Exploring the fate of the tris(pentafluorophenyl)borane radical anion in weakly coordinating solvents

Abstract

We report a kinetic and mechanistic study into the one-electron reduction of the archetypal Lewis acid tris(pentafluorophenyl)borane, B(C₆F₅)₃, in dichloromethane and 1,2-difluorobenzene. Electrochemical experiments, combined with digital simulations, DFT computational studies and multinuclear NMR analysis allow us to obtain thermodynamic, kinetic and mechanistic information relating to the redox activity of B(C₆F₅)₃. We show that tris(pentafluorophenyl)borane undergoes a quasi-reversible one-electron reduction followed by rapid chemical decomposition of the B(C₆F₅)₃˙⁻ radical anion intermediate via a solvolytic radical pathway. The reaction products form various four-coordinate borates of which [B(C₆F₅)₄]⁻ is a very minor product. The rate of the follow-up chemical step has a pseudo-first order rate constant of the order of 6 s⁻¹. This value is three orders of magnitude larger than that found in previous studies performed in the donor solvent, tetrahydrofuran. The standard reduction potential of B(C₆F₅)₃ is reported for the first time as −1.79 ± 0.1 V and −1.65 ± 0.1 V vs. ferrocene/ferrocenium in dichloromethane and 1,2-difluorobenzene respectively.

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Introduction

The preparation of tris(pentafluorophenyl)borane, B(C₆F₅)₃, was first reported by Massey and co-workers in 1963.^1–3 It was noted that the compound formed strong adducts with a number of different Lewis bases. The Lewis acidity of B(C₆F₅)₃ was later measured and determined to be intermediate between BF₃ and BCl₃.^3–5 Unlike the boron trihalides, however, B(C₆F₅)₃ is a relatively thermally stable solid that exhibits a good resistance to hydrolysis.^6,7 B(C₆F₅)₃ therefore offers an unprecedented ease of handling, combined with strong Lewis acidity and adequate steric bulk. It is for this reason that Piers and Chivers described B(C₆F₅)₃ as “the ideal boron-based Lewis acid”.⁷

B(C₆F₅)₃ has been employed as a key component in a number of important applications relating to synthetic organic transformations,^8–12 the preparation of weakly coordinating anions,^13–15 and the activation of olefin polymerization catalysts.^16–20 Since the pioneering work of Stephan et al.²¹ in 2006, B(C₆F₅)₃ has become the archetypal Lewis acid in Frustrated Lewis Pair (FLP) chemistry^22–25 – currently a highly active area of research with applications in hydrogenation reactions^26,27 and small molecule activation.^28–38

In addition to its interesting Lewis acidic properties, the ability of B(C₆F₅)₃ to act as a one-electron oxidant was accidentally discovered by Norton's group in 1999.³⁹ Erker and co-workers had previously demonstrated that B(C₆F₅)₃ could be used to open zirconocycles to generate effective olefin polymerization catalysts.⁴⁰ When Norton and co-workers attempted to extend this concept to heteroatom-substituted zirconocycles, they noted the partial oxidation of their catalyst. Soon afterwards, Green et al. also observed the one-electron oxidation of a η²-vinyl molybdenum complex in the presence of B(C₆F₅)₃.⁴¹ Norton's group then went on to investigate the redox properties of B(C₆F₅)₃ by reducing it using decamethylcobaltocene (Cp*₂Co) and studying the resulting B(C₆F₅)₃˙⁻ intermediate via EPR and UV-vis spectroscopic methods.⁴² The rate of decomposition of the B(C₆F₅)₃˙⁻ species was determined to be ca. 5.7 × 10⁻³ s⁻¹ at 23 °C using UV-vis spectrophotometry (λ_max = 603 nm).⁴² However, this value should be treated with some caution given that the experiments were performed in the donor solvent THF and the formation of the (THF)·B(C₆F₅)₃ adduct is well known.⁴³

