Synergistic effects of CH₃CO₂H and Ca²⁺ on C–H bond activation by MnO₄⁻

Abstract

The activation of metal-oxo species with Lewis acids is of current interest. In this work, the effects of a weak Brønsted acid such as CH₃CO₂H and a weak Lewis acid such as Ca²⁺ on C–H bond activation by KMnO₄ have been investigated. Although MnO₄⁻ is rather non-basic (pK_a of MnO₃(OH) = −2.25), it can be activated by AcOH or Ca²⁺ to oxidize cyclohexane at room temperature to give cyclohexanone as the major product. A synergistic effect occurs when both AcOH and Ca²⁺ are present; the relative rates for the oxidation of cyclohexane by MnO₄⁻/AcOH, MnO₄⁻/Ca²⁺ and MnO₄⁻/AcOH/Ca²⁺ are 1 : 73 : 198. DFT calculations show that in the active intermediate of MnO₄⁻/AcOH/Ca²⁺, MnO₄⁻ is H-bonded to 3 AcOH molecules, while Ca²⁺ is bonded to 3 AcOH molecules as well as to an oxo ligand of MnO₄⁻. Our results also suggest that these synergistic activating effects of a weak Brønsted acid and a weak Lewis acid should be applicable to a variety of metal-oxo species.

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Introduction

The chemistry of metal-oxo species has been of great interest to chemists because of the key roles they play in various chemical and biological systems.^1–4 In recent years, the use of Lewis acids to enhance the reactivity of metal-oxo species has attracted much attention,^5,6 especially after the discovery of a Ca²⁺ in the active site of the oxygen-evolving center (OEC) of photosystem II, which contains a manganese calcium oxo cluster (Mn₄CaO₅).^7,8

Various metal-oxo species are readily activated by strong Lewis acids such as BF₃, Sc(OTf)₃ and other M³⁺ ions toward the oxidation of organic substrates.^9–19 There are also a number of studies on the interaction of relatively weak Lewis acids such as Ca²⁺ and other group II ions with metal-oxo species, and as expected their activating effects are much smaller than those of Sc(OTf)₃ or other M³⁺ ions.^{12,15,20–24}

We recently reported the use of a weak Brønsted acid such as acetic acid in combination with a weak Lewis acid such as a group II ion to activate RuO₄⁻.²⁵ We found a remarkable cooperative activating effect of these two acids on RuO₄⁻ toward the oxidation of alkanes. The oxo ligands in RuO₄⁻ are basic and can be protonated by AcOH to generate [RuO₃(OH)]OAc⁻, which readily abstracts H atoms from alkanes. In the presence of Ca²⁺, a more active intermediate is formed with [RuO₃(OH)]OAc⁻, in which Ca(II) binds to both OAc⁻ (which is H-bonded to OH) and an oxo ligand. In this work, we have carried out experimental and theoretical investigations on the use of AcOH and Ca²⁺ to activate MnO₄⁻ toward cyclohexane oxidation. MnO₄⁻ was chosen because of the relevance of Mn in the OEC of photosystem II. Remarkably, a similar synergistic activating effect of AcOH and Ca²⁺ is also observed for this oxoanion, although MnO₄⁻ is much less basic than RuO₄⁻ and is not protonated by AcOH. Our results suggest that the synergistic effects of a weak Brønsted and a weak Lewis acid should be applicable to a variety of metal-oxo species.

Results and discussion

Effects of AcOH on cyclohexane oxidation by KMnO₄

KMnO₄ does not react with cyclohexane at room temperature. However, upon addition of AcOH (3–9 M) to KMnO₄ (0.011 M) in CH₃CN in the presence of cyclohexane (1.0 M) at 23 °C, the purple color of the solution gradually changed to brown. Analysis of the solution by GC-MS and GC-FID revealed the formation of cyclohexanone (38.8%), together with a small amount of cyclohexanol (2.2%). The yields were calculated based on KMnO₄ acting as a 3-electron oxidant (product is a Mn(IV) species, see below).

