Tuning the Redox Properties of a Nonheme Iron(III)–Peroxo Complex Binding Redox-Inactive Zinc Ions by Water Molecules

Abstract

Redox-inactive metal ions play important roles in tuning chemical properties of metal–oxygen intermediates. Herein we report the effect of water molecules on the redox properties of a nonheme iron(III)–peroxo complex binding redox-inactive metal ions. The coordination of two water molecules to a Zn²⁺ ion in (TMC)Fe^III-(O₂)-Zn(CF₃SO₃)₂ (1-Zn²⁺) decreases the Lewis acidity of the Zn²⁺ ion, resulting in the decrease of the one-electron oxidation and reduction potentials of 1-Zn²⁺. This further changes the reactivities of 1-Zn²⁺ in oxidation and reduction reactions; no reaction occurred upon addition of an oxidant (e.g., cerium(IV) ammonium nitrate (CAN)) to 1-Zn²⁺, whereas 1-Zn²⁺ coordinating two water molecules, (TMC)Fe^III-(O₂)-Zn(CF₃SO₃)₂-(OH₂)₂ [1-Zn²⁺-(OH₂)₂], releases the O₂ unit in the oxidation reaction. In the reduction reactions, 1-Zn²⁺ was converted to its corresponding iron(IV)–oxo species upon addition of a reductant (e.g., a ferrocene derivative), whereas such a reaction occurred at a much slower rate in the case of 1-Zn²⁺-(OH₂)₂. The present results provide the first biomimetic example showing that water molecules at the active sites of metalloenzymes may participate in tuning the redox properties of metal–oxygen intermediates.

There is much current interest in understanding the roles of redox-inactive metal ions in modulating reactivities of metal–oxygen intermediates in enzymatic and biomimetic reactions. For example, a redox-inactive Ca²⁺ ion is required to produce O₂ at the oxygen-evolving complex (OEC) of photosystem II (PSII).^[1–3] In the reactions of high-valent metal–oxo complexes, the stabilities and reactivities of iron(IV)–, manganese(IV or V)–, and cobalt(IV)–oxo complexes are markedly affected by binding redox-inactive metal ions (e.g., Sc³⁺ ion).^[4–6] In O₂-activation reactions, Borovik and co-workers demonstrated that the redox-inactive metal ions accelerate the reduction of O₂ by iron(II) and manganese(II) complexes.^[7]

Very recently, nonheme iron(III)–peroxo complexes binding redox-inactive metal ions, [(TMC)Fe^III(O₂)]⁺-Mⁿ⁺ (1-Mⁿ⁺; Mⁿ⁺ = Sr²⁺, Ca²⁺, Zn²⁺, Lu³⁺, Y³⁺, and Sc³⁺), were successfully synthesized by reacting a nonheme iron(III)–peroxo complex, [(TMC)Fe^III(O₂)]⁺ (1),^[8] with various redox-inactive metal ions.^{[9, 10]} The Lewis acidity of the redox-inactive metal ions was shown to be an important factor that determines the redox potentials and reactivities of 1-Mⁿ⁺ with electron donors and acceptors.^[9b] For example, as the Lewis acidity of the redox-inactive metal ions increased, the one-electron oxidation and reduction potentials of 1-Mⁿ⁺ became more positive. Further, the 1-Mⁿ⁺ complexes binding Ca²⁺ and Sr²⁺ ions showed similar redox potentials and reactivities in the one-electron oxidation and reduction reactions. Furthermore, the 1-Mⁿ⁺ complexes binding Ca²⁺ and Sr²⁺ ions were oxidized by an oxidant [e.g., cerium(IV) ammonium nitrate (CAN)] to release O₂, whereas no release of O₂ occurred for complexes binding stronger Lewis acids (1-Mⁿ⁺; Mⁿ⁺ = Zn²⁺, Lu³⁺, Y³⁺, and Sc³⁺).^[9b]