Despite there being an interest in the redox properties of B(C₆F₅)₃, its direct electrochemical reduction initially proved to be difficult for two reasons. Early attempts to record the cyclic voltammetry of B(C₆F₅)₃ were performed using either coordinating solvents, e.g. THF, and/or common supporting electrolyte salts of ClO₄⁻, PF₆⁻ or BF₄⁻ that can react with B(C₆F₅)₃. These experimental conditions resulted in ill-defined cyclic voltammograms at best, and only enabled predictions of the reduction potential of B(C₆F₅)₃.^39,44 The first direct voltammetric reduction of B(C₆F₅)₃ was reported by this group and collaborators in 2011.⁴⁵ This was achieved by virtue of a carefully selected system, comprising CH₂Cl₂ solvent and a non-coordinating electrolyte based on Kobayashi's anion, [ⁿBu₄N][B(3,5-(CF₃)₂C₆H₃)₄].⁴⁶ However, no further mechanistic or kinetic studies were undertaken at that time. In this report we address this by extracting mechanistic and kinetic parameters for the reduction of B(C₆F₅)₃ in solvents of low donor strength, whilst also determining the reaction products. This will allow for a better understanding of the one-electron redox chemistry of B(C₆F₅)₃.

Experimental section

Materials and general methods

All synthetic reactions and manipulations were performed under a dry N₂ atmosphere (BOC Gases) using either a Saffron glovebox or standard Schlenk-line techniques on a dual manifold vacuum/inert gas line. All glassware was dried under vacuum at 170 °C before use. Diethyl ether and light petroleum either were dried via distillation over Na/benzophenone diketyl; toluene was dried via distillation over molten Na; dichloromethane was dried via distillation over CaH₂; 1,2-difluorobenzene (DFB) was dried by stirring over P₄O₁₀ and triply distilled prior to use. All solvents were sparged with nitrogen gas to remove any trace of dissolved oxygen and stored in ampules over activated 3 Å molecular sieves. Bromopentafluorobenzene was purchased from Fluorochem and used without further purification. Mg turnings were purchased from Alfa Aesar and used as supplied. All other reagents were purchased from SigmaAldrich and were of the highest grade available and used without further purification. Deuterated NMR solvents (CDCl₃, 99.8%; CD₃CN, 99.9%) were purchased from Cambridge Isotope Laboratories Inc. and were dried over P₄O₁₀, degassed using a triple freeze–pump–thaw cycle and stored over activated 3 Å molecular sieves. B(C₆F₅)₃ and [Li(OEt₂)₃][B(C₆F₅)₄] were prepared according to literature methods.^47,48 NMR spectra were recorded using a Bruker Advance DPX-300 MHz spectrometer. Chemical shifts are reported in ppm and are referenced relative to appropriate standards; ¹⁹F is relative to CFCl₃, ¹¹B is relative to Et₂O·BF₃. For NMR experiments performed in non-deuterated solvents a C₆D₆ insert was used. ESI-MS spectra were recorded using a Shimadzu LCMS 2010EV spectrometer in negative ESI mode. EPR spectra were recorded using a Bruker ER200D spectrometer fitted with a dual-mode (ER4116M) X-band cavity and interfaced to an EMX control system. A flow-through cryostat used in conjunction with a Eurotherm (B-VT-2000) variable temperature controller provided temperatures ranging from 80–180 K.

Preparation of [ⁿBu₄N][B(C₆F₅)₄] electrolyte

A solution of [Li(OEt₂)₃][B(C₆F₅)₄] (17.16 g, 18.9 mmol) in dry, degassed CH₂Cl₂ (50 mL) was added to a solution of tetrabutylammonium chloride (5.25 g, 18.9 mmol) in dry, degassed CH₂Cl₂ (50 mL). Rapid formation of a fine off-white precipitate resulted. The reaction mixture was left to stir overnight before filtration and removal of the solvent in vacuo to yield the crude product (16.97 g, 97%) as an off-white solid. [ⁿBu₄N][B(C₆F₅)₄] was recrystallized three times according to the method of LeSuer et. al. prior to use as a supporting electrolyte.⁴⁹¹H NMR (400 MHz, CDCl₃): δ 3.11 (m, 8H), 1.58 (m, 8H), 1.36 (m, 8H), 0.96 (t, J = 8 Hz) ppm; ¹¹B NMR (96 MHz, CDCl₃): δ −16.7 ppm; ¹⁹F NMR (282 MHz, CDCl₃): δ −132.5 (m, 8F, o-F), −162.8 (m, 4F, p-F), −166.7 (m, 8F, m-F) ppm.

