Ligand substitution and electron transfer reactions of trans-(diaqua)(salen)manganese(III) with oxalate: an experimental and computational study

Abstract

The trans-Mn^III(salen)(OH₂)₂⁺ undergoes reversible aqua ligand substitution by HOX⁻ (H₂salen = N,N′-bis(salicylidene)ethane-1,2-diamine; HOX⁻ = ⁻O–COCO₂H) with k₁/dm³ mol⁻¹ s⁻¹ (k₋₁/s⁻¹) = 11.8 ± 0.7 (0.255 ± 0.02), ΔH^≠/kJ mol⁻¹ = 54.6 ± 0.8 (64.2 ± 6.7), ΔS^≠/J K⁻¹ mol⁻¹ = −41.2 ± 2.6 (−40.8 ± 22.7) at 25.0 °C and I = 0.3 mol dm⁻³. The low values of the activation enthalpy and nearly the same and negative values of the activation entropy are ascribed to an associative transition state for this interchange process (I_a mechanism). The redox reaction that follows involves several paths and the products are Mn^II and CO₂ identified by ESR spectroscopy and conventional test, respectively. The rate retardation by acrylamide monomer with no perceptible polymerization during the course of the redox reaction supports the involvement of the radical intermediate, C₂O₄⁻˙ (= CO₂ + CO₂⁻˙) which succeeds in reducing Mn^III species much faster than the dimerisation of its congener, CO₂⁻˙ in keeping with the stoichiometry, |[ΔMn^III]/Δ[OX]| = 2. The trans-[Mn^III(salen)(OH₂)(HOX) and its conjugate base, trans-Mn^III(salen)(OH₂)(OX)⁻ are virtually inert to intramolecular reduction of the Mn^III centre by the bound oxalate species but undergo facile electron transfer by H₂OX, HOX⁻ and very slowly by OX²⁻ following the reactivity sequence, k_H2OX > k_HOX ⋙ k_OX and featuring second order kinetics. The rate retardation by the anionic micelles of SDS (sodium dodecyl sulfate) and rate enhancement by N₃⁻ provide supportive evidence in favor of the proposed mechanistic pathways. The structure optimization of trans-Mn^III(salen)(OH₂)(HOX) (A), trans-Mn^III(salen)(HOX)₂⁻ (B), trans-Mn^III(salen)(OH₂)(OX)⁻ (C), trans-Mn^III(salen)(OH₂)(H₂OX)⁺ (E₁), and trans-Mn^III(salen)(HOX)(H₂OX) (E₂) {all high spin Mn^III(d⁴)} by Density Functional Theory (DFT) reveals that the structural trans-effect of the unidentately bonded OX²⁻ in C is the strongest and Mn^III assumes five coordination with the H₂O molecule (displaced from the Mn^III centre), hydrogen bonded to the phenoxide oxygen moiety. The computational study highlights different modes of H-bonding in structures A–E. The activation parameters for the redox reactions, A + HOX⁻ and A + H₂OX, ΔH^≠/kJ mol⁻¹ (ΔS^≠/J K⁻¹ mol⁻¹): 42.5 ± 6.2, (−106 ± 20) and 71.7 ± 7.7 (+12 ± 25), respectively, are indicative of different degrees of ordering and reorganization of bonds as expected in the case of a proton coupled electron transfer (PCET) process.

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1. Introduction

Mn^III complexes with a salen (H₂salen = N,N′-bis(salicylidene)ethane-1,2-diamine) motif are models for catalase and superoxide dismutase (SOD) activities.^1–4 Many intricate diseases, such as Alzheimer, cancer, multiple sclerosis, heart and brain stokes, are supposed to be related to the superoxide activity and Mn^III(salen) complexes have been shown as promising therapeutics.^1,3,5 The interaction of bioactive ROS (Reactive Oxygen Species) with the Mn^III centre of Mn^III(salen) moiety is a primary feature in such processes. Mn^III(salen) complexes have been used as catalysts in the epoxidation of olefins, and other organic transformations.^6–9 A review of the redox reactions of mononuclear Mn^III complexes, essentially of aminopolycarboxylate ligands, appeared several years ago which describes briefly the mechanistic aspects of such reactions.¹⁰ Nature's water oxidation catalyst (WOC) responsible for photo-catalyzed O₂ evolution from water is a tetra manganese cluster with Mn^III–Mn^IV coupled system which is believed to undergo redox cycling of the oxidation states of the Mn centres.^11–13 Recent investigations on the oxidation of oxalate, a ubiquitous moiety in the biological domain, by oxalate oxidase have revealed the importance of the Mn^III centre in catalyzing the oxidation of oxalate.^14,15 It involves an intermediate Mn^III–monooxalate complex in the monodentate form and electron transfer from the bound oxalate species to the Mn^III centre is considered to be rate limiting. Mn^III(salen) may be used to model this reaction. However, the studies devoted to understand the mechanisms of ligand substitution and redox reactions of Mn^III are limited. In view of the importance of oxalate in biology and Mn^III(salen) as a suitable oxidant for oxalate oxidation, we, in continuation of our other investigations^16,17 of the mechanism of ligand substitution and electron transfer at the Mn^III centre, present here a thorough study on trans-Mn^III(salen)(OH₂)₂⁺ + oxalate system over extended pH and temperature ranges. To the best of our knowledge there is no earlier literature report on such a study, although oxidation of oxalate by higher valent manganese has been the topic of several investigations in the past.^18,19

2 Experimental

2.1 Materials and reagents

Mn^III(salen)Cl·H₂O was received from our earlier work and synthesized when required and purity checked by elemental analysis and UV-Vis, IR spectra.¹⁷ This complex undergoes fast aquation to Mn^III(salen)(OH₂)₂⁺ when dissolved in water.^16,17 The optical spectrum of the diaqua complex in aqueous medium (pH 3) displays λ_max, nm (ε _max, dm³ mol⁻¹ cm⁻¹): 235(40 240), 279 (18 120) which agreed well with the previously reported values.¹⁷ The GR grade (E. Merck) oxalic acid (H₂OX), potassium oxalate, perchloric acid, sodium hydroxide and glacial acetic acid were used as received. All other reagents were of highest grade purity available. Freshly prepared doubly distilled water received from an all glass (borosilicate) distillation set was used to prepare the solutions; the second distillation of water was made through alkaline KMnO₄. NaClO₄ used for ionic strength adjustment was prepared by mixing requisite amounts of the standardized solutions of NaOH and HClO₄. Stock solution of NaClO₄ (1 mol dm⁻³) was prepared from time to time and adjusted to pH 6 and the concentration checked by a combined ion-exchange alkalimetric procedure using Dowex 50W X8 resin in the H⁺ form. The stock solution of the complex (5 × 10⁻³ mol dm⁻³, pH ∼ 5) was protected from light and stored in a refrigerator at ∼20 °C when not in use. It was not allowed to age for more than 24 h.

