H₂O₂ Oxidation by Fe^III–OOH Intermediates and Its Effect on Catalytic Efficiency

Abstract

The oxidation of the C–H and C=C bonds of hydrocarbons with H₂O₂ catalyzed by non-heme iron complexes with pentadentate ligands is widely accepted as involving a reactive Fe^IV=O species such as [(N4Py)Fe^IV=O]²⁺ formed by homolytic cleavage of the O–O bond of an Fe^III–OOH intermediate (where N4Py is 1,1-bis(pyridin-2-yl)-N,N-bis(pyridin-2-ylmethyl)methanamine). We show here that at low H₂O₂ concentrations the Fe^IV=O species formed is detectable in methanol. Furthermore, we show that the decomposition of H₂O₂ to water and O₂ is an important competing pathway that limits efficiency in the terminal oxidant and indeed dominates reactivity except where only sub-/near-stoichiometric amounts of H₂O₂ are present. Although independently prepared [(N4Py)Fe^IV=O]²⁺ oxidizes stoichiometric H₂O₂ rapidly, the rate of formation of Fe^IV=O from the Fe^III–OOH intermediate is too low to account for the rate of H₂O₂ decomposition observed under catalytic conditions. Indeed, with excess H₂O₂, disproportionation to O₂ and H₂O is due to reaction with the Fe^III–OOH intermediate and thereby prevents formation of the Fe^IV=O species. These data rationalize that the activity of these catalysts with respect to hydrocarbon/alkene oxidation is maximized by maintaining sub-/near-stoichiometric steady-state concentrations of H₂O₂, which ensure that the rate of the H₂O₂ oxidation by the Fe^III–OOH intermediate is less than the rate of the O–O bond homolysis and the subsequent reaction of the Fe^IV=O species with a substrate.

Biomimetic analogues play a central role in understanding bioinorganic systems and enzymes, particularly in the identification of reactive intermediates and their role in catalytic processes.¹⁻⁴ In this context, high-valent iron oxo species (i.e., Fe^IV=O) have been studied intensely over the past decade,⁵⁻⁹ especially since their first isolation and crystallographic characterization by Que and co-workers¹⁰ in 2003. The synthetic non-heme Fe^IV=O complexes reported to date show a broad range of reactivity, including C–H oxidation,^7,11−14 with potencies comparable to those of non-heme and heme enzymes, such as Tau-D, and cytochrome P450.¹⁵

High-valent Fe^IV=O species are frequently invoked as the active species engaged in the oxidation of organic substrates by both heme and non-heme enzymes^5,7,16−18 and in biomimetic non-heme iron catalysts. The Fe^IV=O species that have been isolated to date are invaluable in determining their intrinsic reactivity, and their continuous regeneration under catalytic conditions, with H₂O₂ as terminal oxidant, is desirable in achieving turnover in the oxidative transformations that they engage in.

The formation of Fe^IV=O species upon homolytic O–O bond cleavage in the corresponding Fe^III–OOH complexes has been postulated to be a key step for the oxidation of organic substrates by nonheme iron catalysts with H₂O₂:^4,16,19 for example, in the oxidative cleavage of DNA by bleomycin–Fe^III–OOH.^20,21 Notably, in contrast to heme systems, where formation of Fe^V=O species is observed via heterolytic O–O bond cleavage (followed by oxidation of the porphyrin ligand to form compound I). Heterolysis of the O–O bond in low-spin non-heme iron(III)–hydroperoxy species is energetically unfavorable.²²⁻²⁴

However, to the best of our knowledge, this process (Fe^III–OOH → Fe^IV=O) was observed only recently for high-spin Fe^III–OOH species but has not yet been seen for low-spin Fe^III–OOH complexes.²⁵⁻²⁹ Furthermore, the relatively low efficiency of non-heme iron complexes in alkane oxidations with an excess of H₂O₂, together with the known reactivity of Fe^IV=O species with H₂O₂,³⁰ casts doubt on the validity of this paradigm under catalytic conditions.³

The absence of evidence of the formation of Fe^IV=O species and loss of H₂O₂ through unproductive pathways (i.e., disproportionation) can be rationalized by assuming that the generated Fe^IV=O species either reacts with H₂O₂ or engages in, for example, C–H oxidation and hence the reaction of Fe^IV=O with H₂O₂ competes with its reaction with organic substrates. Indeed, Collins and co-workers have shown that the Fe^III(TAML) (TAML = tetraamidato macrocyclic ligand) system disproportionates H₂O₂ through a Fe^IV=O intermediate,² and of direct relevance to the present study, the complex [(N4Py)Fe^IV=O]²⁺ (where N4Py is 1,1-bis(pyridin-2-yl)-N,N-bis(pyridin-2-ylmethyl)methanamine) was shown by Rohde and co-workers to react rapidly with H₂O₂ in acetonitrile.³⁰

In the case of complexes based on pentadentate ligands, e.g., N4Py (Figure 1), the apparent stability of the Fe^III–OOH intermediate and absence of direct spectroscopic evidence for the formation of Fe^IV=O from it make it challenging to identify the actual mechanisms involved in substrate oxidation and H₂O₂ disproportionation.