Kinetics studies were carried out at constant [KMnO₄] (0.011 M) and [cyclohexane] (1.0 M) by monitoring the growth of cyclohexanone using a GC-FID (gas chromatograph-flame ionization detector). The increase in [cyclohexanone] follows pseudo-first-order kinetics (Fig. 1a), and the pseudo-first-order rate constant k_obs at [AcOH] = 3.0 M was found to be 4.3 × 10⁻⁶ s⁻¹. k_obs increases with increasing [AcOH] (3–9 M), and the plot of k_obsversus [AcOH]² is linear (Fig. 1b), suggesting that under these conditions the active oxidizing species involves MnO₄⁻ binding to two molecules of AcOH. The reaction can be represented by eqn (1) and (2).

Fig. 1 (a) Time course for the production of cyclohexanone at various [AcOH]. Conditions: KMnO₄, 0.011 M; cyclohexane, 1.0 M; at 23 °C in CH₃CN. (b) Linear plot of k_obsvs. [AcOH]². Slope = (2.84 ± 0.11) × 10⁻⁵, intercept = (2 ± 5) × 10⁻⁵, and r² = 0.992.

Under the condition that K₁ ≪ 1, the rate-law is given by eqn (3).

d[C₆H₁₀O]/dt = k_obs[MnO₄⁻] = K₁k_a[AcOH]²[c-C₆H₁₂][MnO₄⁻] (3)

From the slope of Fig. 1b, K₁k_a = (4.73 ± 0.19) × 10⁻⁷ M⁻² s⁻¹ at 23.0 °C.

The kinetic isotope effect for cyclohexane oxidation determined by competitive oxidation of an equimolar mixture of c-C₆H₁₂ and c-C₆D₁₂ was found to be 7.5 ± 0.1, indicating C–H bond cleavage in the rate-limiting step.

Effects of Ca²⁺

The oxidation of cyclohexane by KMnO₄ is also activated by Ca²⁺. Upon adding Ca(OTf)₂ (5.5 × 10⁻³ to 4.4 × 10⁻² M) to KMnO₄ (0.011 M) in CH₃CN containing cyclohexane (1.0 M), the purple color of the solution was gradually discharged together with the formation of a brown precipitate (Fig. 2). The UV/vis spectral changes indicate that there is no shift in the wavelengths of the peaks of MnO₄⁻, just a gradual decrease in their absorbances (Fig. S1†). The same result was also observed in our earlier work on the effect of BF₃ on KMnO₄.¹⁸ The solution IR of KMnO₄ and KMnO₄ + Ca(OTf)₂ in CH₃CN was also performed (Fig. S2†). KMnO₄ shows a M=O stretch at 904 cm⁻¹, but this peak was not shifted upon addition of Ca(OTf)₂; presumably the binding of Ca²⁺ to MnO₄⁻ is relatively weak.

Fig. 2 UV-vis spectral changes for the reaction of KMnO₄ (2.35 × 10⁻³ M) with cyclohexane (1.0 M) in the presence of Ca(OTf)₂ (2.35 × 10⁻³ M) in CH₃CN at 25 °C (1 mm UV cell was used in the measurement). The inset shows the change of absorbance at 250 nm and 527 nm versus time.

Analysis of the product solution by GC-FID and GC-MS revealed the formation of cyclohexanone (10.5%) and cyclohexanol (0.8%) in relatively low yields. The yields of products reached a maximum with just 1 equiv. of Ca²⁺.

Kinetics studies were also carried out at various [Ca(OTf)₂] but at constant [KMnO₄] (0.011 M) and [cyclohexane] (1.0 M). The increase in [cyclohexanone] follows pseudo-first-order kinetics (Fig. 3a), but saturation kinetics were observed on increasing [Ca(OTf)₂] (Fig. 3b), and the plot of 1/k_obsversus 1/[Ca(OTf)₂] is linear. Kinetic studies were also conducted at various [cyclohexane] (Fig. 4a), while keeping both Ca(OTf)₂ and KMnO₄ at 0.011 M. The plot of k_obsversus [cyclohexane] is linear with an intercept of (2.87 ± 3.3) × 10⁻⁴ s⁻¹ (Fig. 4b), which should be due to the decomposition of KMnO₄ in the absence of cyclohexane. An independent experiment showed that in the presence of 1 equiv. of Ca(OTf)₂, the absorbance of KMnO₄ (2.82 × 10⁻⁴ M) in CH₃CN at around 527 nm gradually decreased with time with a rate constant of (9.5 ± 1.1) × 10⁻⁴ s⁻¹ at 23 °C. Such a decomposition also accounts for the low yield of products. The kinetic results are consistent with a pre-equilibrium binding of Ca(II) to MnO₄⁻ to form an intermediate, which then oxidizes cyclohexane, as represented by eqn (4) and (5), and the rate law is shown in eqn (6).