2.2 Physical measurements

A PerkinElmer Lambda25 and a Systronics (India) model 118 UV-visible spectrophotometers with a matched pair of 10 mm quartz cells were used for all absorbance measurements. The IR measurements were made on a PerkinElmer FTIR spectrometer, model Spectrum2 using KBr pellet. Fluorescence measurements were made on a JASCO spectrofluorimeter model FP-8200 using Xe/D₂ light sources; band width was set at 10 nm and scan speed was 100 nm min⁻¹. The excitation wavelength (λ_exct) was set at 265 nm and spectral scans covered 340–500 nm. The intensity of emission at any wavelength was normalized as where denotes intensity at the wavelength maximum in absence of SDS. The ESR measurement was performed on a JEOL (Japan) JES-FA 200 ESR spectrometer at room temperature operating in X-band mode (8.75–9.65 GHz, power 1.08 W, sensitivity 7 × 10⁹ spins/0.1 mT, resolution 2.35 μT). The pH measurements were made with a Systronics (India) pH meter model 335 using a glass–Ag/AgCl, Cl⁻ (3 mol dm⁻³ NaCl) electrode CL 51. NBS buffers of pH 4.01, 6.86 and 9.20 prepared from KHphthalate, Na₂HPO₄/KH₂PO₄, and Na₂B₄O₇·10H₂O respectively were used to calibrate the pH meter. The measured pH of the reaction medium was converted to p[H⁺] (= −log[H⁺]) established by a calibration curve using dilute HClO₄ solutions (1.98 × 10⁻² ≤ [H⁺]/mol dm⁻³ ≤ 1.00 × 10⁻⁵) at the same ionic strength as maintained in the reaction media (I = 0.3 mol dm⁻³).²⁰

2.3 Kinetics

The fast kinetics measurements were performed on a KinetAsyst SF-61 SX2 single mixing stopped flow spectrophotometer; the data acquisition and analysis was made using kinetic studio software version 0.94, application version 1.12 (TGK Scientific, U. K). The flow module and the mixing chamber were thermostatted to the desired temperature by circulating water from a refrigerated/heating water bath (Julabo F12-ED). One of the syringes was loaded with the solution of trans-Mn^III(salen)(OH₂)₂⁺ complex (in situ generated) while the other contained the desired mixture of oxalic acid, HClO₄ and NaClO₄ such that after mixing the final ionic strength was set at 0.3 mol dm⁻³. The available pK values of oxalic acid (see later) were used to calculate the ionic composition so as to set the ionic strength at the desired value. Rate measurements were made at 380 nm under pseudo-first order conditions. The absorbance versus time plots were biphasic over extended time scale (see Fig. S1(a and b)†) thus indicating the rapid formation of an intermediate followed by its decay at long time scale. The two processes were treated independently. The initial fast rise of absorbance with time was fitted to eqn (1) to get k^f_obs and A_eq. Data fitting using the software package (see above) for any individual run was within ±1%. At least 6–10 measurements were made for each run and k^f_obs could be reproduced within ±5% (σ(k^f_obs)/k^f_obs ≤ ± 0.05).

A_t = C₁ exp(−k^f_obst) + A_eq (1)

The values of A_eq were dependent on [OX]_T for a given [H⁺] and [complex]_T indicating that the formation of the intermediate was equilibrium controlled. A limited number of runs for the slow phase were made by stopped flow spectrophotometry wherever possible. The absorbance–time data for the slow phase of the reaction also studied under pseudo-first order conditions fitted to a single exponential equation with A_∞ close to zero (A_t = C₂ exp(−k^s_obst) + A_∞) and the corresponding rate constants (k^s_obs) were calculated. Most rate measurements for the slow reactions were conveniently made by batch sampling technique at 25.0–40.0 °C.¹⁷ The concentration of the complex, [Mn^III(salen)(OH₂)₂⁺] was varied as (0.6–1.22) × 10⁻⁴ mol dm⁻³ and that of [OX]_T (= total oxalic acid concentration) in the range 0.0005–0.1 mol dm⁻³. The ionic strength of the medium was fixed at 0.3 mol dm⁻³ (NaClO₄) unless otherwise quoted. The pH of the reaction mixtures was varied by self buffering due to H₂OX/HOX⁻ and HOX⁻/OX²⁻. The observed rate constants (k^s_obs) were calculated by fitting the absorbance (A_t)–time (t) data to a single exponential equation as mentioned above. A_∞ was close to zero for the completion of the reaction which was further verified by simulating the reaction mixture at complete reaction with appropriate solutions made out of Mn^II acetate, oxalic acid and other components at the same pH (for this the medium was 5% MeOH–water v/v as H₂salen was prepared in MeOH). The initial absorbance was in the range 0.4–0.6. For very slow reactions (k^s_obs ∼ 10⁻⁵–10⁻⁶ s⁻¹) the rate constants were evaluated by the method of initial rate as described earlier.¹⁷ σ(k^s_obs)/k^s_obs was generally better than ±2% while the same from the initial rate method was ∼±6%.

2.4 Initial fast reaction

The initial fast reaction was considered to be the reversible complexation of Mn^III(salen)(OH₂)₂⁺ with oxalate species. The k^f_obs values at 20–40 °C are collected in Table S1(a).† In the medium comprising HClO₄ and oxalic acid (see Table S1(a)†), [OX]_T is partitioned between H₂OX and HOX⁻ (pK₁ = 1.00–1.03, pK₂ = 3.59–3.64 for H₂OX at 20–40 °C, I = 0.3 mol dm⁻³ see Table S1(b)†).²¹

However, the concentrations of different oxalate species may be expressed as [HOX⁻] = f₁[OX]_T, [H₂OX] = f₂[OX]_T, and [OX²⁻] = f₃[OX]_T where f₁ = K₁[H⁺]/D, f₂ = [H⁺]²/D, f₃ = K₁K₂/D and D = [H⁺]² + K₁[H⁺] + K₁K₂. The values of [H⁺], [HOX⁻] and [H₂OX] were computed from the initial analytical values of [HClO₄] and [OX]_T considering the first stage acid dissociation of H₂OX, [H⁺] = [HClO₄] + X, X (= [HOX⁻]) being the acceptable solution of eqn (3), and [H₂OX] = [OX]_T − X.

X² + ([HClO₄] + K₁) − K₁[OX]_T = 0. (3)

The k^f_obs data were fitted to eqn (4) valid for Scheme 1 by a least squares computer program.

k^f_obs = (k^f₁f₁+ k^f₂f₂)([OX]_T + Q₁⁻¹/f₁) (4)

Scheme 1 Complexation of trans-Mn^III(salen)(OH₂)₂⁺ by HOX⁻/H₂OX.

In eqn (4) Q₁ denotes the equilibrium constant for the formation of Mn^III(salen)(HOX)(OH₂) (eqn (5)).

Q₁ = [Mn^III(salen)(HOX)(OH₂)]_eq/[Mn^III(salen)(OH₂)₂⁺]_eq[HOX⁻]_eq, (5)

f₁ and f₂ are the fractions of [OX]_T as HOX⁻ and H₂OX, and k^f₁ and k^f₂ are the second order rate constants for the formation of Mn^III(salen)(OH₂)(HOX) by HOX⁻ and H₂OX (see Scheme 1) respectively. The calculated rate, equilibrium and activation parameters are collected in Table S1(c).†

It turns out that k^f₂f₂ term is statistically insignificant at all temperatures. Thus neglecting k^f₂f₂ term and setting k^f₁Q₁⁻¹ = k^f₋₁ eqn (4) can be rearranged to eqn (6).

k^f_obs = k^f₁f₁[OX]_T + k^f₋₁ (6)

A representative plot at 25 °C (see Fig. 1) bears this fact. As a check the equilibrium absorbance data (A_e) from the stopped flow runs for a constant [complex]_T but varying [OX]_T and [H⁺] (see Table S1a†) are used to calculate Q₁ from the linear plots of 1/(A_e − A₀) versus 1/(f₁[OX]_T) (see eqn (7)); here A₀ and A_c denote the absorbances of Mn^III(salen)(OH₂)₂⁺ and Mn^III(salen)(HOX)(OH₂) respectively at the same total concentration of the complex.

1/(A_e − A₀) = 1/[Q₁(A_c − A₀) × f₁[OX]_T] + 1/(A_c − A₀). (7)

Fig. 1 Formation of [Mn^III(salen)(OH₂)(HOX)]. k^f_obs/s⁻¹ vs. 10²f₁[OX]_T/mol dm⁻³ plot at 25 °C.

The calculated values of Q₁ are 40.1 ± 2.0, 42.6 ± 1.0, 41.2 ± 2.3 and 40.6 ± 2.7 dm³ mol⁻¹ at 20.0°, 25.0°, 30.0° and 40.0 °C respectively which compare well with the values obtained from the kinetic data (see Table S1c,† foot note a). As Q₁ shows little variation with temperature its mean value from kinetic and equilibrium measurements (Q₁ = 42.0 ± 0.7 dm³ mol⁻¹ at 20–40 °C) is used for all other calculations. A comparison of the rate and activation parameters for the reversible formation of some Mn^III complexes is made in Table 1.