Figure 1.

Structures of the complexes and intermediates discussed in the present study.

Here, using a combination of time-resolved UV–vis absorption, (resonance) Raman, and EPR spectroscopy and computational chemistry, we demonstrate that, contrary to expectations, the rate of O–O bond homolysis in [(N4Py)Fe^III–OOH]²⁺ to form [(N4Py)Fe^IV=O]²⁺ and a hydroxyl radical is much lower than the rate of H₂O₂ disproportionation observed under reaction conditions. We show that the Fe^III–OOH species is responsible for H₂O₂ decomposition. As a result, the efficiency of substrate oxidation is negatively affected by an increase in the steady-state H₂O₂ concentration, since formation of Fe^IV=O species is uncompetitive.

Experimental Details

Synthesis

The ligand 1,1-bis(pyridin-2-yl)-N,N-bis(pyridin-2-ylmethyl)methanamine (N4Py),³¹ [(N4Py)Fe^II(CH₃CN)](ClO₄)₂ (1),^21,31,32 and [(N4Py)Fe^IV=O](PF₆)₂ (4)³³ were prepared as reported previously. Commercially available chemicals were purchased from Sigma-Aldrich without further purification. All solvents used for spectroscopy were of UVASOL (Merck) grade. H₂O₂ was 50 wt % in H₂O from Sigma-Aldrich and was diluted in methanol as required. The concentration of H₂O₂ in methanol was confirmed by Raman spectroscopy (see Figure S7 for details).

Physical Methods

UV–vis absorption spectra were recorded with a Specord600 (Analytik Jena) spectrophotometer in 1 cm (unless stated otherwise) path length quartz cuvettes. Raman spectra at λ_exc 785 nm were recorded on a PerkinElmer Raman Station at room temperature. Raman spectra at 355 nm (10 mW at source, Cobolt Lasers) were acquired in a 180° backscattering arrangement. Raman scattering was collected by a 2.5 cm diameter plano convex lens (f = 7.5 cm). The collimated Raman scattering passed through an appropriate long pass edge filter (Semrock) and was focused by a second 2.5 cm diameter plano convex lens (f = 15 cm) into a Shamrock500i spectrograph (Andor Technology) 2399 L/mm grating blazed at 300 nm, acquired with an iDus-420-BU2 CCD camera (Andor Technology). The spectral slit width was set to 12 μm. Data were recorded and processed using Solis (Andor Technology) with spectral calibration performed using the Raman spectrum of acetonitrile/toluene, 50/50 (v/v).³⁴ EPR spectra (X-band, 9.46 GHz) were recorded on a Bruker ECS106 spectrometer in liquid nitrogen (77 K) or a Bruker EMX Nano spectrometer. Samples for measurements were transferred to a quartz 3 mm EPR tube (0.5 mL) and flash frozen in liquid nitrogen immediately, concurrent with monitoring by UV–vis absorption spectroscopy.

Computational Details

Computational studies were performed using ADF and QUILD,³⁵⁻³⁷ as reported earlier.³⁸ Briefly, geometry optimization and frequency calculations were performed using the unrestricted density functional BP86-D₃³⁹⁻⁴¹ with a triple-ζ valence plus polarization basis set on iron combined with a double-ζ valence plus polarization on all other atoms (TDZP). Single-point energy calculations on these geometries were made with the S12g spin-state consistent functional^42,43 in a triple-ζ valence plus double polarization (TZ2P) basis set. Free energy (ΔG) corrections were obtained from the BP86-D₃ data and are corrected for zero point energy (ZPE); thermal and entropic corrections were made from frequency calculations at 298 K. The solvation energy was considered using methanol as a solvent with the COSMO solvation model as implemented in ADF.⁴⁴

Caution! The drying or concentration of solutions that potentially contain H₂O₂ should be avoided. Prior to drying or concentrating, the presence of H₂O₂ should be tested for using peroxide test strips followed by neutralization on solid NaHSO₃ or another suitable reducing agent. In work with H₂O₂, suitable protective safeguards should be in place at all times due to the risk of explosion. In experiments where complex 2 is mentioned, it was prepared by dissolution of 1 in methanol (Figure S1).

