Syddansk Universitet

Electrochemically Generated cis-Carboxylato-Coordinated Iron(IV) Oxo Acid-BaseCongeners as Promiscuous Oxidants of Water Pollutants

Poulsen de Sousa, David; Miller, Christopher J; Chang, Yingyue; Waite, T David; McKenzie,ChristinePublished in:Inorganic Chemistry

DOI:10.1021/acs.inorgchem.7b02208

Publication date:2017

Document versionPublisher's PDF, also known as Version of record

Citation for pulished version (APA):de Sousa, D. P., Miller, C. J., Chang, Y., Waite, T. D., & McKenzie, C. J. (2017). Electrochemically Generatedcis-Carboxylato-Coordinated Iron(IV) Oxo Acid-Base Congeners as Promiscuous Oxidants of Water Pollutants.Inorganic Chemistry, 56(24), 14936-14947. DOI: 10.1021/acs.inorgchem.7b02208

General rightsCopyright and moral rights for the publications made accessible in the public portal are retained by the authors and/or other copyright ownersand it is a condition of accessing publications that users recognise and abide by the legal requirements associated with these rights.

• Users may download and print one copy of any publication from the public portal for the purpose of private study or research. • You may not further distribute the material or use it for any profit-making activity or commercial gain • You may freely distribute the URL identifying the publication in the public portal ?

Take down policyIf you believe that this document breaches copyright please contact us providing details, and we will remove access to the work immediatelyand investigate your claim.

Download date: 08. Jan. 2018

Electrochemically Generated cis-Carboxylato-Coordinated Iron(IV)Oxo Acid−Base Congeners as Promiscuous Oxidants of WaterPollutantsDavid P. de Sousa,† Christopher J. Miller,‡ Yingyue Chang,‡ T. David Waite,‡

and Christine J. McKenzie*,†

†Department of Physics, Chemistry and Pharmacy, University of Southern Denmark, Campusvej 55, 5320 Odense, Denmark‡Water Research Centre, School of Civil and Environmental Engineering, University of New South Wales, Sydney, New South Wales2052, Australia

*S Supporting Information

ABSTRACT: The nonheme iron(IV) oxo complex [FeIV(O)-(tpenaH)]2+ and its conjugate base [FeIV(O)(tpena)]+ [tpena−

= N,N,N′-tris(2-pyridylmethyl)ethylenediamine-N′-acetate]have been prepared electrochemically in water by bulkelectrolysis of solutions prepared from [FeIII2(μ-O)(tpenaH)2]-(ClO4)4 at potentials over 1.3 V (vs NHE) using inexpensiveand commercially available carbon-based electrodes. Oncegenerated, these iron(IV) oxo complexes persist at roomtemperature for minutes to half an hour over a wide range ofpH values. They are capable of rapidly decomposing aliphaticand aromatic alcohols, alkanes, formic acid, phenols, and thexanthene dye rhodamine B. The oxidation of formic acid tocarbon dioxide demonstrates the capacity for total mineraliza-tion of organic compounds. A radical hydrogen-atom-abstraction mechanism is proposed with a reactivity profile for the seriesthat is reminiscent of oxidations by the hydroxyl radical. Facile regeneration of [FeIV(O)(tpenaH)]2+/ [FeIV(O)(tpena)]+ andcatalytic turnover in the oxidation of cyclohexanol under continuous electrolysis demonstrates the potential of the application of[FeIII(tpena)]2+ as an electrocatalyst. The promiscuity of the electrochemically generated iron(IV) oxo complexes, in terms of thebroad range of substrates examined, represents an important step toward the goal of cost-effective electrocatalytic waterpurification.

■ INTRODUCTION

According to the World Health Organization, never before haveso many people been without clean drinking water.1 Not only isthe threat to freshwater supplies a concern for developingcountries, but it is also becoming an increasing issue in thewestern world where pesticides, persistent pharmaceuticals,plasticizers, textile dyes, fluorinated coatings, fire retardants,organic solvents, and natural and artificial hormones arecontaminating freshwater reservoirs. The pressure for thereuse of treated wastewaters, potentially containing residualamounts of these contaminants, is intensifying.2,3 Chemicallycontaminated natural waters have been linked to hormonaldisruptions in terrestrial and marine life4 and a decrease inhuman male fertility5 and are also a risk factor in cancerous andcongenital disorders.6 The destruction of persistent organicpollutants from waters for environmental remediation anddrinking water production is challenging because lowbiodegradability means that these contaminants are likely tobe unaffected by conventional biological treatment processes7

and potentially difficult to remove by standard physicochemicalprocesses. Some progress has been achieved through the use of

advanced oxidation processes for degrading pollutants in whichreactive hydroxyl radicals (HO•) are generated from terminalchemical oxidants such as hydrogen peroxide (H2O2), ozone, orpersulfate.8 One of the most common pathways by which HO•

is produced in both natural and engineered systems is viaFenton reactions, whereby H2O2 reacts with [FeII(H2O)6]

2+ toproduce HO• and aquated and/or hydroxylated ferric speciessuch as [FeIII(H2O)6]

3+ and [FeIII(OH)(H2O)5]2+. The

reaction is catalytic under acidic conditions because theiron(III) species can be reduced back to the iron(II) state byreaction with a second equiv of H2O2.