Table 1 Comparison of the rate and activation parameters for the formation/dissociation of some Mn^III complexes

Reaction: R = Mn^III(salen)	k^a (30 °C)	ΔH^≠/kJ mol^−1	ΔS^≠/J K^−1 mol^−1	Ref.
a Units: dm^3 mol^−1 s^−1 (s^−1) for the formation (dissociation) reactions. I = 0.3 mol dm^−3.b This work.c Ref. 16.d I = 0.2 mol dm^−3; hydroquinone (H2Q) and catechol (H2Cat) (ref. 22).e 28 °C; ascorbic acid (H2Asc) (ref. 33).f I = 0.25 mol dm^−3 (ref. 36).
R(OH2)2^+ + HOX^− → Mn^III(salen)(OH2)(HOX)	17.5 ± 1.1	54.6 ± 0.8	−41.2 ± 2.6	^b
R(OH2)(HOX) → Mn^III(salen)(OH2)2^+ + HOX^−	0.47 ± 0.07	64.2 ± 6.7	−40.8 ± 22.7	^b
R(OH2)2^+ + HSO3^− → Mn^III(salen)(OH2)(HSO3)	(3.0 ± 0.3) × 10^2	42.4 ± 0.2	−55.3 ± 0.6	^c
R(OH2)2^+ + SO3^2− → Mn^III(salen)(OH2)(SO3)^−	(1.10 ± 0.08) × 10^3	33.0 ± 3.0	−75 ± 10	^c
R(OH2)(OH) + SO3^2− → R(OH)(SO3)^2−	(2.1 ± 0.2) × 10^3	32.4 ± 0.3	−72.9 ± 0.6	^c
R(OH2)(OH) + H2Q → R(OH)(H2Q)	21.8 ± 0.5			^d
R(OH2)(OH) + HQ^− → R(OH)(HQ)^−	(1.4 ± 0.2) × 10^3			^d
R(OH2)(OH) + H2Cat → R(OH)(H2Cat)	1.91 ± 0.41			^d
R(OH2)(OH) + HCat^− → R(OH)(HCat)^−	(2.2 ± 0.3) × 10^2			^d
R(OH2)2^+ + H2Asc → R(OH2)(H2Asc)^+	1.2	72.6	+2.5	^e
Mn^III(EDTA)(OH2)^− + N3^− → Mn^III(EDTA)(N3)^2−	0.16	57.4 ± 0.9	−71.3 ± 2.9	^f
Mn^III(EDTA)(N3)^2− → Mn^III(EDTA)(OH2)^− + N3^−	4.7 × 10^−3	54.4 ± 0.7	−110 ± 3	^f

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2.5 Redox reaction

2.5.1 Product identification and stoichiometry. The time dependent spectral scans of the reaction mixture containing Mn^III complex and oxalic acid at pH = 4.17 is presented in Fig. S2.† Similar trend is observed at pH = 1.44.

The maxima around 285 and 390 nm (broad) characteristics of the parent complex is lost during the course of the reaction with the development of a maximum at 325 nm (broad). This is in good agreement with the spectrum of the mixture of Mn^II acetate + salen + H₂OX at the same pH and respective concentrations of the reactants, left to equilibration (see Experimental section). The maximum at 325 nm is, however, considerably reduced in intensity at lower pH. Acetate ion/acetic acid at low concentration (∼(1.2–2.4) × 10⁻⁴ mol dm⁻³) originating from Mn(OAc)₂ has no effect. The broad 325 nm peak is attributed to Mn^II(salen)(OH₂)(OX/OXH)^2−/− in equilibrium with Mn^II(salen)(OH₂)₂. The ESR spectrum of the spent reaction mixture displayed 6 line spectrum characteristic of Mn^II (see Fig. 2) thus establishing the Mn^III(salen)/oxalate redox reaction.

Fig. 2 ESR spectrum (X-band) of the product Mn^II in the reduction of trans-Mn^III(salen)(OH₂)₂⁺ by oxalate (pH 1.8, ∼27 °C), g = 1.9968. Intensity versus H/mT plot.

The formation of CO₂ was qualitatively established by the conventional test as follows. The reaction mixture ([Mn^III(salen)(OH₂)₂⁺]_T = 3 × 10⁻³ and [OX]_T = 0.02 mol dm⁻³, pH = 3.5, 40 °C) was set aside for 10t_1/2 and then treated with a slight excess CaCl₂ followed by aqueous NH₃ resulting in a white precipitate. This was allowed to coagulate, collected by filtration, air dried and treated with dilute HCl when a colorless gas with effervescence (characteristic of CO₂) evolved. Our attempts to perform a quantitative analysis of the unreacted oxalate by KMnO₄ titration in acid medium after precipitating it as CaC₂O₄ from ammoniacal solution was unsuccessful due to the presence of the salen ligand. Based on the identified products we propose the following stoichiometry:

2Mn^III(salen)(OH₂)₂⁺ + OXH⁻ + 3H⁺ = 2Mn^II + 2H₂salen + 2CO₂ + 2H₂O. (8)

Similar relation can be written for H₂OX or OX²⁻. It may be also noted that under mild acidic condition H₂salen undergoes hydrolysis to salicylaldehyde and bis-N-protonated ethylenediamine as the final end products.

2.5.2 Analysis of rate data for redox reaction. The rate data for the redox reaction are collected in Tables S2–S5.† A preliminary rate measurement in absence of oxalic acid but in acid medium at [HClO₄] ≤ 0.02 mol dm⁻³ (25–40 °C) indicated that the complex is significantly inert to the acid catalyzed decomposition. This is indicated by the constancy of the molar extinction coefficient of the complex over an extended time period, ε_{380 nm}/dm³ mol⁻¹ cm⁻¹ ([HClO₄]/mol dm⁻³): 5024 ± 5(0.01), 4930 ± 8 (0.02) at 25 °C for 4.63 h; 4970 ± 6 (0.01), 4944 ± 12 (0.01), 5123 ± 13 (0.01) at (30–40) °C for 4.5 h. However, there was a slow H⁺-catalyzed decomposition of the complex at [H⁺] ≥ 0.05 mol dm⁻³, 10⁵k_obs/s⁻¹ (t/°C); 0.32 ± 0.05(25), 0.70 ± 0.020(30), 1.28 ± 0.05(35), and 1.85 ± 0.04(40). Hence all our measurements for redox reaction was restricted to ca. [H⁺] ≤ 0.05 mol dm⁻³ at which the k^s_obs at the lowest [OX]_T (= 0.5 × 10⁻³ mol dm⁻³) was ≥12 times higher than the same for the H⁺-catalyzed decomposition of the complex. The correction of k^s_obs for the H⁺-catalyzed decomposition of the complex was insignificant and hence was neglected in the analysis of the rate data.

The k^s_obs versus [OX]_T plots at [HClO₄] = 0.05 mol dm⁻³ and pH = 3.10 ± 0.09 (see Fig. S3(a and b)†) are distinctly nonlinear. The observed trend shows greater than first-order dependence of k^s_obs on [OX]_T. Further k^s_obs at constant [OX]_T = 0.005 mol dm⁻³ approaches a low limiting value around pH 5 (∼10⁻⁶ s⁻¹, see Fig. S3c†) indicating that OX²⁻ is not an efficient reductant as HOX⁻ and H₂OX.

A limited number of runs at 35 °C were made at 5.04 ≤ pH ≤ 6.72, 0.005 ≤ [OX]_T mol⁻¹ dm⁻¹⁻³ ≤ 0.1 in order to establish the redox activity of trans-Mn^III(salen))(OH₂)(OX)⁻. Due to the solubility limitation of Na₂C₂O₄ the source of oxalate was K₂C₂O₄ and ionic strength was adjusted by KCl. The presence of chloride under the conditions had no perceptible effect on the absorption spectrum of the di-aqua complex discounting its interference. Under this condition [OX]_T = [HOX⁻] + [OX²⁻] and the diaqua Mn^III complex predominates (∼80%). Here again a nonlinear and greater than first order dependence of k^s_obs with [OX]_T at constant pH (= 5.27 ± 0.07) was observed (Fig. S3d†). Considering these facts Scheme 2 is proposed for the overall reaction for which k^s_obs (without considering the reaction stoichiometry; k_icorr = k_i/2, i = 0–7) is given by eqn (9),

where f_is (i = 1–3) are as defined earlier, k′= k₅Q₂ + k₇Q₁, and k_is (i = 0–7), Q₁, Q₂ (= Q₁K^′₂/K₂) and K^′₂ are the rate and equilibrium parameters respectively (see Scheme 2). The k^s_obs data at 1.26 ≤ pH ≤ 5.08 and low [OX]_T (see Tables S2–S4†) could be satisfactorily analyzed by eqn (10) a limiting form of eqn (9), as detailed below.