Caution! In work with perchlorate salts, suitable protective safeguards should be in place at all times due to the risk of explosion. Perchlorate salts should be handled in small (milligram) quantities and used only where necessary.

Results and Discussion

Typically, acetonitrile is the solvent of choice for the reaction of non-heme iron complexes with oxidants such as H₂O₂.⁴⁵ However, in acetonitrile, the formation of [(N4Py)Fe^III–OOH]²⁺ (3) is observed only with a large excess (>50 equiv) of H₂O₂ and the subsequent formation of [(N4Py)Fe^IV=O]²⁺ (4) has not been observed,²¹ despite the fact that 4, prepared independently, is itself stable in acetonitrile even at room temperature. This is in part due to the stability (E_1/2 = 1.2 V vs SCE) of [(N4Py)Fe^II–NCCH₃]²⁺ (1) toward electron transfer oxidation and in part due to the high binding constant of the CH₃CN ligand in comparison with water or H₂O₂. In the present study methanol was chosen to circumvent the formation of such kinetically inert CH₃CN complexes. In methanol, the CH₃CN ligand of 1 exchanges immediately, to form 2 (which is either [(N4Py)Fe^II–OCH₃]⁺or [(N4Py)Fe^II–HOCH₃]²⁺ see the Supporting Information for a discussion), as manifested in a decrease and red shift in the near-UV and visible absorption bands (Figure S1).³² The exchange of the methanol/methoxido ligand for water and H₂O₂ is relatively rapid in both the ferrous and ferric states (vide infra), which is central to enabling observation of other species involved in the reactions discussed and is in stark contrast to the slow ligand exchange seen for 1 in acetonitrile.

Reaction of 2 with Near-Stoichiometric H₂O₂ and Homolysis of O–O Bond of [(N4Py)Fe^III-OOH]²⁺

Addition of 0.6 equiv of H₂O₂ to 2 results in immediate (<2 s) conversion to [(N4Py)Fe^III–OCH₃)]²⁺ (5a) with its characteristic X-band EPR spectrum at g = 2.29, 2.12, and 1.96.⁴⁶ With 1.2 equiv of H₂O₂, [(N4Py)Fe^III–OOH]²⁺ (3) is obtained in minor amounts by both UV–vis absorption and EPR spectroscopy (g = 2.16, 2.11, and 1.98;²¹Figure S2). Addition of 2 equiv of H₂O₂ to 2 results in the formation of [(N4Py)Fe^III–OOH]²⁺ (3) (Figure 2) by ligand exchange over 50 s at room temperature, reaching a maximum of 14% (based on the absorbance at 550 nm, Figure 2-I) before decreasing again over 1000 s. The decrease in absorbance at 550 nm (of 3) proceeds concomitant with an increase in absorbance at 692 nm due to Fe^IV=O (4, Figure 2-II/III). Since 4 reacts rapidly (200 s) with even stoichiometric H₂O₂ (vide infra), its appearance indicates that the concentration of H₂O₂ in solution is already negligible by 80 s (vide infra). The absorbance at 692 nm remains almost constant over 200 s during the decay of 3. The hydroxyl radical formed due to O–O bond homolysis will react with methanol (9.7 × 10^–8 s^–1) to form a methoxy radical that can react with H₂O₂ or other species to yield either methanol or formaldehyde.

Figure 2.

(top) UV–vis absorption spectrum of 2 (0.25 mM) in methanol before (black) and after addition of 2 equiv of H₂O₂ at 21 °C. (bottom) Corresponding change in absorbance over time at 550 and 692 nm. Path length: 2 cm.

Once 3 has been fully consumed, the absorbance at 692 nm then decreases concomitant with the formation of more [(N4Py)Fe^III–OCH₃]²⁺ (5a). These data are consistent with a prior equilibrium between 5a and 3 followed by O–O bond homolysis to form 4. Once sufficient H₂O₂ is consumed, the concentration of 4 is dependent only on the rate of its formation from 3 and the rate of its loss by reaction with methanol (vide infra). The rate of formation of 4 through homolysis of the O–O bond of [(N4Py)Fe^III–OOH]²⁺ (3) under these conditions is low (<2.2 × 10^–4 s^–1, vide infra), which is consistent with the reaction’s endergonicity; calculated (see the Supporting Information) at 19.1 kcal mol^–1. The value is also consistent with the reported value calculated for the related homolytic cleavage in activated Fe–bleomycin.^18,23