9,10 When the pH israised, however, HO• is no longer formed,11 and ironprecipitates as a ferric oxide sludge, which itself must bedisposed of.12 Complexation with polyaminecarboxylato ligandssuch as ethylenediaminetetraacetic acid (EDTA) and relatedligands13 can prevent iron oxide precipitation and allow HO•

production at circumneutral pH;14 however, the resultant FeIII-EDTA complexes are not readily reduced back to the active

Received: August 27, 2017Published: October 17, 2017

Article

pubs.acs.org/IC

© 2017 American Chemical Society 14936 DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

This is an open access article published under an ACS AuthorChoice License, which permitscopying and redistribution of the article or any adaptations for non-commercial purposes.

iron(II) form, in effect rendering them stoichiometric oxidantsonly. The discovery of new ligand systems for promotingtransient generation of catalytically competent high-valent iron-based oxidants in sustainable redox cycles is therefore of primeimportance. Electrochemical activation, rather than terminalchemical oxidant activation, of iron catalysts to form a high-valent species is an alternative but as yet unexploited approachfor the production of iron-based oxidants. In water, potentiron(IV) oxo species could be produced by such a potentiallymore environmentally benign method (using renewableelectricity sources). It would present a significant advancementover chemical activation using excess amounts of atom-uneconomical and logistically difficult terminal oxidants, andcleaner methodologies for the oxidation of organic substrateswould be opened up.The ability to activate chemical oxidants in water with the

production of high-valent iron complexes is a good startingpoint in the search for the appropriate electrocatalysts forachieving this goal. Several ligand classes have beendemonstrated to support high-valent iron. These includeporphyrin, thiocarbamato, amido, and aminopolypyridylligands. Of particular note is the work of Collins and co-workers on the iron complexes of tetraanionic tetraamidomacrocyclic ligands (abbreviated generally as TAMLs) toactivate H2O2. These reactions are proposed to produce mono-and dinuclear high-valent oxo complexes [FeIV/V(O)-(TAML)]2−/− and [FeIV2(μ-O)(TAML)2]

2− and not HO•.The structure of a mononuclear iron(IV) oxo complex with theprototype TAML ligand 3,4,8,9-tetrahydro-3,3,6,6,9,9-hexam-ethyl-1H-1,4,7,10-benzotetraazacyclotridocane-2,5,7,10-(6H,11H)tetrone (H4B*) is shown in Scheme 1. These H2O2-

activating systems have been applied to water remediation. Theprototype shown functions optimally around pH 9; however,since this initial work, the Fe-TAML system has evolvedthrough scaffold modification to be capable of the oxidativedegradation of a range of compounds in aqueous solutionincluding azo- and phthalocyanine dyes,15 the antidepressantsertraline,16 organophosphorus pesticides,17 nitrophenols,18 andchlorinated phenols and metaldehyde,19,20 at lower pH valuesapproaching 7. Electroactivation and consequent wateroxidation was reported for a carbon-immobilized Fe-TAML.21

This result raises the issue of selectivity of the high-valent ironoxidants toward the oxidation of C−H and CO−H bonds of

organic micropollutants over the O−H bonds of water. If wateris the medium, it is likely to outcompete micropollutantmineralization, with the result that too much energy will beconsumed. Fine-tuning the design of the ligand scaffoldsupporting molecular iron catalyst design is crucial.In parallel to the development of Fe-TAML systems, an

understanding of the nature and reactivity of molecularnonheme iron(IV) oxo complexes based on aminopolypyridylligands has developed. Examples of this class of ligand are thepentadentate systems, parent to the system used in the workdescribed below; i.e., N′-benzyl-N,N,N′-tris(2-pyridylmethyl)-ethylenediamine (bztpen)22 and N,N′-bis(2-pyridylmethyl)-N,N′-bis(2-pyridylmethyl)amine (N4Py),23 as depicted inScheme 1. The neutrality of these ligands means that theirhigh-valent iron oxo complexes should inherently be moreelectrophilic relative to the Fe-TAML systems. Unfortunately,these systems are typically incompatible with aqueouschemistry often because of the formation of chemicallyunreactive stable oxo-bridged iron(III) systems as thermody-namic sinks. They are synthesized in organic solvents from thereaction of iron(II) precursors with hypervalent iodine reagentsor alkyl peroxides, in two-electron reactions. Under non-aqueous conditions they have been shown to stoichiometricallyoxidize alkanes, alkenes, sulfides, ketones and alcohols.24−27