Water molecules at the active sites of metalloenzymes play a vital role in shaping their structures and functions.^[11] Recently, the crystal structure of the OEC in PSII shows the presence of water molecules near the Ca²⁺ ion in the OEC.^[1f] In addition, it has been shown that the Lewis acidity of redox-inactive metal ions decreases by coordinating water molecules.^[12] Therefore, we were curious about the effect of water molecules on the redox properties of metal–dioxygen intermediates binding redox-inactive metal ions. Herein we report for the first time the synthesis of a nonheme iron(III)–peroxo–zinc(II) complex coordinating two water molecules, (TMC)Fe^III-(O₂)-Zn²⁺-(OH₂)₂ [1-Zn²⁺-(OH₂)₂] (Figure 1), and the effect of the water coordination on the Lewis acidity of the Zn²⁺ ion and the redox properties and reactivities of the iron(III)–peroxo complex binding zinc(II) ion, (TMC)Fe^III-(O₂)-Zn²⁺ (1-Zn²⁺; Scheme 1).

Figure 1.

DFT-optimized structure of 1-Zn²⁺ binding two water molecules, (TMC)Fe^III-(O₂)-Zn(CF₃SO₃)₂-(OH₂)₂ [1-Zn²⁺-(OH₂)₂] (color code: Fe, orange; Zn, green; N, blue; O, red; S, yellow; C, gray; F, pink) with the distances [Å] of Fe−Zn and Zn−O bonds. The values in parenthesis are the distances [Å] in the absence of water. Blue dotted lines show the hydrogen bonding interaction.

Scheme 1.

Effects of water coordination to Zn²⁺ ion in 1-Zn²⁺.

An iron(III)–peroxo complex binding Zn²⁺ ion (1-Zn²⁺) was prepared by adding Zn²⁺ ion (50 mm) to [(TMC)Fe^III(O₂)]⁺ (1).^{[8, 9b]} Upon addition of H₂O (0–2.8m) to a solution of 1-Zn²⁺ in CH₃CN (MeCN) at 273 K, the absorption band at 650 nm for 1-Zn²⁺ was red-shifted to 760 nm with a clean isosbestic point at 765 nm, and the shift of the absorption band was found to depend on the amount of H₂O added (Figure 2a). The spectral change of 1-Zn²⁺ exhibits a sigmoidal curvature (Figure 2b), suggesting that more than one H₂O molecule is coordinated to 1-Zn²⁺. The observation of the clean isosbestic point indicates that there is an equilibrium between 1-Zn²⁺ and 1-Zn²⁺ binding water molecules. In such a case, the absorption change (ΔA) at 650 nm due to the formation of 1-Zn²⁺ species binding H₂O is given by Equation (1),

$$\mathrm{\Delta }A=\mathrm{\Delta }{A}_0/(1+K{[{\mathrm{H}}_2\mathrm{O}]}^n)$$ (1)

where ΔA₀ is the final absorption change, n is the number of H₂O molecules coordinated to Zn²⁺ ion, and K is the formation constant. The number of H₂O molecules bound and the formation constant were determined to be 2 and 1.9(1) m⁻² at 273 K, respectively (see Figure S1 in the Supporting Information for the detailed calculation method). This indicates that two H₂O molecules are coordinated to the Zn²⁺ ion to form [(TMC)Fe^III(O₂)]⁺-Zn²⁺-(OH₂)₂ [1-Zn²⁺-(OH₂)₂] (Scheme 1; Figure 1 for DFT-optimized structure; Experimental Section for DFT calculations and Tables S1–S3 in the Supporting Information). In the case of the 1-Ca²⁺ complex, a similar trend of red-shift of absorption band was observed (see Figure S2 in the Supporting Information). In the presence of Ca²⁺ ions (50 mm), the number of H₂O molecules bound to 1-Ca²⁺ and the formation constant of 1-Ca²⁺-(OH₂)₂ were determined to be 2 and 0.52(3) m⁻² at 273 K, respectively (see Figures S2 and S3 in the Supporting Information). However, in the case of Fe^III–peroxo complexes binding redox-inactive metal ions with a stronger Lewis acidity than Zn²⁺ ion (i.e., 1-Sc³⁺), addition of H₂O to a solution of 1-Sc³⁺ resulted in the disappearance of the absorption band at 530 nm due to 1-Sc³⁺, accompanied by the appearance of an absorption band at 810 nm due to 1 (see Figure S4a in the Supporting Information), indicating that Sc³⁺ ion was released from the Fe^III–peroxo moiety in the presence of a large concentration of H₂O.