1

Disproportionation of H₂O₂ by 3 in Methanol

Addition of excess H₂O₂ (>40 equiv) to 2 in methanol results in immediate oxidation to 5a (i.e., a complete loss in absorbance at 450 nm within the mixing time, 2 s; Figure S3). The oxidation is followed by full conversion of 5a to [(N4Py)Fe^III–OOH)]²⁺ (3) over 5–10 s. The H₂O₂ concentration was monitored in real time by Raman spectroscopy. The second-order rate constant for the formation of 3 from 5a, determined under pseudo-first-order conditions (2.5–50 mM H₂O₂, Figure 3 and Figure S4), is 10.5(±0.1) M^–1 s^–1 at 21 °C, consistent with the exothermicity (−10.2 kcal mol^–1) and low barrier for the exchange of the sixth ligand.

2

Figure 3.

(top) Concentration of [(N4Py)Fe^III–OOH]²⁺ (3, from absorbance at 550 nm) against log(time) for various amounts of H₂O₂ added (5 (red), 10 (black), 20 (blue), 40 (pink), and 400 (khaki) equiv) to 1 (0.56 mM) at 21 °C in methanol. (bottom) Pseudo-first-order rate constant k_obs for the formation of [(N4Py)Fe^III–OOH)]²⁺ vs concentration of H₂O₂.

EPR spectra of samples flash frozen to 77 K (Figure S5) immediately after addition of an excess of H₂O₂ show two well-resolved S = 1/2 signals, characteristic of 3 (major species) and 5a (minor species). Samples, flash frozen after 18 min, show that the signals of 3 are diminished with the concomitant increase of in the signals of 5a, and at ca. 50 min, the signals of 3 are absent, leaving only a more intense signal from 5a.

Notably, both the maximum extent of formation of 3 and the time between addition of H₂O₂ and the start of the subsequent decrease in the absorbance of 3 are dependent on the initial concentration of H₂O₂ (Figure 3). These data indicate that H₂O₂ consumption is relatively similar to the rate of formation of 3 from 5a. The rate of decrease of the absorbance due to 3 is independent of the initial H₂O₂ concentration (Figure S6), because the decay occurs only after essentially all of the H₂O₂ has been consumed, as confirmed by Raman spectroscopy (λ_exc 785 nm, Figure 4). Time-resolved Raman spectroscopy shows that the concentration of H₂O₂ decreases from t = 0 while the resonantly enhanced bands of 3 (Fe^III–OOH) at 632, 650, 670, and 798 cm^–1 do not decrease in intensity until the signal (ν(O–O)) from H₂O₂ at 872 cm^–1 has decreased to near-stoichiometric levels at least (i.e., below the limit of detection of ca. 10 mM, Figure S7).

Figure 4.

(top) Raman spectra of 1 (ca. 5 mM) in methanol over time after addition of 50 equiv of H₂O₂ at λ_exc 785 nm. (bottom) Change in intensity of Raman bands at 872 (of H₂O₂) and 632 cm^–1 (of 3) over time at 21 °C. Spectra correspond to the data points shown.

Regeneration of Fe^III–OOH and O₂ Evolution

For the absorption at 550 nm and its EPR signals, the characteristic Raman bands of 3 appear within the time resolution of the measurement (<60 s) upon addition of excess H₂O₂ and maintain their intensity until all H₂O₂ has been consumed. These data are consistent with the continuous regeneration of 3 from [(N4Py)Fe^III–OR)]²⁺ (where R = H, CH₃) and H₂O₂. 3 is the resting state in the cycle, and the formation of 3 from 5a is a rapid equilibrium prior to the rate-determining step in the reaction.

Headspace analysis by Raman spectroscopy (Figure 5 and Figure S8) confirms generation of O₂ at a rate corresponding to the rate of decrease of H₂O₂. Details for the quantification of O₂ generated are provided in section 3 of the Supporting Information.

Figure 5.

(left) Raman spectra (λ_exc 532 nm) obtained from the headspace above the reaction mixture containing 1 (0.25 mM) and 200 mM H₂O₂ in methanol at 21 °C. (right) Change in intensity of Raman band at 1555 cm^–1 of O₂ (head space, red, λ_exc 532 nm, internal reference was 2329 cm^–1 band of N₂) and at 872 cm^–1 for H₂O₂ (liquid phase, black, λ_exc 785 nm).