Scheme 2 Reduction of trans-Mn^III(salen)(OH₂)₂⁺ by oxalate.

The species, H₂OX, will not exist in significant concentrations under the pH conditions 2.9 ≤ pH ≤ 5.08. This amounts to a reasonable choice of neglecting the k₂ and k₃ terms of eqn (10) to treat the k^s_obs data at low [OX]_T. On this ground eqn (10) reduces to eqn (11):

The temperature independent value of Q₁ (= 42.0 dm³ mol⁻¹) was used, the fraction f₁ was calculated as mentioned above and the k^s_obs data in Tables S3 and S4† were analyzed by eqn (11) using a nonlinear least squares computer program assigning unit weight to each data point. The calculated values of k₀, k₁ and K^′₂ are collected in Table 2. The k₀ values turned out statistically insignificant. The inclusion of the k₂ term in the calculation did not improve the data fitting. A representative plot at 35 °C using a linearized form of eqn (11) (see Fig. 3) clarifies that contributions from k₀, k₂, and k₃ are statistically insignificant.

Table 2 Summary of the calculated values of the rate and equilibrium constants and activation parameters for the redox reaction

Rate constant^a	Temp./°C
a Rate constants (kis) are not corrected for the stoichiometry factor (i.e. kicorr = ki/2).b Calcd from the rate data in Table S2.c Calcd from the rate data in Tables S3 and S4.d Calcd from the temperature dependence of k1 and k2: k = (kBT/h)exp(−ΔH^≠/RT + ΔS^≠/R).
k0/s^−1	25.0 ± 0.1	30.0 ± 0.1	35.0 ± 0.1	40.0 ± 0.1
	0.05 ± 0.0015^b	0.0045 ± 0.0018^b	0.00002 ± 0.002^b	0.00049 ± 0.0029^b
	0.0000 ± 0.0005^c	0.000 ± 0.00027^c	0.000042 ± 0.00025^c	0.00019 ± 0.00092^c
k1/dm^3 mol^−1 s^−1	0.73 ± 0.12	0.87 ± 0.07	1.16 ± 0.07	1.73 ± 0.24
k2/dm^3 mol^−1 s^−1	7.39 ± 1.03	10.4 ± 1.2	19.6 ± 1.4	26.7 ± 4.6
k3/dm^3 mol^−1 s^−1	0.00017 ± 0.27	0.00093 ± 0.31	0.13 ± 0.47	0.00013 ± 0.47
10^6k4/s^−1			0.03 ± 14.0
k5/dm^3 mol^−1 s^−1			0.077 ± 0.047
10^4 k6/dm^3 mol^−1 s^−1			2.94 ± 2.57
k7/dm^3 mol^−1 s^−1			0.077 ± 0.047
k8/dm^3 mol^−1 s^−1				1.08 ± 0.09
10^4K2′/mol dm^−3	7.14 ± 2.21	4.62 ± 0.54	2.56 ± 0.31	3.52 ± 0.78
Q1/dm^3 mol^−1	42.0	42.0	42.0	42.0
ΔH^≠/kJ mol^−1d	k1 path, 42.5 ± 6.2	k2 path, 71.7 ± 7.7
ΔS^≠/J K^−1 mol^−1^d	−106 ± 20	+11.7 ± 25.4

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Fig. 3 10³k^s_obs(1 + Q₁f₁[OX]_T(1 + K^′₂/[H⁺])/(Q₁f₁[OX]_T) (= F_c/s⁻¹) vs. 10³f₁[OX]_T/mol dm⁻³ plot at 35 °C.

The k^s_obs data in Table S2† were then analyzed by eqn (10) using the values of k₁, K^′₂ (see Table 2) and Q₁. The calculated values of k₀, k₂ and k₃ are also collected in Table 2. Here again, the k₀ and k₃ values turned out statistically insignificant.

At 5.04 ≤ pH ≤ 6.72, and 0.0025 ≤ [OX]_T/mol dm⁻³ ≤ 0.1, OX²⁻ and HOX⁻ were considered to be the reducing species. The first step acid dissociation equilibrium of trans-Mn^III(salen)(OH₂)₂⁺ (pK_M ≥ 7.3 (ref. 16 and 22) was also taken in to account. eqn (9) was recast as eqn (12),

and k^s_obs values in Table S5† were analyzed by eqn (12) with the known values of Q₁, k₁ and the calculated value of Q₂ (= Q₁K^′₂/K₂); pK_M was varied between 6.8–7.3. The calculated values of k₄, k′ (= k₅Q₂ + k₇Q₁) and k₆ turned to be little sensitive to pK_M.

2.5.3 Effect of ionic strength. The ionic strength was varied as 0.01 ≤ I/mol dm⁻³ ≤ 0.3 at 30 °C keeping pH = 3.73 ± 0.06 and [OX]_T = 0.006 mol dm⁻³. The values of k^s_obs did not change significantly with the variation of ionic strength and averaged to (3.52 ± 0.09) × 10⁻⁴ s⁻¹. This is as expected from the consideration that under this condition, the redox process is essentially driven by the reaction between the uncharged Mn^III complex and the anion, HOX⁻ (i.e. Mn^III(salen)(OH₂)(HOX)⁰ + HOX⁻, see Scheme 2). No detailed analysis of the ionic strength effect was further attempted considering the complexity of the reaction.

2.5.4 Effect of acrylamide. The oxidation of oxalate by Mn^III and its complexes including Mn^IIIOX⁺ is known to generate the oxalate radical, C₂O₄⁻˙ which decomposes in a fast step to yield the radical CO₂⁻˙ and CO₂.^18,23 Recently the computer simulation of the complex oxalic acid permanganate reaction which involve the oxalate complexes of Mn^III unearthed the less known crucial role of the radical CO₂⁻˙.²⁴ The dimerisation of CO₂⁻˙ is also fast and regenerates oxalate. If this is a major step of the loss of C₂O₄⁻˙ then the overall stoichiometry of the reaction of Mn^III complex with OX²⁻/HOX⁻/H₂OX will be 1 : 1. On the other hand, the overall stoichiometry 2 : 1 for Mn^III/OX_T reaction should result if Mn^III complex scavenges the radical (C₂O₄⁻˙/CO₂⁻˙) at a rate much faster than the rate of its dimerisation. Keeping that in mind we investigated the effect of acrylamide monomer, a good scavenger of the radical, on the kinetics of the reaction under study at 30 °C with [Mn^III(salen)(OH₂)₂⁺]_T = 1.21 × 10⁻⁴, [OX]_T = 2.5 × 10⁻³, I = 0.3 mol dm⁻³ and pH = 3.22. The anticipated polymer formation (through visual inspection) could not be observed. Interestingly the values of k^s_obs showed a decreasing trend with the increase of [acrylamide]_T (10⁴k^s_obs/s⁻¹ = 1.67 ± 0.02, 1.63 ± 0.02, 1.61 ± 0.02 and 1.57 ± 0.02 at [acrylamide]_T(monomer) = 0, 0.01, 0.02, 0.03 mol dm⁻³ respectively). This indirectly led us to believe that the Mn^III complexes compete very successfully and efficiently in scavenging the C₂O₄⁻˙ or CO₂⁻˙ generated in the redox process maintaining the stoichiometry (= |Δ[Mn^III]|/|Δ[OX]|): 2 : 1, and dimerisation of CO₂⁻˙ at high dilution is of little significance; the possible reason may be the role played by the Coulombic repulsion between the anionic radicals in relation to the favorable electrostatic and non covalent interactions (H-bonding) between the radical anion and Mn^III(salen)(OH₂)₂⁺ and Mn^III(salen)(OH₂)(OXH).