Oxygen-atom-transfer (OAT) and/or hydrogen-atom-transfer(HAT) mechanisms have been proposed. The ability to targetthe oxidation of specific substrates is often limited, andoveroxidation beyond the intended scope occurs. While thismay not be ideal for selective catalysis, such promiscuousreactivity is ideal for applications where oxidative power, andnot selectivity, is desired, such as in wastewater treatment.We have shown that, in contrast to the nitrogen-atom-donor-

only polypyridyl-supported systems, the iron(IV) oxo species[FeIV(O)(tpenaH)]2+ [tpena− = N,N,N′-tris(2-pyridylmethyl)-ethylenediamine-N′-acetate; Scheme 1] can be generated inwater by oxidizing solutions of [FeIII2(μ-O)(tpenaH)2](ClO4)4by the reaction with cerium(IV) ammonium nitrate.28 In fact,[FeIV(O)(tpenaH)]2+ cannot be spectroscopically observed inorganic solvents, presumably because of its spontaneous decayas the result of it oxidizing these solvents. In comparison, it isrelatively stable in water with a half-life of around 2 h. Thechemical generation of an iron(IV) oxo complex by one-electron oxidation of an iron(III) precursor contrasts with thetwo-electron OAT methods used for the preparation of theaforementioned iron(IV) oxo aminopolypyridyl complexes, likethose based on bztpen and N4Py, from their resting iron(II)states. We therefore speculated that [FeIII2(μ-O)(tpenaH)2]

4+

might be an interesting candidate complex for the investigationof its electrooxidation in water to form the iron(IV) oxospecies. Aside from the salient stability of the 3+ state over the2+ oxidation state, an important feature afforded by the Fe-tpena complexes is the unique presence of a carboxylate donoralong with flexible ligand denticity. When fully coordinated, thetpena ligand is hexadentate, as verified by the structurallycharacterized homoleptic [MIII(tpena)]2+ complexes (M =Cr,28 Co,29 and Fe30). Interestingly, the iron complex can beisolated only in the absence of water. Even if trace water ispresent hemi- and monohydrates, [FeIII2(μ-O)(tpenaH)2]

4+

and [FeIII(OH)(tpenaH)]2+, respectively, are formed. Inthese, a pyridine arm is decoordinated and protonated. It isnoteworthy, with respect to the lability of the iron complextoward nucleophiles including water and thus pertinent to thework described below, that these [MIII(tpena)]2+ complexes

Scheme 1. Structure of the Iron(IV) Oxo Complex oftpenaH/tpena− and Related Systems

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14937

show irregular octahedral geometries due to the small bite ofthe ethylenediamine backbone. The distortion from 90° for thechromium(III) complex is exceptional, compared to typicalWerner-type complexes (one cis-N−Cr−N angle is 115°). Thestrain can be relieved by 7-coordination, as seen in the crystalstructures of the high-spin d5 systems ([FeIII(tpena)(OIPh)]2+

and [MnII(tpena)(H2O)]+)31,32 or by decoordination of a

pendant arm, as observed in [VIV(O)(tpenaH)]2+ and thehemihydrate [FeIII2(μ-O)(tpenaH)2]

4+.26 When the latteroccurs, tpena is pentadentate and the decoordinated methyl-pyridyl group protonated. [VIV(O)(tpenaH)]2+ is isovalent withthe iron-based oxidant [FeIV(O)(tpenaH)]2+ and models itsstructure. Protonation on the dangling pyridine of [FeIV(O)-(tpenaH)]2+ (rather than formation of an FeIVOH species) issupported by the reproduction of the Mossbauer spectrum bydensity functional theory calculations.26 The proximity of anuncoordinated pyridinium/pyridine group is reminiscent of thenatural second coordination sphere bases histidine and lysine.Together with the carboxylato donor and a coordination spherecompleted by neutral nitrogen-atom donors, this makes

[FeIV(O)(tpenaH)]2+ one of the most germane nonhemeiron(IV) oxo enzyme mimics reported to date.Electrooxidation is envisaged to be kinetically demanding,

and the Fe-tpena system has the same advantage as the Fe-TAML systems in that a one-electron-only process is requisiteto reach the iron(IV) state from aqueous solution resting states.The fact that [FeIV(O)(tpenaH)]2+ appears to be apromiscuous oxidant that can be chemically generated usingone-electron chemistry from stable aqueous solution iron(III)resting states using an easily accessible solid-state precursor andwater as the oxygen-atom donor has prompted us to investigatethe potential for its generation in water using electrochemicalmethods. We describe here success in this approach, along withthe first systematic study of the reactivity of aqueouselectrochemically generated nonheme iron(IV) oxo complexes.We show that the acid−base congeners [FeIV(O)(tpenaH)]2+

and [FeIV(O)(tpena)]+ are active toward the oxidation andmineralization of a wide range of organic pollutants under bothstoichiometric and electrocatalytic conditions.

