The oxidation potentials of Mn^II in aqueous solutions of bicarbonate, formate, acetate and oxalate are reported as a function of concentration and compared to the rate of photooxidation of these solutions by the Mn-depleted water-oxidizing complex of photosystem II (apo-WOC-PSII) from peas.

This shift is in agreement with the onset of formation of [Mn(OH)]⁺ (pK_a for Mn²⁺_aq is 10.5).

Therefore, in the present experiments at pH lower than 8.35, Mn²⁺_aq, and not [Mn(OH)]⁺, is the predominant Mn^II species detected electrochemically in equilibrium with exogenous anions.

A small cathode wave at 1.12 V and additional two peaks at 0.96 and 0.28 V are observed upon the return scan, reflecting the presence of all the reactions, e.g., (1–5) of Mn^II oxidation in water.

The difference between the anode (E_p^a) and the first cathode (E_p^k) peaks for Mn^II is equal to 60 mV, both in the presence and absence of acetate thus indicating a reversible transfer of one electron per Mn^II during the oxidation step.

The peak current of the cathode peak is much lower than the anode peak current, which indicates that a subsequent transformation of the oxidation product into another product takes place.

In the concentration region from 2.15 × 10⁻³ M to 7.11 × 10⁻³ M (Fig. 3, curve 1) the dependence of the peak potential on log C_X is linear and the slope of this dependence (ΔE/Δlog C_X) is equal to 58 mV, indicating participation of just one ligand in the electrode reaction:Mn^II− e⁻ + Ox²⁻ ⇔ [Mn^III(Ox²⁻)]⁺.Extrapolation of this linear part of the dependence to log C_X = 0 gives the value for E⁰, equal to 0.92 V. Using eqn. (I), the logarithm of the stability constant of the Mn^III-oxalate complex is equal to 10 ± 0.01 which is in agreement with the literature data for log K_st of [Mn^III(C₂O₄²⁻)]⁺ equal to 9..98⁴⁰

Extrapolation of the linear dependence of E_p to log C_X = 0 gives the E⁰ value equal to 0.77 V. The same value is obtained using eqn. (II) and in the available literature data, as follows.

Then the logarithm of K_ox/K_red will be equal to 12.75 that [according to eqn. (II)] gives ΔE⁰ equal to 0.75 V and E⁰_com equal to 0.76 V. This (calculated) value agrees with the value of E⁰ obtained from our experimental work (0.77 V).

The slope is equal to 120 mV which by eqn. (I), gives p = 2 for n = 1 and thus corresponds to a 1∶2 ratio of Mn∶acetate, formulated as [Mn^III(CH₃COO⁻)₂]⁺:Mn^II − e⁻ + 2CH₃COO⁻ → [Mn^III(CH₃COO⁻)₂]⁺.Extrapolation of the potential dependence to log C_X = 0 gives E⁰ = 0.69 V for this process and enables calculation of the stability constant for acetate binding, K_st = 7.91 × 10¹³ at T = 21.5 °C.

The slope is equal to 131 ± 5 mV, which according to eqn. (I) gives p ≅ 2 at n = 1, corresponding to a reaction stoichiometry of 1∶2:Mn^II − e⁻ + 2HCO₂⁻ → [Mn^III(HCO₂⁻)₂]⁺.Extrapolation of the data to log C_X = 0 gives E⁰ = 0.775 V for Mn^III(HCO₂⁻)₂⁺ and a corresponding stability constant K_st equal to 2.86 × 10¹².

(i) The linear dependence of the Mn^II oxidation potential on log C_X at concentration higher than 3 × 10⁻³ M, NaHCO₃ (Fig. 3, curve 4) reflects the oxidation of free aqua Mn^II (in analogy with the other anions) since it extrapolates back to the free Mn^II potential.

The slope of this dependence is equal to 173 ± 5 mV, in contrast to the slope for oxalate, acetate and formate, and thus by eqn. (I) corresponds most closely to p ≅ 3 for n = 1.

This ratio suggests participation of three HCO₃⁻ anions in this electrode reaction:Mn^II − e⁻ + 3HCO₃⁻ ⇔ Mn^III(HCO₃⁻)₃E⁰ is equal to 0.67 V for this reaction (determined by extrapolation to log C = 0).

This value gives a stability constant for Mn^III(HCO₃⁻)₃K_st = 1.73 × 10¹⁴ [T = 21 °C].