2.5.5 Effect of surfactant, sodium dodecyl sulphate (SDS). The UV-Vis absorption spectra of trans-Mn^III(salen)(OH₂)₂⁺ in the presence of varying [SDS] ([complex]_T = 5.70 × 10⁻⁵, [HClO₄] = 1.0 × 10⁻⁴, [SDS]_T = 0, 0.01, 0.02, 0.05, 0.10 mol dm⁻³, 260 ≤ λ, nm ≤ 500, see Fig. S4a†) exhibits a small red shift (∼2 nm) and are virtually super imposable except for [SDS] = 0 (the spectral measurements at [SDS]_T = 0.002–0.006 could not be made due to appearance of a silky white precipitate). Only a small reduction in intensity (≤4% decrease of absorbance at 280 nm) is observed due to the presence of the surfactant. The fluorescence spectra of Mn^III(salen)(OH₂)₂⁺ in aqueous SDS media ([SDS]_T/mol dm⁻³ = 0, 0.02, 0.03, 0.05) is presented in Fig. S4(b).† The emission peak is observed at 415 nm in absence of SDS; it is enhanced with the increase of [SDS] indicating micellar binding of Mn^III(salen)(OH₂)₂⁺. The observed trends in emission and absorption behavior of Mn^III(salen)(OH₂)₂⁺ are reconciled with the partitioning of this cationic complex from the bulk aqueous phase to the dominantly water rich region of the micellar surface.

The rate data (k^s_obs, 30 °C) at constant [OX]_T = 0.006 mol dm⁻³, pH = 3.97 ± 0.03 and 0.00 ≤ [SDS]_T/mol dm⁻³ ≤ 0.10 are collected in Table S6.† There is no significant effect of [SDS] on the rate constant at [SDS]_T ≤ 0.0075 mol dm⁻³; marked retardation is observed beyond [SDS]_T = 0.0075 mol dm⁻³ and k^s_obs tends to attain a low limiting value at high [SDS]. This is reconciled with the fact that the diaqua complex and presumably the neutral complex, trans-Mn^III(salen)(OH)₂(OXH), are partitioned into the micellar pseudo phase of SDS and this protects the Mn^III species from electron transfer involving HOX⁻ which on electrostatic ground exists exclusively in the aqueous pseudo phase. The pH condition is such that [OX]_T exists in equilibrium in the bulk aqueous phase as HOX⁻ and OX²⁻. Interestingly there is substantial rate enhancement on increasing [Na⁺]_T by addition of NaClO₄ (= 0.01–0.3, see Table S6†) at a fixed [SDS]_T = 0.02 which also leads to the decrease of pH. The observed pH perturbation is a clear indication of the prevalence of ion exchange equilibrium involving H⁺/Na⁺ between the anionic micellar pseudo phase and the bulk aqueous phase. We consider the participation of the mono cationic Mn^III complex in this ion exchange equilibrium along with the equilibrium partitioning of the corresponding neutral HOX complex between the bulk aqueous phase and micellar pseudo phase. Scheme 3 is presented to interpret the rate data. Accordingly,

where Q⁰₁, K^0′₂ are the corresponding terms at zero ionic strength, y₊ is the activity coefficient of a mono valent cation, [Na_M⁺] = β[Dn] − [H_M⁺], [H_M⁺] = [H⁺]_T − [H_W⁺] − [OX]_T[H_W⁺]/([H_W⁺] + K⁰₂/y₂), y₂ = activity coefficient of OX⁻², [Dn] = [SDS]_T − cmc, cmc denotes critical micelle concentration, the subscripts M and W stand for the micellar pseudo-phase and bulk aqueous phase respectively (the micellar binding parameter β is approximated to: β = ([Na_M⁺] + [H_M⁺])/[Dn] as [Na_M⁺] + [H_M⁺] ≥ [R⁺]_M + ROXH_M), R⁺ = Mn^III(salen)(OH₂)₂⁺. The ionic strength (I) of the aqueous pseudo phase,

I = 3[Na₂OX]_T +[NaHOX]_T + [NaClO₄] + cmc + 0.5(1 − β)[[SDS]_T − cmc)

was calculated neglecting the contribution from [Mn^III(salen)(OH₂)₂]⁺_W] ([complex]_T = 1.0–1.2 × 10⁻⁴ mol dm⁻³). The activity coefficient at a given ionic strength was calculated using the relationship,

log y_z+ = −0.5Z²(I^1/2/(1 + I^1/2)) + 0.2I (14)

where Z denotes the charge of the ionic species. Since the solution compositions had variable ionic strength, the values of the equilibrium parameters (Q₁, K₂, K^′₂) used in the data fitting were corrected to zero ionic strength. The rate data (k^s_obs) at 0.01 ≤ [SDS]_T/mol dm⁻³ ≤ 0.1 were analyzed by eqn (13) with k_1W = k₁, pK⁰₂ = 4.27,²¹ Q⁰₁ = 72.7 (= Q₁(I = 0.3)/y₊² for y₊ = y₋) by varying cmc and β.

Scheme 3 Mn^III(salen)(OH₂)₂⁺ reduction by HOX⁻ in the presence of SDS micelles, R⁺ = Mn^III(salen)(OH₂)₂⁺, Dn = [SDS]_T − cmc.

The best fit values of K_ex, β and cmc turned out as 10.4 ± 0.8, 0.6 and 0.009 (cmc = 0.008 mol dm⁻³ at 25 °C, in absence of additives)²⁵ respectively while the micellar binding constant Q_m for the partitioning of the neutral species, Mn^III(salen)(OH₂)(HOX), is statistically insignificant (see Tables S6 and S7†).

2.5.6 Effect of azide (N₃⁻). The effect of azide ion on the redox reaction was studied at 40 °C with 0.01 ≤ [NaN₃]_T/mol dm⁻³ ≤ 0.2 at fixed [OX]_T = 0.022, I = 0.3 mol dm⁻³ and constant pH = 4.53 ± 0.03. k^s_obs increases with [N₃⁻]_T tending to attain a limiting value (see Table S8†).

This trend is explicable in terms of the equilibrium formation of trans-Mn^III(salen)(OH₂)(N₃) competitive with the corresponding HOX⁻ complex followed by the reduction of the azido complex by oxalate species as given below. Accordingly k^s_obs is given by eqn (15)

where G = 1.0 + Q₁[HOX⁻](1 + K^′₂/[H⁺]), [N₃⁻] = K_d[N₃⁻]_T/([H⁺] + K_d) (pK_d = 4.35 at 25 °C, I = 0.3 mol dm⁻³ for N₃H)²⁶ and k^′_obs = k^s_obs in absence of azide. The k^s_obs data fitted eqn (15) well and the values of k₈ and Q₄ are collected in Table S8.† As a check a preliminary rate measurement for the reduction of Mn^III(salen)(OH₂)₂⁺ by N₃⁻ ([N₃⁻]_T/mol dm⁻³ = 0.01, 0.10 and 0.2) indicated that such reaction at 40 °C, pH = 4.46 ± 0.06 (I = 0.3 mol dm⁻³) was extremely slow (see absorbance versus time plots in Fig. S5†) yielding k^s_obs = (7.7 ± 0.3) × 10⁻⁶ s⁻¹ for [N₃⁻]_T = 0.2 mol dm⁻³.