Figure 1. (A) Schematics of the pH-dependent iron(III) speciation of [FeIII2(μ-O)(tpenaH)2](ClO4)4 dissolved in water. (B) Selected CVs recordedfor aqueous solutions of [Fe2(O)(tpenaH)2](ClO4)4 (0.8 mM per Fe; 0.1 M NaClO4) at various pH values. Scan rate = 100 mV s−1. The full seriesand assignments are provided in the Supporting Information. (C) Pourbaix (potential−pH) diagram for the FeII/III/IV-tpena system (L = tpena).Open red squares (□) denote anodic peak potentials for the oxidation of [FeIII(OH2)(tpenaH)]

3+ to [FeIV(O)(tpenaH)]2+ (2H+/1e− and Ep = 1.25V). Open blue squares (□) denote the oxidation of [FeII(OH2)(tpenaH)]

2+ to [FeIII(OH)(tpenaH)]2+ (1H+/1e− and Ep = 0.84 V). Open circles(O) denote the cathodic peak potentials for unassigned {Fe2(O)} reductions (Ep ∼ 0.35−0.07 V).

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14938

■ RESULTS AND DISCUSSION

Speciation of the FeIII-tpena Complexes in Water. Thestructurally characterized [FeIII2(μ-O)(tpenaH)2](ClO4)4

31 isthe solid-state starting material for all experiments. [FeIII2(μ-O)(tpenaH)2]

4+ is formally a hemihydrate of [FeIII(tpena)]2+

with the water equivalent divided between the oxygen atom ofthe oxo bridge and the two protons on the dangling pyridyls ofeach tpena. When dissolved in water, a series of pH-dependentand pH-independent protonation/deprotonation and hydrol-ysis/dehydration reactions give rise to the mono- and dinuclearspecies shown in Figure 1A. The structures of these compoundsare proposed on the basis of the geometries observed in thecrystal structures of iron(III), vanadium(IV), manganese(II),cobalt(III), and chromium(III) complexes, vide supra, and fromparallel electron paramagnetic resonance (EPR), UV/visible,and previously reported Mossbauer26 spectroscopic data.Solution-state EPR spectroscopy can be used to monitor theconcentration of the mononuclear species (Figure S2). At pH≤2.5, the spin content suggests that the iron(III) species aredistributed as ∼50% low-spin (S = 1/2) and ∼10% high-spin (S= 5/2) species. The remaining ∼40% is EPR-silent, consistentwith the presence of the strongly magnetically coupled oxo-bridged complex [FeIII2(μ-O)(tpenaH)2]

4+, with this assign-ment corroborated by the Mossbauer spectrum26 recorded inaqueous solution at pH 4 (Mossbauer signature: δ = 0.47 mms−1 and ΔEQ = 1.60 mm s−1). The low-spin species is themonohydrate of [FeIII(tpena)]2+, namely, [FeIII(OH)-(tpenaH)]2+ (Mossbauer signature: δ = 0.16 mm s−1 andΔEQ = 2.19 mm s−1 26), shown in Figure 1A, where the waterequivalent is divided between the hydroxo ligand and theprotonated pyridyl. The high-spin species is proposed to be aprotonated derivative, [FeIII(OH2)(tpenaH)]

3+ (or, alterna-tively, [FeIII(OH)(tpenaH2)]

3+, not shown in Figure 1A). AtpH >5, a high-spin species, [FeIII(OH)(tpena)]+, develops, withthis evidenced by its signature in the Mossbauer spectrum (δ =0.46 mm s−1 and ΔEQ = 0.71 mm s−1).26 Additional support forthe speciation profile is gleaned from electrospray ionizationmass spectrometry (ESI-MS), where the dominant ionsproduced from [FeIII2(μ-O)(tpenaH)2]

4+ are [FeIII(OH)-(tpena)]+ (m/z 463.13), [FeIII2(μ-O)(tpena)2]

2+ (m/z454.14), and [FeII(tpena)]+ (m/z 446.13).26 The latter gas-phase ferrous complex [FeII(tpena)]+ is not derived from theaqueous solution but is produced during ionization bydehydration and reduction of [FeIII(OH)(tpenaH)]2+.Electrochemical Generation of [FeIV(O)(tpenaH)]2+.