The relation between the rate of consumption of H₂O₂ and concentration of 3 is apparent when H₂O₂ is present in excess (>50 equiv). The concentration of 3 remains constant (>80% of total iron concentration) for a period of time, the duration of which is dependent on the initial concentration of 2 (Figure S8). The concentration of H₂O₂, determined by Raman spectroscopy, during this period shows an exponential decay (Figure 6 and Figure S9). The observed rate constant (k_obs) for the decomposition of H₂O₂ is linearly dependent on the catalyst concentration (i.e., [3], Figure 6), with a second-order rate constant of 0.8 M^–1 s^–1 at 21 °C (Figure S9). The rate constant is less than that for the formation of 3 (10.5(±0.1) M^–1 s^–1) and is thus in agreement with 3 as the resting state in the catalytic cycle under steady-state conditions.

Figure 6.

(a) Concentration of H₂O₂ with time following addition of H₂O₂ (200 mM) to 2 (0.25 mM) at 21 °C. (b) Plot of the pseudo-first-order rate k_obs versus concentration of 2.

Reaction of [(N4Py)Fe^IV=O]²⁺ (4) with Methanol and H₂O₂

The self-decay rate of 4 (prepared independently),³³ due to reaction with solvent, is low in acetonitrile³⁰ but is significant in methanol (Figure 7). In methanol, the NIR absorbance of 4 decays exponentially over 1000 s with the concomitant production of 1 equiv of 5a (Fe^III–OCH₃) and 0.5 equiv of formaldehyde (see the Supporting Information for details). The kinetic isotope effect for this decay in CD₃OD is ca. 10 (Figure 7).³³ OH/OD exchange does not affect this rate, which is consistent with the competence of 4 in the oxidation of methanol with a rate-determining hydrogen atom abstraction (HAT) step at the C–H bond.

Figure 7.

Normalized absorbance of 4 (1 mM) at 692 nm in CH₃OH (blue), CH₃OD (red), and CD₃OD (black) with time in the (a) absence and (b) presence of 1 equiv of D₂O₂. Corresponding UV–vis spectra are shown in Figure S10.

Addition of 1 equiv of H₂O₂ to 4 (Fe^IV=O) in methanol results in conversion to 5a (Fe^III–OCH₃) within 200 s, in agreement with data reported in acetonitrile (second-order rate constant of 8 M^–1 s^–1 at 21 °C),³⁰ and is ca. 10 times faster than the reaction of 4 with CH₃OH. However, in stark contrast with the 2:1 ratio of 4 to H₂O₂ required in acetonitrile³⁰ for full reduction of 4 to the Fe^III state, in methanol only 1 equiv of H₂O₂ is required (Figure S10). In both solvents the need for excess H₂O₂ (>0.5 equiv) indicates that H₂O₂ is consumed by other pathways.

3

The OH/D kinetic isotope effect observed in the reduction of 4 with H₂O₂ is masked to some extent by the competing reaction of 4 with CH₃OD (vide supra) but is nevertheless consistent with an HAT mechanism.

In contrast to CH₃OH, in CD₃OD the decay of 4 upon addition of 1 equiv of D₂O₂ is biphasic. Deuterium atom abstraction (from C–D in CD₃OD) by 4 is much slower than the reaction of 4 with 1 equiv of D₂O₂. Consequently, the initial rate (i.e., within 10 min after addition of D₂O₂, 1.13 × 10^–3 s^–1) of decay in the absorbance of 4 is due to reaction with the peroxide only. After this period the rate of decay decreases (to 9.5 × 10^–5 s^–1), which corresponds to the decay of 4 in CD₃OD alone (Figure 7b). A biphasic decay is observed in CH₃OD also but is much less pronounced due to the relatively rapid rate of reaction of 4 with CH₃OD also.

These data indicate that, in addition to reaction with 4, D₂O₂ is consumed through a second pathway, i.e. by the Fe^III species formed initially, which is only apparent when the background reaction of 4 with solvent is slow (i.e., in the case of CD₃OD). The decreased extent of reduction of 4 with 1 equiv of D₂O₂ is similar to that observed in acetonitrile earlier.³⁰ Although calculation of the KIE for reaction of 4 with H₂O₂/D₂O₂ is estimated as close to 10, indicating that HAT is likely to be rate limiting, the occurrence of several reactions in parallel precludes mechanistic interpretation of this value.

With excess H₂O₂ (40 equiv) in CH₃OH, the characteristic NIR absorbance of 4 disappears over 10 s, while that of 3 (Fe^III–OOH, Figure S11) appears concomitantly. These data indicate that the reduction of 4 to 5a by H₂O₂ is followed by ligand exchange with H₂O₂ to form 3. Thereafter, the spectral changes are essentially the same as those observed upon addition of H₂O₂ to 2 in methanol.