2.5.7 Molecular modelling and structure optimization. All calculations were performed with the program package TURBOMOLE 6.4 using density functional theory (DFT).^27,28 The BP86 functional and def2-TZVPP basis set together with the resolution-of-the-identity (RI) approximation^29–31 (RI-BP86/def2-TZVPP in short) was employed for the structure optimization procedure. Numerical frequency calculations of the optimized structures were done to ensure that the optimized structures were true minima not the transition states. The “freeh” script of turbomole was used to calculate the free energies of the complexes at 25 °C (298 K) and 1 atmospheric pressure. For the graphical presentation and the bond distance and angle measurements Mercury 3.0 was used.³²

3.1 Results and discussion

3.1.1 Formation/dissociation of trans-Mn^III(salen)(OH₂)(HOX). The kinetic data could not detect outer sphere association of the diaqua Mn^III complex with HOX⁻ or H₂OX. Also H₂OX is not an effective reactant for the aqua ligand substitution (k^f₂ ∼ 0, see Table S1c†) at the Mn^III centre. Similar observation has been reported for hydroquinone/catechol (H₂Q/H₂Cat).²² However, such a path has been detected for ascorbic. The comparison made in Table 1 reflects the dependence of the rate constants on the nature of the incoming ligand and the electrostatic interaction between the reacting partners. The rate constants may be contrasted with the water exchange rate constant of [Mn^III(OH₂)₆]³⁺, k_ex ∼ 10⁵ s⁻¹.³⁴ The labilizing effect of the coordinated hydroxide (see Table 1) is not large. Similar behaviour has been recently reported for the water exchange reactions of trans-(diaqua)/(aqua/hydroxo)Mn^III(porphyrins) (at 25 °C values of k_ex/s⁻¹, ΔH^≠/kJ mol⁻¹, ΔS^≠/J K⁻¹ mol⁻¹ for trans-[Mn^III(TPPS)(OH₂)₂]³⁻ and its (aqua)(hydroxo) isomer are: (1.4 ± 0.1) × 10⁷, 32.7 ± 1.1, 1.65 ± 3.0 and (2.5 ± 0.7) × 10⁷, 24.1 ± 2.0, −22.9 ± 9.9 respectively).³⁵ The ΔH^≠ values for the formation and dissociation of HOX complex are similar in magnitude (ΔΔH^≠(k^f₁ − k^f₋₁) = −9.6 ± 6.7 kJ mol⁻¹). This is expected as the bond that is broken and reformed in the process are alike (i.e. Mn^III–O). However, a higher value of ΔH^≠ for the reverse reaction (k^f₋₁) may be reconciled with the relatively stronger columbic interaction of the outgoing HOX⁻ with the Mn^III centre. Such a trend is not seen in the activation enthalpy data for the reversible formation of the anionic azido complex, Mn^III(EDTA)(N₃)²⁻ (see Table 1) reported by Suwyn and Hamm.³⁶ Strikingly the activation entropies for the formation and dissociation of trans-[(OH₂)Mn^III(salen)(HOX)] (see Table 1) are virtually the same and moderately negative; similar trend is also observed for Mn^III(EDTA)N₃²⁻. This leads us to suggest an associative transition state (ts, I_a mechanism) involved in the aqua ligand substitution at the Mn^III centre.

The seven coordinate Mn^III centre in the transition state does not seem to be unusual considering the seven coordinate complexes, Mn^III(EDTA)(OH₂)⁻ and Mn^III(EDTA)(N₃)²⁻ reported by Suwyn and Hamm.³⁶

3.2.2 Redox reaction. In the context of the redox reaction it is pertinent to discuss the DFT optimized structures of different oxalato complexes and their structural parameters. Fig. 4 depicts the structures of trans-Mn^III(salen)(OH₂)(HOX) (A), Mn^III(salen)(HOX)₂⁻ (B) and Mn^III(salen)(OH₂)(OX)⁻ (C₁ and C₂, the two isomeric forms with respect to the disposition of H₂O).

Fig. 4 RI-BP86/def2-TZVPP optimized structures of trans-Mn^III(salen)(OH₂)(HOX) (A), Mn^III(salen)(HOX)₂⁻ (B), and Mn^III(salen)(OH₂)(OX)⁻ (C₁, C₂), HOX⁻(OX²⁻) denotes ⁻O–C(=O)CO₂H/⁻O–C(=O)CO₂⁻.

Relevant bond distances and bond angles are collected in Table 3. In all these structures, the (N₂, O₂) square plane around Mn^III suffers little distortion while the axial Y–Mn^III–X bonds are strongly distorted and also differently for A, B, and C₁, C₂. For example the Mn–O (38O–COCO₂H) bond in A is 0.177 Å shorter than the same in B. Also the Mn–O (43 W) bond for the trans H₂O in A is substantially elongated (2.895 Å). The free carboxyl group(s) in both A and B are bridged to the boned carboxylate group(s) through hydrogen bonds with a relatively shorter H-bond length (δ_{(H-bond)B–A} = 0.083 Å) for A. This might be traced to the relatively stronger binding of HOX⁻ to Mn^III in A in tune with comparatively shorter Mn–O bond length (38O–COCO₂H) in this case. In all such structures Mn^III is situated above the partially distorted (N₂, O₂) square frame of the salen motif. The computed structures of trans-Mn^III(salen)(OH₂)(HOX) with possible H-bonding of the O–H function of the unbound carboxyl group are depicted in D₁–D₄ (Fig. 5). It turns out that D₁ with a long hydrogen bond (O⋯H bond 2.831 Å) between the phenoxide oxygen and the free carboxyl group is energetically unfavourable; D₂ (same as A in Fig. 4) is the most acceptable one.

Table 3 Selected bond distances (Å) and bond angles (degree) for RI-BP86/def2-TVZPP optimized structures of trans-Mn^III(salen)(OH₂)(HOX) (A), Mn^III(salen)(HOX)₂⁻ (B) and Mn^III(salen)(OH₂)(OX)⁻ (C₁, C₂)

A	B	C1	C2
Bond distances (Å)
Mn–O(2 phenoxide) 1.909	Mn–O(2 phenoxide) 1.908	Mn–O(2 phenoxide) 1.913	Mn–O(2 phenoxide) 1.900
Mn–O(3 phenoxide) 1.895	Mn–O(3 phenoxide) 1.908	Mn–O(3 phenoxide) 1.921	Mn–O(3 phenoxide) 1.909
Mn–O(43 W) 2.895	Mn–O(43 O–COCO2H) 2.257	Mn–O(42 W) too long	Mn–O(42 W) too long
Mn–O(38 O–COCO2H) 2.080	Mn–O(38 O–COCO2H) 2.257	Mn–O(38 O–COCO2^−) 2.020	Mn–O(38 O–COCO2^−) 2.003
Mn–N(4 imine) 1.992	Mn–N(4 imine) 2.00	Mn–N(4 imine) 1.984	Mn–N(4 imine) 1.993
Mn–N(5 imine) 1.990	Mn–N(5 imine) 2.00	Mn–N(5 imine) 1.963	Mn–N(5 imine) 1.967
O(38)⋯H(42) 1.867	O(38)⋯H(42) 1.784, O(43)⋯H(49) 1.784	O(2 phenoxide)⋯H(43 W), 2.081, O(3 phenoxide)⋯H(44 W), 2.116	O(3 phenoxide)⋯H(44 W) too long, O(2 phenoxide)⋯H(44 W), 2.343, O(41 O–COCO2)⋯H(43 W) 1.815
Bond angles (degree)
O(2)–Mn–O(38) 95.1	O(2)–Mn–O(38) 89.2	O(2)–Mn–O(38) 96.7	O(2)–Mn–O(38) 95.9
O(2)–Mn–O(3) 93.7	O(2)–Mn–O(3) 95.3	O(2)–Mn–O(3) 93.1	O(2)–Mn–O(3) 92.5
O(2)–Mn–N(4) 89.6	O(2)–Mn–N(4) 91.1	O(2)–Mn–N(4) 89.9	O(2)–Mn–N(4) 90.5
O(3)–Mn–N(5) 90.8	O(3)–Mn–N(5) 91.1	O(3)–Mn–N(5) 90.3	O(3)–Mn–N(5) 90.3
O(2)–Mn–N(5) 164.8	O(2)–Mn–N(5) 173.5	O(2)–Mn–N(5) 166.7	O(2)–Mn–N(5) 167.5
O(3)–Mn–N(4) 163.9	O(3)–Mn–N(4) 173.5	O(3)–Mn–N(4) 159.9	O(3)–Mn–N(4) 157.3
O(3)–Mn–O(43) 81.6	O(3)–Mn–O(43) 89.2
O(2)–Mn–O(43) 71.0	O(2)–Mn–O(43) 91.7
N(4)–Mn–N(5) 82.3	N(4)–Mn–N(5) 82.5	N(4)–Mn–N(5) 82.6
N(4)–Mn–O(38) 97.1	N(4)–Mn–O(38) 89.5	N(4)–Mn–O(38) 96.5	N(4)–Mn–O(38) 97.6
N(4)–Mn–O(43) 84.6	N(4)–Mn–O(43) 89.4
N(5)–Mn–O(38) 98.7	N(5)–Mn–O(38) 89.4	N(5)–Mn–O(38) 95.0	N(5)–Mn–O(38) 95.2
N(5)–Mn–O(43) 95.3	N(5)–Mn–O(43) 89.5
Mn–O(2)–C(6) 126.1	Mn–O(2)–C(6) 129.3	Mn–O(2)–C(6) 129.3	Mn–O(2)–C(6) 129.8
Mn–O(3)–C(21) 131.2	Mn–O(3)–C(21) 129.3	Mn–O(3)–C(21) 130.9	Mn–O(3)–C(21) 131.2
Mn–N(4)–C(12) 124.7	Mn–N(4)–C(12) 125.7	Mn–N(4)–C(12) 126.2	Mn–N(4)–C(12) 126.1
Mn–N(5)–C(15) 126.6	Mn–N(5)–C(15) 125.7	Mn–N(5)–C(15) 127.8	Mn–N(5)–C(15) 127.8