Bulk electrolysis of aqueous solutions containing dissolved[FeIII2(μ-O)(tpenaH)2](ClO4)4 (native pH ∼3.6) using graph-ite felt or reticulated vitreous carbon electrodes produces[FeIV(O)(tpenaH)]2+ within ∼30 min at potentials above +1.3V vs NHE. Charge-transfer curves indicate an approximately90% conversion. The electrolysis is accompanied by a colorchange from yellow-brown to pale-bluish-green, with UV/visible spectra showing the formation of a weak near-IRabsorption band at λmax = 730 nm (ε ∼ 260 M−1 cm−1). Thisband is characteristic for d−d transitions of nonheme iron(IV)oxo complexes,33,34 with the spectrum identical with that for[FeIV(O)(tpenaH)]2+ produced by chemical oxidation usingcerium(IV) in water.26 The extent of conversion and timerequired is dependent on the amount of iron(III) to beelectrolyzed, the potential or current applied, and the size of theworking electrode (WE; Figure 2). During [FeIV(O)-(tpenaH)]2+ formation, a drop in the pH from ∼3.6 to 2.8 is

observed, along with slow gas (H2) evolution at the cathode.The applied voltage is higher than the theoretical value forwater oxidation (1.03 V at pH 3); thus, it is possible that somewater oxidation accompanies the reaction, potentially catalyzedby the complex. However, because [FeIV(O)(tpenaH)]2+ isspectroscopically observable, this competing reaction must besignificantly slower than [FeIV(O)(tpenaH)]2+ formation.

pH Dependence of the Reduction and Oxidation ofthe tpena Complexes of Iron(III). Cyclic voltammograms(CVs) of solutions containing [Fe2(μ-O)(tpenaH)2](ClO4)4were recorded over a pH window of 1.6−9.5. As can be seenfrom Figure 1B, FeIII → FeIV oxidation waves of varyingintensity are present over the whole pH range. The completeset of 24 CVs recorded between pH 1.9 and 9.1 can be found inFigure S3. A Pourbaix diagram summarizing the dominantspeciation and redox events is shown in Figure 1C, and themain electrochemical processes are described below. A moreextensive description of the analysis behind the construction ofthe Pourbaix diagram is provided in the SupportingInformation. Under very acidic conditions (pH ≤2.5), twoFeIII → FeIV waves coexist. The wave at Ep = 1.21 V has a slopeof −104 mV pH−1, which suggests a two-proton/one-electronprocess. This is consistent with [FeIV(O)(tpenaH)]2+ beinggenerated by the oxidation of [FeIII(tpenaH)(OH2)]

3+. TheFeIII → FeIV wave at Ep = 1.36 V shows a pH dependence of−67 mV pH−1, suggesting a one-proton/one-electron oxida-tion, and is assigned to the oxidation of [FeIII(tpenaH)(OH)]2+

to [FeIV(O)(tpenaH)]2+. The change in the slope of the FeIII →FeIV peak potential at pH 3.5 suggests a change in theprotonation state of either [FeIII(tpenaH)(OH)]2+ or[FeIV(O)(tpenaH)]2+. At intermediate pH (pH 4−7), the

Figure 2. (A) Time-resolved UV/visible spectra showing theformation of [FeIV(O)(tpenaH)]2+ during electrolysis of aqueoussolutions of [Fe2(O)(tpenaH)2](ClO4)4 (0.8 mM per Fe and 0.1 MNaClO4) at +1.42 V over a 10.1 or 15.8 cm2 graphite felt WE. Spectrawere recorded every ∼5 min. Inset: Time traces for λ = 730 nm,showing the effect of varying size of the WE. (B) AccompanyingFaradaic charge-transfer curves.

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14939

higher potential FeIII → FeIV wave shows no pH dependence,with this result consistent with the oxidation of [FeIII(OH)-(tpena)]+ to [FeIV(O)(tpenaH)]2+. Over this same pH range,there is a shift in the pH dependence of the reversible FeIII/FeII

couple from E1/2 = 0.63 to 0.47 V, which suggests a change inthe protonation state of the species involved in this reaction(e.g., [FeII(tpena)(OH2)]

+ to [FeII(OH)(tpena)] with a pKavalue of 5.3). Under alkaline conditions (pH >7), the FeIII →FeIV peak again becomes pH-dependent with a slope of −54mV pH−1. This can be rationalized by [FeIII(OH)(tpena)]+

being oxidized to [FeIV(O)(tpena)]+. A pKa value of 6.7 for[FeIV(O)(tpenaH)]2+ is assigned to deprotonation of thependant protonated pyridyl arm to give the conjugate base[FeIV(O)(tpena)]+.The potential for protonation of the electrochemically

generated [FeIV(O)(tpenaH)]2+ was explored by pH manipu-lation (using HClO4 or NaOH) in the pH range of 1.6−2.4.The absence of any changes in the UV/visible spectra or decayrate suggests that protonation of the oxo ligand to give aniron(IV) hydroxide complex, e.g., [FeIV(OH)(tpenaH)]3+, doesnot occur. Similarly, Borovik and co-workers have failed to findany evidence for protonation of the oxo ligand in [FeIV(O)-(H3buea)]