K₂ = y-intercept/slope of Fig. 3b = (52.1 ± 3.5) M⁻¹ and k_b = 1/y-intercept = (1.1 ± 0.1) × 10⁻¹ M⁻³ s⁻¹ at 23 °C.

Fig. 3 (a) Time course for the production of cyclohexanone at various [Ca(OTf)₂]. Conditions: KMnO₄, 0.011 M; cyclohexane, 1.0 M; at 23 °C in CH₃CN. (b) Plot of k_obsversus [Ca(OTf)₂]. Inset: linear plot of 1/k_obsvs. 1/[Ca(OTf)₂]; slope = 0.303 ± 0.018, intercept = 15.8 ± 1.9, and r² = 0.99.

Fig. 4 (a) Time course for the production of cyclohexanone at various [cyclohexane]. Conditions: KMnO₄ (0.011 M), Ca(OTf)₂ (0.011 M) and cyclohexane (0.238–0.951 M) in CH₃CN at 23 °C. (b) Plot of k_obsvs. [cyclohexane]. Slope = (6.64 ± 0.29) × 10⁻³, intercept = (1.72 ± 0.02) × 10⁻², and r² = 0.994.

Synergistic effects of AcOH and Ca²⁺

The oxidation of cyclohexane in the presence of both AcOH and Ca²⁺ was investigated. As described above, the rate of oxidation of cyclohexane by KMnO₄/AcOH is second-order in [AcOH]. However, in the presence of just 1 equiv. of Ca(OTf)₂, the rate becomes first-order in [AcOH] (Fig. 5).

Fig. 5 (a) Time course for cyclohexanone production by KMnO₄ (0.011 M) in the presence of cyclohexane (1.0 M), 1 equiv. of Ca(OTf)₂ (0.011 M) and various [AcOH] in CH₃CN at 23.0 °C. (b) Plot of k_obsvs. [AcOH]. Slope = (1.56 ± 0.06) × 10⁻², y-intercept = 0.025 ± 0.003, and r² = 0.995.

On the other hand, at constant [AcOH] (2.79 M) and [cyclohexane] (1.0 M), saturation kinetics were observed on increasing [Ca(OTf)₂], and the plot of 1/k_obsversus 1/[Ca(OTf)₂] is linear (Fig. 6).

Fig. 6 (a) Time course for cyclohexane production in CH₃CN at various [Ca(OTf)₂]. Conditions: cyclohexane, 1.0 M; KMnO₄, 0.011 M; AcOH, 2.79 M. T = 23 °C. (b) Plot of k_obsversus [Ca(OTf)₂]. Inset: plot of 1/k_obsversus 1/[Ca(OTf)₂]; slope = (6.3 ± 0.2) × 10⁻², y-intercept = 11.9 ± 0.2, and r² = 0.995.

These results are consistent with the reaction scheme shown in eqn (7)–(9).

For K₃ ≪ 1, the rate law is

K₄ = y-intercept/slope of Fig. 6b insett = 188.9 ± 5.6 M⁻¹. From the slope of Fig. 5b and using [Ca(OTf)₂] = 0.011 M, K₃k_c was calculated to be (3.85 ± 0.15) × 10⁻⁴ M⁻¹ s⁻¹ (23 °C).

The kinetic isotope effect for cyclohexane oxidation by the KMnO₄/AcOH/Ca(OTf)₂ system was determined to be 5.7 ± 0.1 by competitive oxidation of an equimolar mixture of c-C₆H₁₂ and c-C₆D₁₂, indicating that C–H bond cleavage is the rate limiting step.

A comparison of the activating effects of AcOH, Ca(OTf)₂ and AcOH + Ca(OTf)₂ on cyclohexane oxidation by KMnO₄ is shown in Table 1. The rate constant (k_obs) for the oxidation of cyclohexane (1.0 M) by KMnO₄ (0.011 M) in the presence of AcOH (3.0 M) at 23 °C was found to be 5.73 × 10⁻⁶ s⁻¹. When AcOH was replaced by 1 equiv. of Ca(OTf)₂ (0.011 M), the rate constant (4.21 × 10⁻⁴ s⁻¹) is 73 times faster, but the yield is much lower (10.5%). On the other hand, in the presence of both Ca(OTf)₂ and AcOH, the rate (1.14 × 10⁻³ s⁻¹) increases by 198 fold relative to AcOH alone, and the product yield is also substantially increased to 57%, indicating synergistic effects of AcOH and Ca²⁺ in the activation of MnO₄⁻ toward C–H bond activation.