---

Fig. 5 RI-BP86/def2-TZVPP optimized structures of trans-Mn^III(salen)(OH₂)(HOX), D₁–D₄ denote the hydrogen bonded conformational isomers.

Notable is the fact that OX²⁻ coordination to Mn^III leads to very strong axial distortion resulting in expulsion of the bound H₂O from the Mn^III centre (C₁, C₂). Essentially Mn^III in trans-Mn^III(salen)(OH₂)(OX)⁻ assumes five coordination. Fig. 6 depicts the computed structures of trans-Mn^III(salen)(OH₂)(H₂OX)⁺ (E₁), Mn^III(salen)(HOX)(H₂OX) (E₂), and Mn^III(salen)(OX)(H₂OX)⁻ (E₃). In E₁ oxalic acid (H₂OX) molecule is bound to Mn^III centre by C=O function with long Mn–O bond (2.430 Å); one H-bond between the phenoxide group and the carboxyl proton (H⋯O bond 1.707 Å) of the bound carboxyl group completes a six membered ring structure imparting stability. It may be noted that both the M–O bond lengths trans to each other do not differ significantly although they are assembled differently (H₂O–Mn–O=C–). In E₂ the hydrogen bond length between the phenoxide oxygen and the carboxylic proton decreases to 1.627 Å with elongation of the Mn^III–O bond of Mn^III–O=C– to 3.090 and concurrent compression of the trans-Mn^III–O (Mn^III–OXH) to 2.068 Å. Interestingly H₂OX in E₃ is H-bonded to the phenoxide group (O⋯H = 1.465 Å) with a long Mn^III–O distance (4.337 Å) of Mn^III–O=C (no bonding situation) with simultaneous shortening of the trans-Mn–O (1.986 Å) reflecting the strong trans effect of the coordinated OX²⁻. Although these computationally accessible species, Mn^III(salen)(X)(H₂OX)⁺ (X = H₂O, HOX⁻ and OX²⁻) are not experimentally identified in aqueous medium, they provide sufficient insight to understand the reaction mechanism. Thus the computational study shows that there is a possibility of inner sphere reduction of Mn^III by HOX⁻ and H₂OX in the reactions of Mn^III(salen)(OH₂)(HOX) with HOX⁻ and H₂OX (k₁ and k₂ paths). The corresponding reactions between trans-Mn^III(salen)(OH₂)(OX)⁻ and HOX⁻ (k₅ path), trans-Mn^III(salen)(OH₂)(OX) and OX²⁻ (k₆ path), and trans-Mn^III(salen)(OH₂)(HOX) + OX²⁻ (k₇ path) in all probability be considered as outer sphere type as coordination of OX²⁻ to Mn^III will lead to the displacement of the trans-ligand (C₁ and C₂) due to the predominant structural trans effect (STE) of OX²⁻.

Fig. 6 RI-BP86/def2-TZVPP optimized structures of trans-Mn^III(salen)(OH₂)(H₂OX)⁺ (E₁), Mn^III(salen)(HOX)(H₂OX) (E₂), and Mn^III(salen)(OX)(H₂OX) (E₃); H₂OX is oxalic acid molecule, and HOX⁻, OX²⁻ denote its mono- and di-anions respectively.

We observe that the intramolecular electron transfer reactions from the bound oxalate to Mn^III centre in trans-Mn^III(salen)(OH₂)(HOX) (k₀) and Mn^III(salen)(OH₂)(OX)⁻ (k₄) are too slow, if at all occur, under the experimental conditions. This is what we observed in an earlier study of the reduction of Mn^III by the bound glyoxylate (gem diol form) in trans-Mn^III(salen)(OH₂)(O₂CH(OH)₂) (10⁵k₀/s⁻¹ = 0.05 ± 0.53 at 45 °C).¹⁷ For such complexes the values of their formation equilibrium constant {Q₁/mol dm⁻³ = 33.2–25.0, 42.0, 49.5 (25–40 °C, I/mol dm⁻³ = 0.3) for CH(OH)₂CO₂⁻, HOX⁻ and OX²⁻ respectively} and the pK′ data (−log K′ = 3.38 ± 0.20 see Table 2) of Mn^III(salen)(OH₂)(HOX) suggests that the carboxylate ligand binds the Mn^III centre in a monodentate fashion; this is further supported on a comparative basis by a value of pK = 2.0 (28 °C, I = 0.3 mol dm⁻³) for the unbound carboxyl group of bioxalatopentaamminecobalt(III), (NH₃)₅CoOCOCO₂H²⁺.³⁷ Despite this fact the carboxylate ligand binding results in substantial stabilization of the Mn^III state towards intramolecular redox process. The DFT calculations indicate that there is a finite possibility of a 1 : 1 species, Mn^III(salen)(H₂O)(H₂OX)⁺. We, however, did not observe [H₂OX] dependent rate of complexation of Mn^III(salen)(OH₂)₂⁺. Also the rate constant of reduction of Mn^III(salen)(OH₂)₂⁺ by H₂OX (k₃) turned out statistically insignificant. The possible reason for not being able to detect Mn^III(salen)(H₂O)(H₂OX)⁺ is due to the very low pK₁ of oxalic acid (pK₁ = 1.0–1.04 at 20–40 °C). Further lowering of the pK₁ of the coordinated oxalic acid favours complete H⁺ dissociation under the experimental conditions and in consequence Mn^III(salen)(OH₂)(H₂OX)⁺, if formed, gets converted to its HOX analogue. Presumably the experimental identification of Mn^III(salen)(OH₂)(H₂OX)⁺ in aqueous medium demands still higher [H⁺] and [H₂OX]. However, the acid catalyzed decomposition of Mn^III(salen)(OH₂)₂⁺ at high [H⁺] set a limit to our investigation to [H⁺] ≤ 0.05 mol dm⁻³.

The predominant reactions are the reduction of Mn^III in Mn^III(salen)(OH₂)(HOX) by HOX⁻ (k₁ path) and H₂OX (k₂ path), H₂OX being a better reducing agent than HOX⁻ (k₂/k₁ ≥ 10 at 25–40 °C). This trend appears surprising as HOX⁻ is a better electron donor than its conjugated acid, H₂OX. To account for this difference we suggest that the reaction in the k₂ path is preferably an efficient inner-sphere electro-protic reaction involving proton coupled electron transfer, PCET, (Fig. 7) or concerted proton electron transfer, CPET, process as observed in the oxidation of aromatic phenols by Mn^III(OH)(dpaq)⁺ (dpaq = 2-[bis(pyridine-2-ylmethyl)]amino-N-quinolin-8-yl-acetamidate) in MeCN.³⁸

Fig. 7 Probable mechanism of proton coupled electron transfer (PCET) involved in k₁ and k₂ paths.