It is clearly seen (Fig. 3, curve 4) that the plot of the dependence of Mn^II oxidation potential on log C_X is free of any breaks, thus indicating that a single species is detected, formulated as the electro-neutral species Mn^III(HCO₃⁻)₃.

(ii) The voltage–current curves in Fig. 4 show that a second precursor is revealed (as a “new” wave of oxidation) at increasing bicarbonate concentrations that is more easily oxidized (appears at less positive potential) than the previous species.

The experimental dependence of the oxidation potential of this new wave on log C_X (from −1.4 to −1.08) reveals a linear region with slope equal to 60 ± 2 mV (Fig. 3, curve 5), suggesting the involvement of one additional bicarbonate ion per electron in the oxidation of this new complex.

Extrapolation of the data in Fig. 3 (curve 5) to log C_X = 0 gives the E⁰ = 0.61 V (complex I).

(iii) Upon further increase of the NaHCO₃ concentration (to 0.1 M and higher) there is a break in the dependence of E_P on log C_X (Fig. 3, curve 5) with a linear slope equal to 120 ± 10 mV, suggesting the involvement of two additional bicarbonate ions per electron in the oxidation of this new complex.

In this case, the slope extrapolates to E⁰ = 0.52 V at log C_X = 0.

The appearance of this new section is evidently related to oxidation of another Mn^II-BC complex (complex II).

In this region (curves 7–12) Mn_ag^II is oxidized by the reaction:Mn^II − e⁻ + 3HCO₃⁻ ⇒ Mn^III(HCO₃⁻)₃.The observed equilibria are the same as found in Fig. 4.

This stimulation occurs using Mn^II concentrations equal to the native stoichiometry of 4 Mn/PSII and over a 1000-fold range in ligand concentrations.

The oxidation potential for Mn^III/Mn^II decreases from 1.18 V for Mn_aq^II (in 0.1 M LiClO₄, pH independent in the region pH 5.0–pH 8.5) to less positive values upon formation of Mn^III complexes with acetate, formate, oxalate and bicarbonate, respectively:Mn^II − e⁻ + 2CH₃COO⁻ ⇔ [Mn^III(CH₃COO⁻)₂]⁺E⁰ = 0.69 V, (K_st = 7.9 × 10¹³)Mn^II − e⁻ + 2HCO₂⁻ ⇔ [Mn^III(HCO₂⁻)₂]⁺E⁰ = 0.775 V, (K_st = 2.9 × 10¹²)Mn^II − e⁻ + Ox²⁻ ⇔ [Mn^III(Ox²⁻)]⁺E⁰ = 0.92 V, (K_st = 1.0 × 10¹⁰)Mn^II − e⁻ + 3HCO₃⁻ ⇔ Mn^III(HCO₃⁻)₃E⁰ = 0.67 V, (K_st = 1.7 × 10¹⁴)In the bicarbonate stimulated oxidation process the stoichiometry of 3 bicarbonates reflects the total number of bicarbonate molecules involved in the process and thus includes several potential Mn^III species differing in both ligation and deprotonation of bicarbonate and hydroxide.

In case of acetate and formate, a positively charged complex, Mn^IIIL₂, forms upon oxidation of Mn_aq^II.

Only for bicarbonate does the neutral complex Mn^III(HCO₃⁻)₃ form upon oxidation of Mn_aq^II_..

The absence of breaks in the plot of E_Pvs. log of the concentration of acetate and formate shows that, in spite of the possibility of formation of Mn^II-complexes with the added ligands (log K_st for the Mn^II–(CH₃COO⁻)⁺ is equal to 1.2),⁴⁰ there is no additional wave related to the oxidation of the Mn^IIL complexes.

The E⁰ value for Mn_aq^II is equal to 1.18 V, which is higher than the oxidation potential of the Y_Z˙ radical estimated to 0.95–1.0 V.⁴¹

Acetate and formate are not active in stimulation of electron transfer from Mn^II to apo-WOC-PSII, although both anions considerably lower the reduction potential of Mn^III/Mn^II.

Bicarbonate efficiently stimulates the process of Mn^II-dependent electron transfer to apo-WOC-PSII, and is also known to reverse the formate-induced inhibition of electron transport from Mn^II and O₂ evolution of functionally active PSII with high specificity.^14–18