Table 1 Effects of AcOH, Ca(OTf)₂ and AcOH + Ca(OTf)₂ on the oxidation of cyclohexane by KMnO₄ in CH₃CN at 23 °C^a,b

	3.0 M AcOH (275 equiv.)	1 equiv. Ca(OTf)2 (0.011 M)	1 equiv. Ca(OTf)2 + 3.0 M of AcOH
a Reaction conditions: KMnO4, 0.011 M, cyclohexane, 1.0 M, T = 23 °C. b KMnO4 alone does not oxidize cyclohexane for >48 h at 23 °C. c The final oxidation state of the Mn product after oxidation was determined by iodometric titration. d % Yield was calculated based on KMnO4 acting as a 3-electron oxidant. Cyclohexanol is the minor product in all cases.
k obs (relative rate)	5.73 × 10^−6 s^−1 (1)	4.21 × 10^−4 s^−1 (73)	1.14 × 10^−3 s^−1 (198)
Oxidation state of Mn product^c	4.1 ± 0.2	4.1 ± 0.2	4.0 ± 0.2
% Yield of cyclohexanone^d	38.8	10.5	57.3
Cyclohexanone/cyclohexanol	18 ± 1	14 ± 1	7 ± 1

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The rate constant for the oxidation of 1.0 M toluene by (ⁿBu₄N)MnO₄ was reported to be 4.2 × 10⁻⁷ s⁻¹ at 25.0 °C.²⁶ As a comparison, we also investigated the oxidation of 1.0 M toluene by KMnO₄ in the presence of 1 equiv. Ca(OTf)₂ and 3.0 M AcOH, and the rate constant was found to be 5.2 × 10⁻³ s⁻¹ at 23 °C (Fig. S3†). Based on these data, the oxidation of toluene by KMnO₄/Ca²⁺/AcOH is over 4 orders of magnitude faster than that of MnO₄⁻ alone.

Effects of other metal ions

The effects of other metal salts, including Ba(OTf)₂, Mg(OTf)₂, and Sc(OTf)₃ on cyclohexane oxidation by KMnO₄ in the presence of AcOH have also been investigated. As expected, both the rate and yield increase with increasing Lewis acidity of the metal ion (Fig. 7a and S4–S6†). The ratio of the rate constants obtained by GC for the 4 metal ions is found to be k(Ba²⁺) : k(Ca²⁺) : k(Mg²⁺) : k(Sc³⁺) = 1 : 4.8 : 10.4 : 228; and the yields of cyclohexanone are 51%(Ba²⁺), 56%(Ca²⁺), 61%(Mg²⁺), and 75%(Sc³⁺). There is a fairly linear correlation of log(rate constant) with the pK_a of the metal ion in water, Sc³⁺(pK_a = 4.3), Mg²⁺(pK_a = 11.2), Ca²⁺(pK_a = 12.7), and Ba²⁺(pK_a = 13.4),²⁷ which is a measure of their Lewis acidity (Fig. 7b). Similar correlations of Lewis acid effects with their acidities have also been reported.^15,23

Fig. 7 (a)Time course for the production of cyclohexanone by KMnO₄ (0.01 M)/AcOH (2.78 M)/c-C₆H₁₂ (1.0 M) in the presence of 0.01 M of various metal salts. Pseudo-first-order fits give k_obs: Sc(OTf)₃ = 5.4 × 10⁻² s⁻¹, Mg(OTf)₂ = 2.3 × 10⁻³ s⁻¹, Ca(OTf)₂ = 1.1 × 10⁻³ s⁻¹, and Ba(OTf)₂ = 2 × 10⁻⁴ s⁻¹. The inset shows the time course for Sc(OTf)₃. (b) Plot of log(k_obs) vs. pK_a of metal ions. Sc³⁺(pK_a = 4.3), Mg²⁺(pK_a = 11.2), Ca²⁺(pK_a = 12.7), and Ba²⁺(pK_a = 13.4). Slope = −0.24 ± 0.04, intercept = −0.2 ± 0.5, and r² = 0.90.

