. A procedure identical to that described in the previous section gave rise to the same observations. EPR (toluene): g = 1.995 (multiplet), aMo = 34.5 G. The values of the other hyperfine constants was not optimized by digital fitting. The appearance of this spectrum, however, was essentially identical with that of the butadiene analogue (vide supra), except for a slightly worse resolution.Reaction ofCpMo( 4 -C4H6)(CH3)2 with B(C6F5)3 in the presence of PMe3. Following the same procedure as described above, PMe3 (1 mL, 1 M in toluene, 1 mmol) was added via a syringe at room temperature. No spectral modification were observed by ESR. Formation of [CpMo( 4 -C4H6)(µ-Cl)(µ-CH2)(O)MoCp][B(C6F5)CH3]. After carrying out a reaction between CpMo( 4 -C4H6)(CH3)2 and B(C6F5)3 as indicated above, the toluene structure refinement for compound [CpMo( 4 -C4H6)(µ-Cl)(µ-CH2)(O)MoCp][B(C6F5)3CH3] are summarized in Table

Figure 1 .

 1 Figure 1. EPR monitoring of the reaction between CpMo( 4 -C4H6)Me2 and B(C6F5)3 in toluene

  Figure 1. EPR monitoring of the reaction between CpMo( 4 -C4H6)Me2 and B(C6F5)3 in toluene

Figure 2 .Figure 3 . 23 Figure 1 .

 23231 Figure 2. An ORTEP view of cation (above) and anion (below) of compound [CpMo( 4 -

  Figure 2

Table 2 .

 2 Crystal data and structure refinement for [CpMo( 4 -C4H6)(µ-Cl)(µ-CH2)(O)MoCp]-

	[B(C6F5)CH3]
	Empirical formula	C34H19BClF15Mo2O
	Formula weight	968.63
	Temperature	173(2) K
	Crystal system	Monoclinic
	Space group	P21/c
	a (Å).	14.2045(12)
	b (Å)	18.5778(14)
	c (Å)	13.1036(10)
	 (°).	107.380(2)
	Volume (Å 3

Acknowledgement. We are grateful to the CNRS (Programme Catalyse pour l'Industrie et l'Environnement) and COST D17 (Working group on "Environmentally friendly catalysts for olefin polymerization") for support of this work. E.L.G thanks the MENRT for a doctoral fellowship. Thanks are also extended to the Servei Central d'Instrumentació Científica (CSIC) of the university Jaume I for providing us with mass spectrometry and x-ray facilities.

solvent was completely evaporated. The residue was then extracted with 3 mL of CH2Cl2, followed by filtration through Celite (0.5x1cm). A slow diffusion of pentane in this solution at room temperature afforded red crystals. Yield = 148 mg; 90 %. Anal. Calcd. for