Figure 2.

a) Absorption spectral changes of 1-Zn²⁺ (0.50 mm; blue line) upon addition of H₂O (0–2.8m with interval of 0.28m) in the presence of Zn²⁺ ions (50 mm) in MeCN at 273 K. b) Plot of absorbance at 650 nm against concentration of H₂O in MeCN at 273 K. Red line shows the fitting line, which is drawn using Equation (1); where n = 2 and K = 1.9m⁻².

The EPR spectrum of 1-Zn²⁺-(OH₂)₂ exhibits signals at g = 9.4 and 4.3 (see Figure S5 in the Supporting Information), which is indicative of high-spin (S = 5/2) Fe^III species. It is worth noting that the EPR feature of 1-Zn²⁺-(OH₂)₂ is completely different from those of 1^[8] and 1-Zn^2+[9b] but similar to that of 1-Sr²⁺ (see Figure S5 in the Supporting Information). The binding of Zn²⁺ ion to 1 in the absence and presence of H₂O molecules was also confirmed by recording a coldspray ionization time-of-flight mass (CSI-TOF MS) spectrum of 1-Zn²⁺ (see Figure S6 in the Supporting Information). However, H₂O molecules in the 1-Zn²⁺-(OH₂)₂ species were not detected under the CSI-TOF MS conditions.

It has been shown that the Lewis acidity of metal ions can be quantitatively determined from the g_zz values of EPR spectra of O₂^•−-metal ion complexes,^[13] since the g_zz values decrease with the increase of the Lewis acidity of metal ions according to Equation (2),

$$g_{\mathrm{z}\mathrm{z}}=g_{\mathrm{e}}+\lambda /\mathrm{\Delta }E$$ (2)

where g_e (= 2.0023) is the g value of free spin, λ (= 0.014 eV) is the spin-orbit coupling constant of oxygen, and ΔE is the energy splitting value of the π_g orbital due to the binding of metal ions to O₂^•−, which can be used as a quantitative measure of the Lewis acidity of metal ions.^{[13, 14]} The Lewis acidity of the Zn²⁺ ion in the presence of 1.4m of H₂O was determined to be (0.57±0.01) eV from the ΔE value (see Figure S7 in the Supporting Information),^{[9b, 13]} which is between the ΔE value of the Ca²⁺ ion (0.58 eV) and the Sr²⁺ ion (0.53 eV; vide infra). We thus conclude that the Lewis acidity of the Zn²⁺ ion in 1-Zn²⁺ decreases by coordinating water molecules (Scheme 1, note 1) and that the Lewis acidity of Zn²⁺ in 1-Zn²⁺-(OH₂)₂ is similar to those of metal ions in 1-Ca²⁺ and 1-Sr²⁺; the Lewis acidities of the metal ions in the latter species are shown to be similar.^[9b] The decrease of Lewis acidity of metal ion upon addition of water was also confirmed by the fluorescence spectral change of the acridone/Zn²⁺ complex (see Figures S8–S10 in the Supporting Information). The fluorescence maximum of the acridone/Zn²⁺ complex is blue-shifted with increasing concentration of water to that of the acridone/Zn²⁺-(OH₂)₂ complex, which is similar to that of the acridone–Sr²⁺ complex (see Figure S9 in the Supporting Information). In addition, the fluorescence maximum did not return back to that of free acridone, being consistent with the result of UV/Vis titration (see Figure 2).