In summary, the rate of reaction of 4 follows the order H₂O₂ (in CH₃OH) > D₂O₂ (in CH₃OD) > D₂O₂ (in CD₃OD). Notably, the presence of even 1 equiv of H₂O₂ precludes the presence of 4 in methanol, rationalizing the fact that 4 can be observed only when the concentration of H₂O₂ is substoichiometric. The rate of reduction of 4 by H₂O₂ in acetonitrile was reported by Braymer et al.³⁰ to be insensitive to deuteration (i.e., D₂O₂). In retrospect this observation can be understood by considering the need for an excess of H₂O₂ in that case and the fact that 4 is not the sole species capable of reacting with H₂O₂.

Mechanistic Considerations

The paradigm for oxidation catalysis with complexes such as 2 and H₂O₂ is rapid oxidation to the ferric state and formation of hydroperoxido complexes (e.g., [(N4Py)Fe^III–OOH)]²⁺, 3). Homolytic cleavage of the O–O bond in 3 yields [(N4Py)Fe^IV=O]²⁺ (4) and a hydroxyl radical, both of which are responsible for oxidation of organic substrates. In the present case, only 3, and not 4, is observed in the presence of excess H₂O₂ (Figures 2 and 3), which is consistent with the homolytic cleavage of the O–O bond in 3 being rate determining.

The efficiency in the oxidation of organic substrates is diminished substantially in the presence of excess H₂O₂ due to disproportionation to H₂O and O₂. In the present study, independently prepared 4 is shown to be reduced to the ferric state in methanol rapidly upon addition of H₂O₂. Hence, the fact that 4 is not observed in the presence of excess H₂O₂ can be ascribed to this reaction pathway (b in Scheme 1). Indeed, [(N4Py)Fe^IV=O]²⁺ (4) reacts with H₂O₂ (k = 8 M^–1 s^–1) much more rapidly than the observed rate of decomposition of H₂O₂ (k = 0.8 M^–1 s^–1). However, pathway b (Scheme 1) will be kinetically possible only if the rate of the O–O bond homolysis of 3 is sufficiently rapid to account for the rate of decomposition of H₂O₂.

Scheme 1. Possible Mechanisms for Reaction of 2 with Excess H₂O₂.

In the present study several observations cast doubt on the validity of pathway b. In methanol, the formation of 4 from 3 is observed once (nearly) all H₂O₂ has been consumed; however, the rate of this reaction is much lower (<3.0 × 10^–3 s^–1, Figure 2) than expected. DFT calculations (vide infra) indicate that the cleavage of the O–O bond is substantially uphill and is accompanied by a low barrier to return to 3 (and hence has an intrinsically substantial thermal barrier). Consequently, the rate of formation of [(N4Py)Fe^IV=O]²⁺ (4) is insufficient to account for the decomposition of H₂O₂ when H₂O₂ is present in excess. This conclusion holds the further consequence that the formation of 4 and hence the oxidation of organic substrates by 4 is not competitive with oxidation of H₂O₂ by 3 (pathway a, Scheme 1). The consequence of this is that the oxidation of organic substrates is only competitive under conditions of low H₂O₂ concentration.

DFT Calculations

The mechanism and comparison of two possible pathways for the reaction of 3 with H₂O₂ were explored through DFT calculations. Geometry optimization and frequency calculations were carried out at the BP86-D₃/TDZP level, with subsequent single point energies at the S12g/TZ2P level, including COSMO-ZORA self-consistently at all stages. All data are given in the Supporting Information.

The doublet ground state calculated for 3 is in accordance with experiment. However, for consistency, the reaction pathways a and b (Scheme 1) were calculated in all three possible spin states: doublet, quartet, and sextet (Figure S12). For both pathways the reaction barriers are much lower in the low-spin state in comparison to those in the other two spin states, and hence the discussion below considers only the low-spin states (Figures 8 and 9).

Figure 8.

Comparison of energy profiles (in kcal mol^–1) of pathway a (catalase, in red) and pathway b (homolysis, in black), as obtained at the S12g/TZ2P//BP86-D₃/TDZP level. (The complete structures indicated in the profile can be also found in Figure 9.)

Figure 9.

Geometries (bond distance in Å) for key species for both path A (catalase) and path B (homolysis).