− (buea3− = tris[2-(N′-tert-butylureaylato)ethyl]-aminato).35 The intensity of the FeIII → FeIV wave correlateswith the amount of conversion that takes place during bulkelectrolysis at a given pH; the highest concentrations ofiron(IV) oxo species are generated when the initial pH isaround ∼2−3. Above this and up to neutral pH, lowerconcentrations of iron(IV) oxo species are produced. If the pHis adjusted to above 7, no observable amounts of iron(IV) oxospecies are generated by bulk electrolysis. At first glance, this issurprising because CVs at these pH values show relativelyintense FeIII → FeIV waves. This effect can, however, berationalized by considering that observable iron(IV) oxo speciesare a balance of both the electrochemical formation rate andsubsequent decay, which, in the absence of substrates,presumably occurs due to water oxidation. The decay of theiron(IV) oxo species was followed by the generation of[FeIV(O)(tpenaH)]2+ at pH ∼ 3−4 and then monitoring of theloss of absorbance at λmax ∼ 730 nm after adjustment of the pH.At pH 3−4, the half-life of [FeIV(O)(tpenaH)]2+ is around

τ1/2 ∼ 30 min; as the pH is increased, the decay rate alsoincreases, with the half-life at pH 6−7 being shorter than τ1/2 ∼5 min. The inability to accumulate the iron(IV) oxo species athigher pH is thus presumably due to the fact that [FeIV(O)-(tpena)]+ is more reactive than its acid congener, [FeIV(O)-(tpenaH)]2+. The rate constants for water oxidation of ∼1.2 ×10−3 and 4 × 10−4 M−1 s−1 were deduced, respectively, at pH6−7 and 3−4. The assignment of the rate constant for thereaction of [FeIV(O)(tpena)]+ with H2O is, however, uncertainbecause the observed self-decay rate of the iron(IV) oxo speciesincreases markedly at pH >8 because of much more rapidoxidation of OH− than H2O, with a rate constant for thereaction of [FeIV(O)(tpena)]+ with OH− determined to be 3 ×102 M−1 s−1, i.e., over 5 orders of magnitude greater than therate for H2O oxidation (detailed derivation available in theSupporting Information). These results are in accordance withthe prediction from the Pourbaix diagram of a pKa = 6.7 fordeprotonation of the dangling pyridyl group to give [FeIV(O)-(tpena)]+ and provide an explanation for the inability togenerate bulk amounts of [FeIV(O)(tpena)]+ at pH >7. Theresults further suggest that the second coordination sphere

pyridyl base has a strong modulating effect on the reactivity ofthe iron(IV) oxo moiety in this system.

Scope of Substrate Oxidation by ElectrochemicallyGenerated [FeIV(O)(tpenaH)]2+ and [FeIV(O)(tpena)]+.When water-soluble organic compounds are added to solutionsof electrochemically generated [FeIV(O)(tpenaH)]2+ and[FeIV(O)(tpena)]+, the decay is accelerated by several ordersof magnitude. Figure 3A shows the disappearance of the 730nm absorption before and after the addition of isopropylalcohol at pH 4. The decay is substantially faster (τ1/2 < 5 s)when phenol or 4-chlorophenol is added (Figure S4),suggesting that it is a C−H bond in the aliphatic alcoholsand a O−H bond in the aromatic alcohols, respectively, that aretargeted by the iron(IV) oxidant, with this result consistentwith their relative bond dissociation enthalpies (BDEs). Gaschromatography (GC) analyses show that the stoichiometricreactions with both aliphatic and aromatic alcohols produce thederivative ketones as the main products (Figures S5−S9) withthe second-order rate constants summarized in Figure 3B andTable 1. Solutions were buffered with acetate or phosphatebuffers at pH 4 and 7, respectively. The oxidation oftetrahydrofuran (THF) and cyclohexanol are also relativelyfast with rate constants of 2.7 × 10−2 M−1 s−1 and 4.7 × 10−2,respectively. Given that the BDE of the weakest C−H bond inTHF is 92.0 versus 92.8 kcal mol−1 for cyclohexanol,36 it issurprising that the oxidation of cyclohexanol is almost twice asfast as the oxidation of THF. One rationalization for this is adifference in the kinetic barriers. The presence of a hydroxylgroup in a substrate might mean that a transition-state ironoxidant−substrate supramolecular adduct for cyclohexanolmight be more easily accessed than one for THF. The ratesfor oxidation of the simple alcoholsisopropyl alcohol (i-PrOH; 1.6 × 10−2 M−1 s−1), ethanol (EtOH; 5.5 × 10−3 M−1

s−1), and methanol (MeOH; 7.1 × 10−4 M−1 s−1) at pH 4follow the order predicted from their BDEs.36 The spread inthe magnitude of the rate constants is larger than what isexpected from the Bell−Evans−Polanyi (BEP) relationship.Corroborating the increase in the decay rate at higher pH, inthe absence of substrates, the oxidative power of [FeIV(O)-(tpena)]+ is also enhanced at higher pH. Thus, the rate ofmethanol oxidation at pH 7 triples from that at pH 4 to k2 = 2.0× 10−3 M−1 s−1.Changes to the protonation state of the substrate also