We then investigated the electrochemical oxidation of 1-Zn²⁺ by varying the amounts of H₂O added to the solution of 1-Zn²⁺ (Figure 3a). As shown in the cyclic voltammogram of 1-Zn²⁺ in the one-electron oxidation process (black line), no oxidation peak was observed up to 1.7 V versus SCE.^[9b] In the presence of 0.28m of H₂O, however, the anodic current peak (E_pa) due to the oxidation of 1-Zn²⁺ was observed at 1.45 V versus SCE (pink line in Figure 3a). No corresponding cathodic current peak (E_pc) was observed at the reverse scan due to the release of O₂ upon the one-electron oxidation.^[9b] In the presence of 0.70m of H₂O, the E_pa value of 1-Zn²⁺ was significantly shifted to the negative direction (blue line in Figure 3a; E_pa = 1.06 V vs. SCE). The E_pa value was further shifted to the negative direction upon addition of 1.4 and 2.8m of H₂O (red and green lines in Figure 3a; E_pa = 0.95 and 0.93V vs. SCE), and these values are nearly the same as those of 1-Ca²⁺ (E_pa = 0.96 V vs. SCE) and 1-Sr²⁺ (E_pa = 0.94 V vs. SCE).^[9b] In the presence of > 2.8m of H₂O, no further shift of the E_pa value was observed.

Figure 3.

Cyclic voltammograms of 1-Zn²⁺ in the absence (black lines) and presence of H₂O [0.28m (pink), 0.70m (blue), 1.4m (red), and 2.8m (green)] in one-electron oxidation (a) and one-electron reduction (b) processes in MeCN at 273 K.

In the case of the one-electron reduction process of 1-Zn²⁺, the cathodic current peak (E_pc) due to the reduction of 1-Zn²⁺ was observed at −0.16 V versus SCE in the absence of H₂O (Figure 3b, black line),^[9b] but no corresponding anodic current peak (E_pa) was observed at the reverse scan due to the O−O bond cleavage of the one-electron reduced species, which resulted in the formation of the corresponding iron(IV)–oxo species, [(TMC)Fe^IV(O)]²⁺ (2).^{[9, 10]} In the presence of 0.70m of H₂O, the E_pc value was shifted to the negative direction (blue line in Figure 3b; E_pc = −0.31 V vs. SCE). The E_pc values in the presence of 1.4 and 2.8m of H₂O (E_pc = −0.38V vs. SCE) were further shifted to the negative direction (red line in Figure 3b and Figure S11 in the Supporting Information for 1.4 and 2.8m of H₂O, respectively). These results clearly indicate that the one-electron oxidation and reduction potentials of 1-Zn²⁺ decrease as the amount of added H₂O increases (Scheme 1, note 2). Moreover, the oxidation and reduction potentials of 1-Zn²⁺-(OH₂)₂ are similar to those of 1-Ca²⁺ and 1-Sr²⁺.^[9b]

The electron-transfer oxidation of 1-Zn²⁺ in the absence and presence of H₂O was examined by using CAN as an oxidant. Addition of one equivalent of CAN to the solution of 1-Zn²⁺ resulted in the immediate disappearance of the absorption band at 650 nm due to 1-Zn²⁺ with the concomitant appearance of the absorption band at 585 nm (see Figure S12 in the Supporting Information). In this reaction, no release of O₂ was detected by GC and mass spectrometry, as reported previously in the reactions of 1-Mⁿ⁺ (Mⁿ⁺ = Zn²⁺, Lu³⁺, Y³⁺, and Sc³⁺) with CAN.^[9b] We may assign the new species with an absorption band at 585 nm as a cerium(IV) ion-bound iron(III)–peroxo species (1-Ce⁴⁺); we have shown previously that the Zn²⁺ ion in 1-Zn²⁺ can be replaced by redox-inactive metal ions with a greater Lewis acidity and the Lewis acidity of the Ce⁴⁺ ion is higher than that of the Zn²⁺ ion.^[9b]