Electrochemistry

All electrochemical experiments were performed under an inert atmosphere using an Autolab PGSTAT 30 computer-controlled potentiostat. Cyclic voltammetry (CV) was performed using a three-electrode configuration consisting of either a Pt macrodisk working electrode (GoodFellow, Cambridge, UK; 99.99%; area 1.4 ± 0.5 × 10⁻³ cm²) or a Pt microdisk working electrode (GoodFellow, Cambridge, UK; 99.99%; radius 30.5 ± 0.5 μm), combined with a Pt wire counter electrode and a Ag wire pseudoreference electrode. The Pt working electrodes were polished between experiments using successive grades of alumina slurries (from 1.0 to 0.3 μm), rinsed in distilled water and subjected to brief ultrasonication to remove any adhered alumina microparticles. The electrodes were then dried in an oven at 120 °C to remove any residual traces of water. The Pt working electrode areas were calibrated for each experiment using a 5.0 mM ferrocene solution in either CH₃CN or CH₂Cl₂ solvent containing 0.1 M [ⁿBu₄N][PF₆] as the supporting electrolyte. The Pt macrodisk working electrode area was accurately determined by construction of a Randles–Sevcik plot from cyclic voltammograms recorded at varying scan rates (50–750 mV s⁻¹).⁵⁰ The Pt microdisk working electrode area was accurately determined from the steady state current, measured using linear sweep voltammetry (scan rate = 5 mV s⁻¹).⁵⁰ The Ag wire pseudoreference electrodes were calibrated to the ferrocene/ferrocenium couple in CH₂Cl₂ at the end of each run to allow for any drift in potential, following IUPAC recommendations.⁵¹ All electrochemical measurements were performed at ambient temperatures under an inert N₂ atmosphere in either CH₂Cl₂ or 1,2-difluorobenzene (DFB) containing 0.05 M [ⁿBu₄N][B(C₆F₅)₄] as the supporting electrolyte, and iR-compensated using positive-feedback to within 85 ± 5% of the solution uncompensated resistance. CV simulations were performed using DigiElch – Professional (v 7.030) software.

Computational modelling

All calculations were performed using the Gaussian 09 computational package.⁵² Geometry optimisation, vibration frequencies and spin distribution calculations have been carried out using the three-parameter exchange functional of Becke⁵³ (B3) and the correlation functional of Lee, Yang, and Parr (LYP), B3LYP.⁵⁴ In each case an all electron 6-311+G(d,p) basis set has been implemented. Structures were geometry optimised in the gas phase with the default convergence criteria and confirmed as minima through frequency calculations. Zero-point energies and thermodynamic properties were calculated at 298.15 K/1 atm. All calculations have been performed at spin-unrestricted level of theory.

Results and discussion

Electrochemical experiments

The direct voltammetric reduction of B(C₆F₅)₃ was explored at a macrodisk electrode using cyclic voltammetry (Fig. 1 and ESI 1†) with weakly coordinating electrolyte systems comprising of a solution of [ⁿBu₄N][B(C₆F₅)₄] in either CH₂Cl₂ or DFB low-donor solvents. Similar voltammetric behaviour was observed in either solvent. Upon first scanning towards more negative potentials a reduction wave was observed at −1.82 and −1.67 V vs. Cp₂Fe^0/+ (at 100 mV s⁻¹) for CH₂Cl₂ and DFB respectively. At slow scan rates the reduction wave appears to be irreversible, with no corresponding oxidation peak observed upon reversing the scan direction. However, at faster scan rates (up to 5000 mV s⁻¹) a small oxidation wave was observable as the scan rate was increased.

Fig. 1 Overlaid cyclic voltammograms of B(C₆F₅)₃ in DFB (5.1 mM, 0.05 M [ⁿBu₄N][B(C₆F₅)₄]) recorded at scan rates of 100–5000 mV s⁻¹ at a Pt macrodisk working electrode.