Nature of the manganese product

Addition of Ca(OTf)₂ to KMnO₄ in CH₃CN gradually produced a brown solid and a colorless solution, irrespective of whether cyclohexane was present. When AcOH (3 M) was added to the resulting solution after cyclohexane oxidation, the brown solid dissolved but with only a small increase of 1–2% cyclohexanone. The oxidation state of Mn in the brown solution was determined to be 4.1 ± 0.2 by iodometric titration, which involves the addition of NBu₄I to reduce the Mnⁿ product to Mn^II and then determining the amount of I₃⁻ formed by UV-vis spectrophotometry.²⁸

Based on CHN elemental analysis, as well as K, Mn and Ca analysis by ICP-AES, it was found that the brown solids obtained from MnO₄⁻/Ca(OTf)₂ (1) and MnO₄⁻/Ca(OTf)₂/c-C₆H₁₂ (2) have similar compositions and are probably a Mn^IV–μ(O)–μ(Ca) polymer with an empirical formula close to “Ca₃Mn₇O₁₄(H₂O)₁₀(OTf)₆(CH₃CN)₂” (see the ESI†). The reason for the conversion of Mn^VII to Mn^IV in the absence of c-C₆H₁₂ is not clear, presumably due to oxidation of the solvent.

We have carried out X-ray photoelectron spectroscopy (XPS) of 1 and 2 (Fig. S7 and S8†). In the Mn3s XPS spectra (Fig. S8†), the peak splitting of ΔE = 4.8 eV indicates that the oxidation state of Mn is +4.²⁹

We have also determined the magnetic susceptibility of the solid samples of 1 and 2 using a magnetic balance. The μ_eff for each Mn was found to be 3.96 and 4.05 μ_B for 1 and 2, respectively, consistent with that of high spin Mn(IV) with S = 3/2.

No brown precipitate was observed when cyclohexane oxidation was carried out in the presence of AcOH, irrespective of whether Ca(OTf)₂ was added. The Mn oxidation state of the brown solution determined by iodometric titration is again 4.0 ± 0.2. Presumably the brown solution contains colloidal MnO₂, as observed in alkane oxidation by BF₃/MnO₄⁻.¹⁸

Theoretical calculations

DFT calculations on the mechanisms of the oxidation of cyclohexane by MnO₄⁻, MnO₄⁻/AcOH, MnO₄⁻/Ca(OTf)⁺ and MnO₄⁻/AcOH/Ca(OTf)⁺ were performed. Solvent (CH₃CN) effects have been included in the calculations.

The potential energy surfaces (PESs) for all calculated reactions are summarized in Fig. S9.† In the c-C₆H₁₂ oxidation by MnO₄⁻ in CH₃CN (Fig. 8), c-C₆H₁₂ and MnO₄⁻ first bind weakly together to form an intermediate [MnO₄⋯C₆H₁₂]⁻, INT1(MnO₄⁻). Hydrogen atom transfer (HAT) then occurs from c-C₆H₁₂ to Mn=O via a transition state, TS1(MnO₄⁻), to give a second intermediate, INT2(MnO₄⁻). The HAT step is rate-determining with a barrier height of 22.7 kcal mol⁻¹. Such a high barrier suggests that MnO₄⁻ will hardly oxidize c-C₆H₁₂ at room temperature. The carbon atom in the cyclohexyl radical of INT2(MnO₄⁻) bears a −1.01 electron spin, consistent with the spin distribution for a HAT. In the next step, the cyclohexyl radical binds to another Mn=O to generate an alkoxo intermediate [MnO₂(OH)(OC₆H₁₁)]⁻via TS2(MnO₄⁻). Subsequently, proton transfer from Mn–OH to the alkoxide in INT3(MnO₄⁻) occurs via TS3(MnO₄⁻) to generate the cyclohexanol product. The reaction mechanism is similar to that of c-C₆H₁₂ oxidation by RuO₄⁻ found in our previous study.²⁵

Fig. 8 PES and structures for cyclohexane oxidation by [MnO₄]⁻ at the B3LYP-D3(BJ)/def2-SVPD level. The relative 298 K Gibbs free energies in acetonitrile are given in kcal mol⁻¹.