In sharp contrast to the result of 1-Zn²⁺ (i.e., no H₂O added), addition of one equivalent of CAN to the solution of 1-Zn²⁺-(OH₂)₂, which was prepared by adding 1.4m of H₂O to 1-Zn²⁺, caused the immediate disappearance of the absorption band at (715±5) nm due to 1-Zn²⁺-(OH₂)₂ (Figure 4a). The analysis of the reaction solution revealed the formation of [(TMC)Fe^II]²⁺ species by EPR and electrospray ionization mass spectrometry (ESI MS; see Figure S13 in the Supporting Information).^[9b] In this case, the release of O₂ was detected by GC and mass spectrometry ((83±5)% yield based on 1-Zn²⁺ used; see Experimental Section in the Supporting Information), showing that 1-Zn²⁺-(OH₂)₂ was oxidized by CAN to produce O₂ and the Fe^II complex, as reported previously in the reactions of 1-Ca²⁺ and 1-Sr²⁺.^[9b] To confirm the source of O₂, ¹⁸O-labeled 1-Zn²⁺-(OH₂)₂ complex was prepared and allowed to react with CAN. The resulting mass spectra in Figure 4b clearly show the evolution of ¹⁸O₂, demonstrating unambiguously that the ¹⁸O-labeled 1-Zn²⁺-(OH₂)₂ complex released its binding ¹⁸O₂ by the oxidation with CAN (Scheme 1, note 3).

Figure 4.

a) Absorption spectral change observed in the reaction of 1-Zn²⁺-(OH₂)₂ (0.50 mm) and CAN (0.50 mm; 1.0 equiv) in the presence of H₂O (1.4m) in MeCN at 253 K. Absorption band at 715 nm (red line) due to 1-Zn²⁺-(OH₂)₂ (i.e., in the presence of 1.4m of H₂O) was decayed to black line upon addition of CAN in MeCN at 253 K. b) Mass spectra monitoring O₂ isotopes [32 (¹⁶O–¹⁶O; black circles), 34 (¹⁶O–¹⁸O; blue circles), and 36 (¹⁸O–¹⁸O; red circles)] produced in the reactions of ¹⁸O-labeled [(TMC)Fe^III(¹⁸O₂)]⁺-Zn²⁺ (0.25 mm) with CAN (0.25 mm) in the presence of H₂O (1.4m) in MeCN at 253 K.

The electron-transfer reduction of 1-Zn²⁺ in the absence and presence of H₂O was also examined by using 1,1′-dimethylferrocene (Me₂Fc) as a reductant. Electron transfer from Me₂Fc to 1-Zn²⁺ in the absence of H₂O resulted in the formation of [(TMC)Fe^IV(O)]²⁺ (2) via heterolytic O−O bond cleavage of the peroxide ligand (see Figure S14 in the Supporting Information).^{[9, 10, 15]} The rate of electron transfer obeyed first-order kinetics and the pseudo-first-order rate constant increased linearly with the increase of the Me₂Fc concentration. The second-order rate constant (k_et) of electron transfer was determined to be 23(2) m⁻¹ s⁻¹ (see Figure S14a in the Supporting Information). The k_et value decreased with increasing the H₂O concentration (Figure 5 and Figure S15 in the Supporting Information);^[9] the k^et value determined in the presence of 2.8m of H₂O was approximately 50 times smaller than that determined in the absence of H₂O. It should be noted that no reaction between 1 and Me₂Fc occurred without Zn²⁺ ion in the absence and presence of H₂O. These results clearly indicate that the rate of the one-electron reduction of 1-Zn²⁺ depends on the H₂O concentration in solution. These results are also in line with our previous observation that the one-electron reduction of 1-Mⁿ⁺ depends on the Lewis acidity of the Mⁿ⁺ ion;^[9b] the O−O bond cleavage of 1-Zn²⁺ slows down as the Lewis acidity of the Zn²⁺ ion in 1-Zn²⁺ decreases by coordinating water molecules (Scheme 1, note 4).

Figure 5.

Plot of k_et against concentration of H₂O in electron transfer from Me₂Fc to 1-Zn²⁺ (0.50 mm) in the absence and presence of H₂O (0m, 0.56m, 1.1m, 1.7m, 2.2m and 2.8m) in MeCN at 273 K.

In conclusion, we have demonstrated the significant effect of water molecules on the redox properties and reactivities of a nonheme iron(III)–peroxo complex binding redox-inactive Zn²⁺ ion (1-Zn²⁺; Scheme 1). We have interpreted these results with the change of the Lewis acidity of the Zn²⁺ ion owing to the coordination of water molecules to the Zn²⁺ ion in 1-Zn²⁺. The present results suggest that water molecules at the active sites of metalloenzymes play important roles in fine-tuning of the redox properties of metal–oxygen intermediates.