The observed voltammetric behaviour is indicative of an EC process, using Testa–Reinmuth notation.⁵⁵ B(C₆F₅)₃ undergoes a heterogeneous, electrochemically quasi-reversible reduction (E-step) at the electrode. This is rapidly followed by an irreversible, homogeneous chemical step in the solution (C-step) to form an electroinactive product. As the scan rate is increased, the kinetics of the chemical follow-up step begin to be outrun on the voltammetric timescale, and the re-oxidation of the B(C₆F₅)₃˙⁻ intermediate back to the neutral B(C₆F₅)₃ parent compound is observed (Scheme 1).

Scheme 1 Postulated EC mechanism of B(C₆F₅)₃ reduction.

Upon closer inspection, additional small reduction and corresponding oxidation waves are also observed at more cathodic potentials than the main B(C₆F₅)₃ reduction peak. Their broad, symmetric wave shape appears to be characteristic of surface-adsorbed species. In light of the NMR analysis of the reaction products (discussed below) we tentatively attribute this to the formation of radical species on the electrode surface during the decomposition process of B(C₆F₅)₃˙⁻.

In order to quantitatively investigate the mechanism of B(C₆F₅)₃ reduction, we first need to determine the number of electrons (n) involved in the reduction process. The diffusion coefficient (D) of the neutral B(C₆F₅)₃ is also required. Values of n and D were determined simultaneously by performing potential-step chronoamperometry at a microdisk electrode and numerically fitting the experimental data using the Shoup–Szabo approximation.⁵⁶ This accurately predicts the current–time response over the entire time domain to a maximum error of less than 0.6% provided that both the concentration of the redox active species and the radius of the microelectrode are known.⁵⁶ Chronoamperomograms were recorded for the reduction of B(C₆F₅)₃ in both CH₂Cl₂ and DFB, and are shown in Fig. 2 and ESI 2† along with the Shoup–Szabo best fits calculated in Origin™. The Shoup–Szabo best fits confirm that B(C₆F₅)₃ undergoes a one-electron (n = 1) reduction in both solvent systems, with diffusion coefficients of 8.5 ± 0.1 × 10⁻⁶ and 3.9 ± 0.1 × 10⁻⁶ cm² s⁻¹ for CH₂Cl₂ and DFB respectively. The difference in the value of the diffusion coefficient between CH₂Cl₂ and DFB likely reflects the greater viscosity of DFB.

Fig. 2 Experimental chronoamperogram (crosses) and Shoup–Szabo best fit (solid line) for the reduction of B(C₆F₅)₃ in CH₂Cl₂ (4.8 mM, 0.05 M [ⁿBu₄N][B(C₆F₅)₄]) at a 31 μm radius Pt microdisk.

To confirm the diffusion coefficients of B(C₆F₅)₃, steady-state (scan rate = 5 mV s⁻¹) linear sweep voltammetry was performed at a microdisk electrode in both solvent systems (ESI 3†). Assuming a one-electron reduction process, the diffusion coefficient of B(C₆F₅)₃ can be determined from the measured steady-state limiting current,⁵⁰ and was found to be 8.4 ± 0.1 × 10⁻⁶ and 4.7 ± 0.1 × 10⁻⁶ cm² s⁻¹ for CH₂Cl₂ and DFB respectively. Considering the experimental error encountered in accurately measuring a steady state current, these D values are in excellent agreement with those obtained using chronoamperometry.

Digital simulation of mechanistic and kinetic parameters

Having ascertained both the number of electrons transferred during the reduction of B(C₆F₅)₃ and its diffusion coefficient in the solvents studied, we next performed digital simulations of the experimentally observed cyclic voltammetric data in order to extract kinetic and thermodynamic parameters. A variety of plausible mechanisms for the decay of B(C₆F₅)₃˙⁻ were simulated. These included unimolecular chemical decomposition to form further redox active products (ECE), disproportionation (DISP), and bimolecular radical recombination (EC₂ or EC₂E) mechanisms. None of these mechanisms were found to fit the observed voltammetry. However, the voltammetric reduction of B(C₆F₅)₃ produced very good fits between simulation and experiment, as shown in Fig. 3a and 3b, when modelled as an EC process (Scheme 1). The globally optimized parameters for the electrochemical reduction (standard potential, E⁰; charge transfer coefficient, α; and standard electron transfer rate constant, k⁰) and the pseudo-first order rate constant, k_f, for the homogeneous chemical decay step are given in Table 1.