Substantially large differences in the activation enthalpy (ΔΔH^≠_(k2−k1) = 29.2 ± 9.3 kJ mol⁻¹) and entropy (ΔΔS^≠_(k2−k1) = +117.5 ± 32.4 J K⁻¹ mol⁻¹) are observed for the two paths. The observed activation enthalpy arises due to (i) the enthalpy changes associated with the pre-equilibria, and (ii) energy demand for the rearrangement of bonds in the activation process. Out of these two contributing factors the later will predominate as the nature of the bond broken and reformed in the equilibrium pre-association are alike (Mn^III–O). However, the ΔS^≠ term depends on solvation of the reacting species in the initial states and the transition states. In the present case the substantially negative value of the activation entropy for k₁ path might indicate the involvement of an ordered transition state. There is likely to be much greater degree of rearrangement of bonds in the k₂ path (see Fig. 7) which accounts for the relatively higher values of ΔH^≠ and ΔS^≠ for this path. A simple calculation of the thermodynamic parameters for the k₁ and k₂ paths, [X(TS^≠₂) − X(TS^≠₁)], where X = G, H, S and TS^≠₂ and TS^≠₁ denote the activated states for the electron transfer respectively might clarify this picture. The differential values of the thermodynamic parameters can be expressed as,

G(TS^≠₂) − G(TS^≠₁) = −2.303RT[pK₁ + log k₂/k₁] + G⁰(H⁺)

H(TS^≠₂) − H(TS^≠₁) = ΔH^≠(k₂) − ΔH^≠(k₁) − ΔH⁰(K₁) + H⁰(H⁺)

S(TS^≠₂) − S(TS^≠₁) = ΔS^≠(k₂) − ΔS^≠(k₁) − T⁻¹ΔH⁰(K₁) + 2.303RpK₁ + S⁰(H⁺)

where X⁰(H⁺) (X = G, H, S) denotes the thermodynamic function of H⁺ in the standard state, and all other terms have their usual meaning. Using pK₁ = 1.252, ΔH⁰(K₁) = 3.76 ± 0.42 kJ mol⁻¹ (I = 0, 25 °C),²¹ X⁰(H⁺) = 0 (by convention) the activation parameter data yield G(TS^≠₂) − G(TS^≠₁) = −12.8 ± 0.2 kJ mol⁻¹, H(TS^≠₂) − H(TS^≠₁) = 26.4 ± 9.3 kJ mol⁻¹ and S(TS^≠₂) − S(TS^≠₁) = +130 ± 32 J K⁻¹ mol⁻¹ (25 °C). This reflects how the enthalpy and entropy factors control the stabilities of the transition states. In other words TS^≠₂ in comparison to TS^≠₁ is assembled with greater energy demand which is more than compensated by the entropy gain.

The intramolecular electron transfer for Mn^III(salen)(OH₂)(OX)⁻ (10⁶k₄ = 0.03 ± 14 at 35 °C) is virtually insignificant. The same is also true for the reaction: Mn^III(salen)(OH₂)₂⁺ + H₂OX (k₃ = 0.13 ± 0.47 dm³ mol⁻¹ s⁻¹, 35 °C, see Table 2). We attempted to examine the reactions via k₀ and k₃ paths using SDS as a probe. Our aim was to selectively partition the neutral H₂OX moiety along with Mn^III(salen)(OH₂)₂⁺ and Mn^III(salen)(OH₂)(HOX) into the anionic micellar pseudo phase and make these reaction paths feasible at a site (i.e. micelle layer) away from the bulk aqueous phase. Our results, however, did not reveal any partitioning of Mn^III(salen)(OH₂)(HOX) and H₂OX into the anionic micelles at the experimental pH (= 3.9). Higher acidity could not be maintained to avoid the competitive pseudo phase ion-exchange equilibrium for (Mn^III(salen)(OH₂)₂⁺)_M/H_W⁺. However, it enabled us to assess the micellar binding of the diaqua complex via ion-exchange equilibrium. The calculated value of the equilibrium constant, K_ex (= 10.4 ± 0.8) is marginally higher than the same for the Na_M⁺/H_W⁺ exchange reported earlier (K_ex = 1, 2–4) (ref. 39 and 40) and comparable to that of trans-(OH₂)Co^III(SO₃)⁺ (K_ex = 10 for β = 0.6 at 25 °C).⁴¹

We undertook the study of the effect of azide ion (N₃⁻) on the redox reaction as a probe for the reaction mechanism. Under the experimental conditions N₃⁻ is not a reductant like HOX⁻. The value of the binding constant (Q₄) for N₃⁻ compares well with that of HOX⁻ (Q₁). The difference in the values of k₁ and k₈ (k₁/k₈ = 1.6 ± 0.3 at 40 °C) is not appreciable. This further suggests that binding of N₃⁻ or HOX⁻ to Mn^III does not lead to any significant difference in the rate of reduction by HOX⁻ that follows. Hence the study of the effect of N₃⁻ validates the fact that trans-(H₂O)Mn^III(salen)(HOX) is kinetically stable to intramolecular electron transfer, .

4. Conclusion

The electron transfer between trans-Mn^III(salen)(OH₂)₂⁺ and oxalate species (H₂OX, HOX⁻, OX²⁻) is preceded by fast and reversible aqua ligand substitution by HOX⁻ resulting in trans-Mn^III(salen)(OH₂)(HOX). The activation parameters for the formation and dissociation of the HOX complex are in support of associative interchange mechanism (I_a). The trans-Mn^III(salen)(OH₂)(HOX) and its conjugate base analogue, trans-MnIII(salen)(OH₂)(OX)⁻ are kinetically inert to intramolecular reduction of Mn^III centre by the bound HOX⁻/OX²⁻; the former undergoes facile redox reaction with HOX⁻ and H₂OX while the latter by only HOX⁻ under mild acidic condition (2 ≤ pH ≤ 6). Structure optimization by DFT shows that there is a strong structural trans-effect of OX²⁻ in trans-Mn^III(salen)(OH₂)(OX)⁻ (unlike for its conjugate acid form) resulting in complete expulsion of H₂O from the Mn^III centre and resulting in a five coordinate species, Mn^III(salen)(OX)⁻ with the H₂O molecule hydrogen bonded to the phenoxide moiety. The trans-(H₂O)Mn^III(salen)(HOX), however, assumes a distorted octahedral structure with a long Mn^III–OH₂ bond. It is possible for the trans-Mn^III(salen)(OH₂)(HOX) to add HOX/H₂OX forming the trans-(HOX)Mn^III(salen)(HOX/H₂OX)^−/0 although such species are not accessible under the experimental conditions. However, the computational study supports the possibility of inner sphere electron transfer processes in the k₁ and k₂ paths. The computed structure of trans-(HOX)Mn^III(salen)(H₂OX) with intramolecular hydrogen bond involving the bound phenoxide and H₂OX, suggests that the intimate redox step in the k₂ path is essentially a proton controlled electron transfer process. The study of the effects of anionic micelles of SDS (sodium dodecyl sulphate) and azide ion provide further support in favor the proposed reaction paths.

Acknowledgements

Financial support from the University grants Commission (UGC), New Delhi in terms of a Teacher Fellowship to AKK (Ref. T. F.OU3-007-1/10-11 (ERO)) is acknowledged. AKK thanks the Odisha Education Department and the authority of Sishu Ananta Mahavidyalaya, Khurda, Odisha, India for granting study leave. Authors are thankful to Prof. Gautam K. Lahiri, Indian Institute of Technology, Mumbai for ESR measurement. HSB acknowledges DST-INSPIRE Faculty fellowship award (IFA11-CH-01) for the financial support.

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Footnote

† Electronic supplementary information (ESI) available: Fig. S1(a and b), S2, S3(a–d), S4(a and b), S5; Tables S1(a–c), S(2–8). See DOI: 10.1039/c4ra10324f

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