For the disproportionation pathway a, the reactants, 3 + H₂O₂, initially form a reactant complex (RC) where the peroxide is bound to the iron complex weakly. This step is followed (in TS1) by hydrogen atom abstraction from the peroxide toward the distal OH group of 3, and simultaneously cleavage of the O–O bond of 3 takes place with a barrier of only 3.2 kcal mol^–1. In this TS, the O–O bond in 3 elongates from 1.60 to 2.14 Å, together with a shortening of the H–(OH) distance to 1.45 Å. This is followed by a highly exergonic (−32.6 kcal mol^–1) completion of the hydrogen atom transfer process to form H₂O in the intermediate (INT). Simultaneously, the O–O bond of the peroxide shortens from 1.49 to1.35 Å. Formation to the product from INT involves a second hydrogen atom abstraction (barrierless in terms of Gibbs free energy (−1.0 kcal mol^–1) and electronic energy (+1.1 kcal mol^–1)) in which the remaining hydrogen of the peroxide is transferred to the oxygen coordinated to iron. This second HAT is exergonic by ca. −17.0 kcal mol^–1 and finally leads to the products Fe^III–OH + H₂O + O₂.

In contrast, the homolysis pathway b initially forms a similar weakly bound complex in the RC_homo. However, the activation barrier (10.7 kcal mol^–1) for homolytic cleaving of the O–O bond of 3 in TS_homo alone to form Fe^IV=O is much higher than that in pathway a. The O–O bond in 3 elongates from 1.51 to 2.63 Å with hardly any change in the structure of the peroxide: i.e., the peroxide does not participate actively in the reaction but merely acts as a hydrogen-bond donor. More importantly, the product for this homolytic pathway b, P_homo, is so close in energy to the TS_homo (<2 kcal mol^–1) that it readily undergoes the reverse reaction to the initial reactants.

These data are consistent with the observed low rate at which 4 forms from 3 and the rapid consumption of H₂O₂ by direct reaction of 3 with hydrogen peroxide in the disproportionation pathway a.

In summary, there are two pathways that should be considered for the decay of 3. The first is a unimolecular homolysis to form 4 (Fe^IV=O) and a hydroxyl radical. This process is slow and only occurs when the concentration of H₂O₂ is sufficiently low such that it is outside of the solvation sphere of 3. In this case, both 4 (Fe^IV=O) and HO^• are eventually formed and are responsible for the oxidation of organic substrates (i.e., methanol), and hence it is a productive reaction.

At higher concentrations of H₂O₂, i.e. where H₂O₂ is likely to be within the solvation sphere of 3, the formation of two H bonds supports the breaking of the O–O bond of 3, and either (Scheme 2) (i) stabilizes the formation of 4 (“insertion” of H₂O₂ into the O–O bond of 3) or (ii) undergoes HAT to form water and HOO^• from H₂O₂ and, in a subsequent step, a second HAT from HOO^• by Fe^IV=O to form [(L)Fe^III(OH)]²⁺ (5b) and dioxygen.

Scheme 2. Homolysis of the O–O Bond in 3 To Form 4 vs Reaction of 3 with H₂O₂.

Our computational data show that the barrier (Scheme 2) to pathway i is substantial and endergonic (10.7 kcal mol^–1), even with stabilization through H bonding with H₂O₂. The barrier to pathway ii is much lower (ca. 3.2 kcal mol^–1), and leads to the generation of dioxygen (observed experimentally). Hence, in the presence of H₂O₂, 3 is almost exclusively transformed into 5b, with subsequent solvent exchange to 5a (and subsequently through the exchange of methoxido by another H₂O₂ back to 3). Therefore, in the presence of excess H₂O₂, disproportionation into H₂O and O₂ is the more energetically favored pathway.

Regardless of the pathway, the observed reactivity presents a dichotomy toward the use of complexes such as 2 for oxidation catalysis. 3 (Fe^III–OOH) does not appear to react directly with organic substrates (Figure S13),⁴⁷ and hence formation of 4 (Fe^IV=O) and a hydroxyl radical from 3 through O–O bond homolysis is required. However, both 3 (Fe^III–OOH) and 4 (Fe^IV=O) react with H₂O₂ more readily than with methanol. Therefore, ideally the steady-state concentration of H₂O₂ should be held as low as possible, yet still sufficiently high to generate 3 and subsequently 4 (Fe^IV=O)/HO^•. Hence, the rate of addition of H₂O₂ should affect the relative efficiency of 2 in the oxidation of organic substrates, as shown below for the oxidation of methanol.