Fig. 3 Experimental (solid line) and simulated (open circles) overlaid cyclic voltammograms of B(C₆F₅)₃ (4.9 mM, 0.05 M [ⁿBu₄N][B(C₆F₅)₄]) at scan rates of ν = 100, 500, 1000, 2000 and 5000 mV s⁻¹ for dichloromethane (top) and 1,2-difluorobenzene (bottom) electrolyte solvents. Inset: plots comparing simulated (open circles) and experimental (closed circles) peak potentials (E_p) and peak currents (I_p) vs. the logarithm of scan rate (ν).

Table 1 Simulated parameters for the voltammetric reduction of B(C₆F₅)₃

Parameter	CH2Cl2	DFB
E ^0/V vs. Cp2Fe^0/+	−1.79 ± 0.01	−1.65 ± 0.01
α	0.49 ± 0.02	0.52 ± 0.02
k ^0/10^−3 cm^2 s^−1	13 ± 0.3	11 ± 0.2
DB(C6F5)3/10^−6 cm^2 s^−1	8.5 ± 0.1	3.9 ± 0.1
DB(C6F5)3˙^−/10^−6 cm^2 s^−1	8.4 ± 0.1	3.9 ± 0.1
k f/s^−1	6.1 ± 0.1	7.7 ± 0.2

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The values for the standard reduction potential of B(C₆F₅)₃ are within the error range of the value predicted by Cummings et al.⁴⁴ and the value reported in our earlier work for the direct measurement of the B(C₆F₅)₃ reduction potential.⁴⁵ The ca. 200 mV difference between the value we previously reported and those herein reflects the subtle but important difference between the mid-peak potential, E_mid, that we reported previously, and the thermodynamic standard potential, E⁰, obtained via simulation. The modest value of the standard electron transfer rate constant (k°) suggests that the reduction of B(C₆F₅)₃ is an electrochemically quasi-reversible process (vide infra). However, the chemical reactivity of B(C₆F₅)₃˙⁻ limits the observation of the corresponding (oxidative) back peak, except at relatively fast scan rates. B(C₆F₅)₃˙⁻ undergoes a rapid follow-up chemical reaction with pseudo-first order rate constants (k_f) of 6.1 ± 0.1 and 7.7 ± 0.1 s⁻¹ in CH₂Cl₂ and DFB respectively. These values obtained in weakly coordinating solvents are ca. three orders of magnitude larger than the decomposition rate constant reported by Norton et al. using EPR measurements in the donor solvent, THF.⁴² Indeed the follow-up reaction in CH₂Cl₂ or DFB occurs so rapidly as to preclude any kinetic measurements using EPR techniques.⁵⁷

Computational modelling of B(C₆F₅)₃˙⁻

Given the Lewis acidity of B(C₆F₅)₃ it is somewhat curious that the rate of electron transfer, k⁰, is only of the order of 10⁻² cm s⁻¹ – i.e. it exhibits quasi-reversible electron transfer kinetics. We therefore performed DFT computational modelling of the B(C₆F₅)₃ and B(C₆F₅)₃˙⁻ species to ascertain the optimised (gas phase) geometry, SOMO, and comparative charge and spin density distributions. For spin-unrestricted type of calculations the unpaired molecular orbital of the B(C₆F₅)₃˙⁻ complex is best represented by a spin down (beta) SOMO shown in ESI 4a.†Fig. 4 shows the resulting spin density distribution in B(C₆F₅)₃˙⁻ and ESI 4b and ESI 4c† show the charge distribution based on Mulliken electron population analysis of B(C₆F₅)₃ and B(C₆F₅)₃˙⁻ respectively.