In the presence of an acetic acid molecule (Fig. S10†), AcOH and MnO₄⁻ first form INT1(AcOH), [MnO₄(AcOH)⋯C₆H₁₂]⁻, through a hydrogen bond. The Mn=O that is H-bonded to AcOH is elongated to 1.611 Å, while the other three Mn=O bonds become more electrophilic and shorter with distances of 1.587–1.590 Å, compared with the bond length of 1.594 Å in MnO₄⁻. In TS1(AcOH), HAT occurs from c-C₆H₁₂ to another free but shorter Mn=O bond (bond length = 1.588 Å). The of TS1(AcOH) is 20.4 kcal mol⁻¹, which is 2.3 kcal mol⁻¹ lower than that of TS1(MnO₄⁻). The subsequent steps are similar to those in c-C₆H₁₂ oxidation by MnO₄⁻.

In the presence of three AcOH molecules, the AcOH molecules form hydrogen bonds with three Mn=O to give an intermediate, INT1(3AcOH) (Fig. 9). The bond lengths of the H-bonded Mn=O are slightly elongated from 1.591 to 1.601 Å. HAT then occurs from c-C₆H₁₂ to the free and more electrophilic Mn=O bond (bond length = 1.580 Å), with a = 16.8 kcal mol⁻¹. Compared with c-C₆H₁₂ oxidation by MnO₄⁻, the of HAT is further reduced by 5.9 kcal mol⁻¹.

Fig. 9 PES and structures for cyclohexane oxidation by [MnO₄(AcOH)₃]⁻ at the B3LYP-D3(BJ)/def2-SVPD level. The relative 298 K Gibbs free energies in acetonitrile are given in kcal mol⁻¹.

In the presence of [Ca(OTf)]₂ (Fig. 10), Ca(II) binds to two of the Mn=O bonds to form INT1(CaOTf), [MnO₄(CaOTf)⋯C₆H₁₂]. The two Mn=O bonds that bind to Ca(II) in INT1(CaOTf) are longer (1.611 and 1.615 Å) than the unbound Mn=O bonds (1.576 Å) in INT1(MnO₄⁻), as a result of the electron withdrawing effect of Ca(II). Subsequent HAT from c-C₆H₁₂ to an unbound Mn=O occurs with a of 12.7 kcal mol⁻¹via TS1(CaOTf), which is lower than that of MnO₄⁻ and MnO₄⁻/3AcOH by 10.0 and 4.1 kcal mol⁻¹, respectively; consistent with the experimental observation that the accelerating effect of Ca(II) is much higher than that of AcOH.

Fig. 10 PES and structures for cyclohexane oxidation by [MnO₄(CaOTf)] at the B3LYP-D3(BJ)/def2-SVPD level. The relative 298 K Gibbs free energies in acetonitrile are given in kcal mol⁻¹.

The activation barrier is further lowered when both [Ca(OTf)]⁺ and AcOH are present. In the presence of one molecule of [Ca(OTf)]⁺ and three molecules of AcOH, each AcOH is H-bonded to one Mn=O, while the [Ca(OTf)]⁺ is bonded to each AcOH and one Mn=O to form a relatively stable intermediate, INT1(CaOTf/3AcOH), as shown in Fig. 11. HAT then occurs from c-C₆H₁₂ to the free and shortest Mn=O bond with the lowest of 9.7 kcal mol⁻¹ ( = 11.5 kcal mol⁻¹ for one Ca(OTf)⁺ and one AcOH, Fig. S11†), in agreement with the observed synergistic activating effects of AcOH and Ca(II). It should be noted that in this case HAT results in the direct formation of the alkoxo intermediate without going through a cyclohexyl radical intermediate.

Fig. 11 PES and structures for cyclohexane oxidation by [MnO₄(CaOTf)(AcOH)₃] at the B3LYP-D3(BJ)/def2-SVPD level. The relative 298 K Gibbs free energies in acetonitrile are given in kcal mol⁻¹.

was also determined experimentally by performing kinetics studies at 20–50 °C. From the Arrhenius plot was found to be 10.4 ± 0.6 kcal mol⁻¹, in agreement with the calculated value (Fig. S12†).

We have also calculated the TS barrier heights for hydroxylation via a rebound mechanism and they are higher than those of the TS forming the alkoxo intermediate by 2–3.4 kcal mol⁻¹ (Table S1†).