Competition between the Oxidation of Methanol and H₂O₂ Disproportionation Catalyzed by 2

The oxidation of methanol to methanal occurs concomitantly with the conversion of 4 (Fe^IV=O) to [(N4Py)Fe^II–OCH₃)]⁺ (vide supra). However, the rate of this reaction is sufficiently low to exclude it as being an important pathway for 2 in the presence of excess H₂O₂ (i.e., both Fe^III–OOH and 4 (Fe^IV=O) react much more rapidly with H₂O₂ than with methanol). A number of competing kinetically competent pathways are thus available in the reaction of 2 with H₂O₂, and variation in the steady-state concentrations of reaction components should indicate the relative importance of each of these pathways.

With an 800-fold excess of H₂O₂, ca. 27% of H₂O₂ is disproportionated to H₂O and O₂ (see the Supporting Information for a detailed O₂ concentration calculation), with only 2% generating formaldehyde (Figure 10, bar on the far left). Addition of fewer equivalents of H₂O₂ (second bar from the left) results in a substantial increase in the efficiency in the use of H₂O₂ to oxidize methanol, which increases further by addition of the same amount of H₂O₂ slowly (Figure 10, the two rightmost bars). Adding fewer equivalents slowly over the same time does not increase efficiency further, nor does a change in the concentration of the catalyst (0.25 vs 0.125 mM), since overall conversion rates are controlled by the rate of addition. These data are consistent with the self-decay rate of 4 (Fe^IV=O) (vide supra), setting the upper limit for the rate of addition of oxidant to achieve maximum efficiency.

Figure 10.

Oxidation of methanol (solvent) with H₂O₂ catalyzed by 1 (0.25 mM). HCHO was quantified colorimetrically (see the Supporting Information for details). The number of equivalents of H₂O₂ is with respect to 2. Slow addition of H₂O₂ indicates a rate of addition of 0.4 equiv min^–1 (the two rightmost bars).

Conclusions

The species accepted, i.e. Fe^IV=O, to be responsible for the oxidation of organic substrates by most non-heme iron catalysts is formed from an Fe^III–OOH precursor through O–O bond homolysis, liberating a hydroxyl radical concomitantly. Roelfes and co-workers³ have noted that, in systems where low-spin Fe^III–OOH species are generated (with excess H₂O₂) and observed, the corresponding Fe^IV=O species is not observed.

Here we show that, in the case of non-heme N5 coordinated iron complexes that form observable Fe^III–OOH species, two key reasons can be invoked to rationalize the absence of a corresponding Fe^IV=O species. The first is that, even if it does form, it reacts rapidly and unproductively with H₂O₂ rather than with an organic substrate. Second, and unexpectedly, Fe^III–OOH (3) reacts more rapidly with H₂O₂ in comparison to the rate that it undergoes O–O bond homolysis to form an Fe^IV=O species in the first place.

Ligand exchange, i.e. Fe^III–OR to Fe^III–OOH (3), precedes the oxidation of both organic substrates (eq 2) and H₂O₂ (eq 3). In the present study we show that O–O bond homolysis is relatively slow and is not competitive with the oxidation of H₂O₂ to O₂ by 3. The reaction bifurcation seen for Fe^III–OOH presents a dichotomy in that H₂O₂ must be present in order to form Fe^IV=O (4), but the steady-state concentration of H₂O₂ should be less than that of Fe^III–OOH (3), in order that O–O bond homolysis to form Fe^IV=O and a HO^• radical can take place and hence oxidation of organic substrates to occur. Hence, efficiency with respect to oxidation of organic substrates is increased, as observed in the present study, by maintaining a low steady-state concentrations of H₂O₂.

In conclusion, we show that the oxidation of organic substrates by reactive iron species competes with the reaction of these same species with H₂O₂ and hence wasteful disproportionation of the terminal oxidant. A substantial increase in oxidant efficiency is achieved by maintaining a pseudo-steady-state concentration of H₂O₂ that is below that of the catalyst itself. Furthermore, far from only being a metastable intermediate on route to an Fe^IV=O species, the Fe^III–OOH complex is kinetically competent in its reaction with H₂O₂. The conclusions reached in the present study have implications with regard to our approach to oxidation catalysis with iron catalysts with pentadentate ligands and in a wider perspective hold implications for the mechanisms invoked for catalase type reactions in both biomimetic and bioinorganic systems.

Supporting Information Available

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acscatal.8b02326.

• Details of physical and computational methods, additional spectroscopic data, and coordinates for calculated species (PDF)

Author Present Address

^# A.D.: Department of Chemistry and Center for Metals in Biocatalysis, 207 Pleasant St. SE, University of Minnesota, Minneapolis, Minnesota 55455, United States.

The authors declare no competing financial interest.