Fig. 4 The spin density distribution calculated using the spin-unrestricted B3LYP 6-311+G(d,p) basis set for the corresponding SOMO of the B(C₆F₅)₃˙⁻ radical anion.

The optimised geometries of B(C₆F₅)₃ and B(C₆F₅)₃˙⁻ reveal little deviation from planarity around the trigonal planar boron centre, although the torsional angle between the aryl rings and the central plane containing the boron atom is reduced from 37°, in the case of B(C₆F₅)₃, to 34°, in the B(C₆F₅)₃˙⁻ species. This is due to delocalisation of some spin density onto the perfluoroaryl rings, within the SOMO.

Marcus theory describes the rate of adiabatic electron transfer in terms of the reorganisation energy (λ). This is comprised of contributions from inner (λ_i) and outer (λ_o) sphere electron transfer. λ_i describes changes in bond strength and bond angles during electron transfer, and λ_o depends on the reorientation of solvent dipoles and electronic polarization within the solvent molecules.⁵⁰ Given that DFT calculations indicate there is no significant change between the structures of B(C₆F₅)₃ and B(C₆F₅)₃˙⁻, we infer that the solvent reorganisation energy (λ_o) is the rate-limiting factor during electron transfer. The relationship between Marcus theory and Butler–Volmer kinetics applied in our voltammetric simulation are described via the charge transfer coefficient, α:

Given that we obtain values of α that are close to 0.5 in either solvent system (Table 1), it is confirmed that the reorganisation energy is very much larger than the Gibbs energy for this reaction.

DFT modelling shows that, in its reduced form, both spin and charge density are predominantly located on the central boron atom of the B(C₆F₅)₃˙⁻ radical anion. Together with the indication that solvent reorganisation is strongly coupled to the electron transfer, these findings may indicate that decomposition of the B(C₆F₅)₃˙⁻ radical anion predominantly proceeds via solvolysis at the boron centre to form four-coordinate borate species (vide infra).

NMR and MS characterisation of B(C₆F₅)₃˙⁻ decomposition products

Finally, we attempted to elucidate the reaction products resulting from the decay of B(C₆F₅)₃˙⁻via¹¹B and ¹⁹F NMR spectroscopy. A colourless solution of B(C₆F₅)₃ (49 mg, 0.1 mmol) in dry, degassed CH₂Cl₂ or DFB (4 mL) was added to a saturated brown/yellow solution of Cp*₂Co (33 mg, 0.1 mmol) in either CH₂Cl₂ or DFB (4 mL) under an inert N₂ atmosphere. Immediately upon mixing, the solution initially turned a deep blue colour, indicative of the B(C₆F₅)₃˙⁻ radical anion, which then very rapidly formed a dark yellow/brown solution upon standing. An aliquot was taken and NMR analysis performed directly on the reaction mixture using a C₆D₆ insert. Then, the solvent was removed and rigorously dried in vacuo to yield a brown residue, which was taken up in either CDCl₃ or CD₃CN (0.8 mL) for further NMR analysis.

The ¹¹B NMR (96.3 MHz, CH₂Cl₂ with C₆D₆ insert) spectrum obtained after the chemical reduction of B(C₆F₅)₃ in CH₂Cl₂ reveal a mixture of five radical decomposition products formed via reaction with the solvent. These are listed in Table 2. The identity of each product has also been tentatively assigned, where possible, by comparison with known literature values.^58–60 The doublet observed at δ −0.52 ppm has a coupling constant of 77 Hz, hence we assign this to an as yet unidentified four-coordinate borate species containing one B–H bond (vide infra) representing ca. 18% of the products formed.