influence the reactivity, with the rate constant for the oxidationof formate/formic acid to carbon dioxide (CO2; Figure 3C)observed to increase from 0.30 M−1 s−1 at pH 3 to 1.2 M−1 s−1

at pH 5.5. The pH dependence is quantitatively consistent witha higher reactivity of HCOO− compared to HCOOH (pKa =3.75). We speculate that the positive charge of [Fe(O)-(tpenaH)]2+ and the hydrogen-bonding potential are advanta-geous for enhanced reactivity toward negatively chargedHCOO− over neutral HCOOH. The energetics of theelectron-transfer process are likely to be similar; however, ahigher outer-sphere association constant for the 2+/1− ion pairwould be expected relative to the 2+/neutral case.45 Formationof CO2 as a result of formate oxidation was confirmed by using14C-labeled formate, with the generated 14CO2 recovered in aNaOH trap and quantified by scintillation counting, unambig-uously proving that the CO2 evolved derives from the substrateand not from ligand breakdown.

Mechanism. While we observe a tendency for faster ratesfor the oxidation of substrates with weaker BDEs by[FeIV(O)(tpenaH)]2+/[FeIV(O)(tpena)]+, as mentioned

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14940

above, a BEP correlation does not apply. This outcomecontrasts with the BEP relationships that have been reportedfor hydrocarbon and alcohol oxidations by iron(IV) oxocomplexes in nonaqueous solution.46−49 The BEP relationshipbuilds on the assumption that the entropic contributions to theactivation barriers, for a series of reactions using closely relatedsubstrates, can be ignored. This assumption may beinappropriate for the Fe-tpena system. The mechanisticalternative to a HAT pathway, a proton-coupled electron-transfer mechanism,50,51 is easily ruled out because thismechanism implies the formation of an iron(III) oxo complexand alcohol radical cations (ROH•+), both unlikely inter-mediates. Mayer and co-workers have shown that the Marcuscross-relation, which uses bond dissociation free energies(BDFEs) instead of BDEs, provides a better description ofHAT reactivity in such cases.52−54 Unfortunately, the BDFEsfor most of the substrates used in this study are not known. Toprobe the size of the entropic activation barrier, we measuredthe rates for the i-PrOH oxidation by [FeIV(O)(tpenaH)]2+ atdifferent temperatures. As shown in Figure 3D, this yielded anegative activation entropy of ΔS⧧ = −23 ± 4 cal mol−1 K−1 (|TΔS⧧| = 7 kcal mol−1 at 25 °C vs ΔH⧧ = 13 ± 1 kcal mol−1),suggesting a significant entropic component to these HATreactions. The combination of an aqueous medium, a danglingpyridine/pyridinium, and the carboxylate group furnisheshydrogen-bonding opportunities. These have the potential oflowering the entropic penalty associated with desolvation of asubstrate in the transition state.54 It is worth noting that,although slower by up to 107 times, the relative order for therate constants for the oxidation of the various substrates by[FeIV(O)(tpenaH)]2+ is the same as those reported for thehydroxyl radical (HO•; Figure 4). Our interpretation is that aHAT mechanism is at play for both oxidants. The transitionstate for HAT reactions by HO• involves the transfer of anegative charge to the oxygen atom of HO• and positive chargegeneration on the hydrogen atom of the substrate. This isstabilized by interaction with a polar solvent, and particularly byhydrogen bonding with water, leading to significant rateacceleration in water.55 It is conceivable that the tpena-derivediron(IV) oxo moiety shows radical character analogous to thatof HO• and also participates in HAT chemistry54 (e.g., an“iron(III) oxyl”-type species). Involvement of the lone pair ofthe uncoordinated pyridine in stabilizing the positive charge onthe substrate H in the transition state may be contributing tothe increased rates observed at higher pH (Scheme 2). Afterhydrogen-atom abstraction, the immediate FeIII-OH productcan itself be intramolecularly stabilized through hydrogenbonding to the second coordination sphere base. Formic acid isa notable exception to the correlation observed in Figure 4,exhibiting a rate constant for the reaction with [FeIV(O)-(tpenaH)]2+ of 0.2 M−1 s−1, which is many orders of magnitudegreater than that predicted from the correlation (10−7 M−1 s−1).The reaction of HO• with formic acid is proposed to involvethe acidic hydrogen atom in a cyclic hydrogen-bonded{HO•:HOOCH} adduct transition state.56 An analogousinteraction is, however, not feasible for the [FeIV(O)-(tpenaH)]2+ oxidant, which is instead more likely to reactwith the weaker C−H bond (BDE of 96.6 kcal mol−1; cf. 112.3kcal mol−1). Thus, we can provide a rationalization for why thisparticular substrate does not fit into the parallel of reactivity for[FeIV(O)(tpenaH)]2+ and HO•. It should be noted that theformate ion, which has only the C−H hydrogen available for