Activating effects of Ca²⁺/AcOH on MnO₄⁻vs. RuO₄⁻

Although MnO₄⁻ and RuO₄⁻ are isostructural with similar E⁰ (0.56 V and 0.59 V, respectively), there is a great difference in the basicity of their oxo ligands; MnO₄⁻ is a weak base (pK_a of HMnO₄ is −2.25),³⁰ whereas RuO₄⁻ is a much stronger base. Although the pK_a of HRuO₄ is not known, RuO₄⁻ is readily protonated by weak acids such as AcOH to give [RuO₃(OH)(AcO)⁻], which readily undergoes HAT with c-C₆H₁₂ with a barrier that is lower than that of RuO₄⁻ by >10 kcal mol⁻¹ (Table 2).²⁵ In contrast, MnO₄⁻ is not protonated by AcOH. Nevertheless, AcOH can still exert a moderate activating effect on HAT by MnO₄⁻ through H-bonding with the oxo ligands, with the activation barrier lowered by 6 kcal mol⁻¹ compared with MnO₄⁻ alone. On the other hand, Ca²⁺ has a higher activating effect on MnO₄⁻ than on RuO₄⁻ (lowering of the barrier is 10 vs. 8 kcal mol⁻¹). Remarkably, for both complexes the presence of AcOH and Ca²⁺ together shows a similar synergistic activating effect, resulting in a lowering of the HAT barrier by 16 and 13 kcal mol⁻¹ for RuO₄⁻ and MnO₄⁻, respectively. DFT calculations show that Ca(II) forms a MO₄⁻/AcOH/Ca(II) active intermediate by binding to AcO⁻/AcOH and an oxo ligand (Fig. 12).

Table 2 Activation barrier (, kcal mol⁻¹) for the oxidation of c-C₆H₁₂ by RuO₄⁻ and MnO₄⁻ under various conditions

	RuO4^−	MnO4^−
No additive	26.8	22.7
AcOH	15.2	16.8
Ca^2+	18.5	12.7
AcOH + Ca^2+	10.8	9.7

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Fig. 12 Structures of the intermediates (a) RuO₄⁻/AcOH/Ca(OTf)₂ and (b) MnO₄⁻/AcOH/Ca(OTf)₂.

Conclusions

In conclusion, we have demonstrated synergistic activating effects of a weak Brønsted acid such as acetic acid and a weak Lewis acid such as Ca²⁺ on MnO₄⁻. Although MnO₄⁻ is rather non-basic, it can readily be activated by AcOH + Ca²⁺ toward HAT of c-C₆H₁₂, resulting in a lowering of the barrier by 13 kcal mol⁻¹. Remarkably, such a cooperative effect on MnO₄⁻ is similar to that on RuO₄⁻, despite a great difference in the basicity of the two metal-oxo species, which should lead to vastly different affinities toward acids. Our studies suggest that such a synergistic activation is a general phenomenon which may occur in various other metal-oxo species.

The combination of a weak Brønsted and a weak Lewis acid is a potentially useful strategy for the activation of a metal-oxo species toward oxidation of various organic compounds, particularly for substrates that may contain functional groups that are sensitive to strong acids. It is possible that this synergistic activating effect on metal-oxo species may also occur in biological systems, where only relatively mild Brønsted acids and Lewis acids (metal ions) are present.

Data availability

The datasets supporting this article have been uploaded as part of the ESI.†

Author contributions

C. K. M. designed and carried out the experiments, analysed the data and wrote the manuscript. T. C. L. designed the experiments and wrote the manuscript. H. S. performed the experiments and helped in the analysis of the data and preparation of the manuscript. L. C., Y. P. and K. C. L. performed computational studies and wrote the manuscript.

Conflicts of interest

There are no conflicts to declare.

Note added after first publication

This article replaces the version published on 26/09/2022, which contained errors in the position of the equilibrium constants in equations 1, 4, 7 and 8.

Acknowledgements

This work was supported by the NSFC/RGC Joint Research Scheme (N_CityU111/20) and the Laboratory for Synthetic Chemistry and Chemical Biology Limited, LSCCB. The computational studies were carried out using the computational facilities at Burgundy at City University of Hong Kong.

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Footnotes

† Electronic supplementary information (ESI) available. See https://doi.org/10.1039/d2sc03089f

‡ These authors have equal contribution.

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