Table 2 ¹¹B NMR data and analysis for the reduction of B(C₆F₅)₃ with Cp*₂Co in dichloromethane

δ/ppm	Multiplicity	Assignment
a Speculative – see text. b Ref. 58. c Ref. 59. d See Experimental section. e Ref. 60.
−0.52	d; ^1JB–H = 77 Hz	[HClB(C6F5)2]^−^a
−3.82	s	[Cl2B(C6F5)2]^−^b
−7.04	br s	[ClB(C6F5)3]^−^c
−13.40	s	[B(C6F5)4]^−^d
−25.40	d; ^1JB–H = 92 Hz	[HB(C6F5)3]^−^e

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The corresponding ¹⁹F NMR (282 MHz, CH₂Cl₂ with C₆D₆ insert) spectrum of this same sample is complex. Five signals were observed as doublets of multiplets (arising from second-order spin–spin coupling) between δ −132.0 and −135.9 ppm, corresponding to ortho-F nuclei on the aryl rings. A further series of broad overlapping multiplets were observed from δ −162.0 to −165.5 ppm and from δ −165.7 to −168.4 ppm, corresponding to aryl fluorine nuclei in the para- and meta-positions respectively. Whilst these latter overlapping signals could not be assigned, the ortho-F signals are listed in Table 3 together with their relative product distribution determined by integration of the peaks. A tentative assignment has been made by comparison to literature values.^58–60

Table 3 ¹⁹F NMR data and analysis for the reduction of B(C₆F₅)₃ with Cp*₂Co in dichloromethane

δ/ppm	Multiplicity	Relative % product distribution	Assignment
a Ref. 59. b Ref. 60. c Ref. 58. d See Experimental section. e Speculative – see text.
−132.4	dm	39 ± 1	[ClB(C6F5)3]^−^a
−133.7	dm	14 ± 2	[HB(C6F5)3]^−^b
−134.4	dm	21 ± 2	[Cl2B(C6F5)2]^−^c
−135.4	m	6 ± 1	[B(C6F5)4]^−^d
−135.8	dm	20 ± 3	[HClB(C6F5)2]^−^e

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The ¹⁹F peak at δ −135.8 ppm is as yet unassigned, but it is likely to correspond to the unidentified product giving rise to the doublet at δ −0.52 ppm in the ¹¹B NMR spectrum. Based on chemical intuition, if we speculate that this is in fact a product of the form [HClB(C₆F₅)₂]⁻ then it forms ca. 20% of the product distribution (based on integration of the ¹⁹F NMR signals).

When the reduction of B(C₆F₅)₃ is performed in DFB three signals are observed in the ¹¹B NMR spectrum (96.3 MHz, DFB with C₆D₆ insert) at: δ −3.88 (sharp s, unassigned), −13.28 (s, [B(C₆F₅)₄]⁻), and −0.28 to 1.16 (br m, unassigned) ppm. The relative product distribution of these peaks is ca. 55, 5, and 40% respectively. The broad multiplet between −0.28 and 1.16 ppm comprising ca. 40% of products most likely corresponds to several structurally related products giving rise to overlapping signals. This is further borne out upon examination of the ¹⁹F NMR spectrum (282 MHz, DFB with C₆D₆ insert) whereby a complex series of overlapping muliplets is observed between δ −131.7 to −135.4 (dm, ortho-aryl F), −161.4 to −164.7 (tm, para-aryl F), and −166.0 to −166.6 (m, meta-aryl F) ppm, indicative of multiple products containing fluorinated aryl rings. No significant change in any of the NMR spectra was observed upon exchanging the solvent to CDCl₃ or CD₃CN.

The lack of NMR data for borate and borane species in DFB hindered full assignment of the products. We can only assign the tetrakis(pentafluorophenyl)borate species with any certainty. However, given that the rate of decomposition of B(C₆F₅)₃˙⁻ is pseudo-first order and similar in either solvent system, it is likely that solvolysis occurs in DFB as it does in dichloromethane – the solvent being in vast excess in both cases. Comparison of the ¹¹B and ¹⁹F NMR spectra of authentic samples of [^tBu₄N][BF₄] and [^tBu₄N][FB(C₆F₅)₃] revealed no evidence for the formation of any reaction products containing B–F bonds in either the dichloromethane or DFB solvent systems. Hence, we speculate that the unidentified (major) products of the decomposition of B(C₆F₅)₃˙⁻ in DFB are likely to be borate species of the form [(C₆F₅)_3−xB(C₆H₄F)_x]⁻. Furthermore, ¹H NMR (CDCl₃) analysis of the products from either solvent reaction system revealed no evidence of H-abstraction from the Cp*₂Co.

Whilst the decomposition of the B(C₆F₅)₃˙⁻ radical anion via a solvolytic pathway may not be surprising, the key point to note is that, in contradiction to Norton's earlier suggestion,⁴² decamethylcobaltocenium tetrakis(pentafluorophenyl)borate is the very minor (ca. 5%) product of this reaction. Further, whilst CH₂Cl₂ is known to be prone to radical attack, DFB is usually considered to be less susceptible.⁶¹ Yet, the B(C₆F₅)₃˙⁻ radical anion intermediate must be a sufficiently reactive species as to decompose via solvolytic pathways at a similar rate in either weakly coordinating solvent.

Interestingly mass spectrometric characterisation (ESI-MS) of the reaction products from either DFB or dichloromethane could only detect one product with molecular ion peaks at m/z values of 678.90 (100%, M⁻), 677.80 (24.98%) and 679.90 (25.75%) Da. This is indicative of the tetrakis(pentafluorophenyl)borate anion; however, given the likelihood of fragmentation and recombination reactions in the mass spectrometer this observation must be interpreted with some caution.

Conclusions

We have studied the direct voltammetric reduction of tris(pentafluorophenyl)borane, B(C₆F₅)₃, in two weakly coordinating solvents, dichloromethane and 1,2-difluorobenzene. In either case cyclic voltammetric combined with digital simulations indicate that the reaction follows an EC mechanism whereby the electro-generated B(C₆F₅)₃˙⁻ radical anion undergoes a rapid chemical step in solution to form redox inactive products. Multinuclear NMR analysis of these products generated by the chemical reduction of B(C₆F₅)₃ with decamethylcobaltocene are indicative of the formation of several four-coordinate borate species arising from solvolytic radical reaction pathways. Solvolysis resulting in reactivity predominantly at the boron centre is further supported by spin density and charge distribution DFT calculations.

Chronoamperometry at a microdisk electrode has allowed us to ascertain that the reduction of B(C₆F₅)₃ is indeed a one-electron process and report diffusion coefficients in each solvent.

Values of the pseudo-first order rate constants for the chemical decomposition of the B(C₆F₅)₃˙⁻ radical anion were determined to be 6.1 ± 0.1 s⁻¹ and 7.7 ± 0.2 s⁻¹ in dichloromethane and 1,2-difluorobenzene respectively. The thermodynamic standard potential, E°, of B(C₆F₅)₃ was also extracted for the first time with values of −1.79 ± 0.1 V and −1.65 0.1 V vs. ferrocene/ferrocenium in dichloromethane and 1,2-difluorobenzene respectively. These values are in close agreement with previous estimates based on either the reduction peak potentials⁴⁴ or the measured mid-peak potentials,⁴⁵ which do not strictly correspond to the thermodynamic potential, E°.

The rate of decomposition of the radical anion is sufficiently fast in solvents of low donor strength that we were unable to measure a signal from B(C₆F₅)₃˙⁻ using EPR experiments, even at low temperatures – in stark contrast to previous reports using strong donor solvents.⁴²

Thus, almost fifty years after tris(pentafluorophenyl)borane was first discovered, we are able to report thermodynamic and kinetic parameters relating to its redox properties in selected weakly donor solvents. Once again, we emphasise the importance of carefully considering the choice of solvent when attempting to study the chemistry of this and its related electrophilic and Lewis acidic boranes.

Acknowledgements

EJL thanks the EPSRC for financial support via a Doctoral Training Account studentship. GGW thanks the Royal Society for support via a University Research Fellowship. AEA thanks Imperial College, London for support via a Junior Research Fellowship and the Royal Commission for the Exhibition of 1851 for a Research Fellowship. The authors thank Dr Myles R. Cheesman and the Henry Welcome Unit for Biological EPR, University of East Anglia for assistance with EPR studies.

Notes and references

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Footnote

† Electronic supplementary information (ESI) available: Cyclic voltammogram, chronoamperometry, linear sweep voltammetry, DFT calculations (PDF). See DOI: 10.1039/c2dt31622f

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