Figure 3. (A) Time-resolved UV/visible spectra of the reaction of[FeIV(O)(tpenaH)]2+ (0.8 mM per Fe, 0.1 M NaClO4) with i-PrOH(60 equiv). Inset: Time trace for λ = 730 nm before and after theaddition of i-PrOH and exponential fits to extract the observed rateconstants. (B) Plot of observed rate constants corrected forbackground decay (k′obs = kobs − kdecay) against the substrateconcentrations to extract the second-order rate constants (k2). Thedotted line is for i-PrOH oxidation by [FeIV(O)(bztpen)]2+. (C) Plotof the apparent rate constants for the pH-dependent formate/formicacid oxidation [kapp = (kobs − kdecay)/CHCOOH] with the fitted curve forkHCOOH = 0.20 ± 0.02 M−1 s−1, kHCOO− = 1.2 ± 0.1 M−1 s−1,pKa(HCOOH) = 3.75. The dashed line shows 95% confidenceinterval. (D) Eyring plot for the i-PrOH oxidation by [FeIV(O)-(tpenaH)]2+ at pH 4.

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14941

abstraction, cannot proceed through such a pathway and doesfit with the correlation in Figure 4.Comparison of the Iron(IV) Oxo Potencies in Aqueous

Oxidations of Organic Substrates. The rates we measurefor the oxidation of benzyl alcohol by electrochemicallygenerated [FeIV(O)(tpenaH)]2+ are higher than those thathave been reported for chemically generated [FeIV(O)-(bztpen)]2+ 25 and [FeIV(O)(N4Py)]2+ 25 (Scheme 1) inacetonitrile. Unfortunately, apart from the aforementioned

Fe-TAML system, there is a paucity of oxidative reactivitystudies of iron(IV) oxo systems in aqueous solution, and theknown iron(IV) oxo complexes have not been synthesized inthis medium. To address this, Wang et al.57 used isolatedpowders of [FeIV(O)(N4Py)](ClO4)2 recovered from itssynthesis in acetonitrile and, subsequently, redissolved thepowder in water at pH ∼7. When MeOH and THF wereadded, the decay of [FeIV(O)(N4Py)]2+ occurred with rateconstants of k2 = 5 × 10−5 M−1 s−1 and k2 = 1.7 × 10−3 M−1 s−1,respectively. A comparison with our data in water indicates that

Table 1. Second-Order Rate Constants for Substrate Oxidations by [FeIV(O)(tpenaH)]2+ at pH 4 and Those Compiled from theLiterature for HO• Oxidation of the Same Substrate along with Gas-Phase BDEs for the Weakest C−H or O−H Bond, asIndicated by the Bold Hydrogen Atom

substrate k2/M−1 s−1 BDEa/kcal mol−1 k2(HO

•)/M−1 s−1

OxidanesHOH water 4 × 10−4 118.8OH− hydroxide 3 × 102 110.2b 1.2 × 1010 c

Aliphatic AlcoholsCH3OH 105.2CH3OH methanol 7.1 × 10−4 96.1 9.7 × 108 d

CH3CH2OH ethanol 5.5 × 10−3 95.9 1.9 × 109 d

(CH3)2CHOH isopropyl alcohol 1.6 × 10−2 94.8 1.9 × 109 d

(CH2)5CHOH cyclohexanol 4.7 × 10−2 92.8 1.7 × 109 d

Heterocyclic Ether(CH2)4O tetrahydrofuran 2.7 × 10−2 92.1 4.0 × 109 d

Aromatic Alcohols(C6H5)CH2OH benzyl alcohol 2.1 × 100 79.2 8.4 × 109 d

(C6H5)OH phenol 3.9 × 102 90.4e 1.4 × 1010 f

(C6H4Cl)OH 4-chlorophenol 2.4 × 102 90.7e,g 9.3× 109 h

Aliphatic AcidHCOO− formate 1.2 × 100 3.2 × 109 d

HCOOH formic acid 2.0 × 10−1 96.6HCOOH formic acid 112 1.3 × 108 d

AlkaneC10H15OSNaO3

− i 10-camphorsulfonate 4.7 × 10−2 4.1× 109 d,j

Xanthene DyeC28H30N2O3 rhodamine B 6.0 × 101 1.7 × 1010 k

aGas-phase BDE for the weakest bond, taken from Luo36 unless indicated otherwise. bValue given in Blanksby and Ellison,37 from the work of Ruscicet al.38 cValue from Buxton.39 dValues from the review by Buxton et al.40 eThe BDE is for the O−H bond. fValues from Land and Ebert.41 gAverageof the five listed in Luo.36 hValue from Stafford et al.42 iSodium salt. jValue for camphor used was an estimate. kAverage of the values reported byBrutz and Sommermeyer43 and Kucherenko.44

Figure 4. Correlation between the second-order rate constants foroxidation reactions of [FeIV(O)(tpenaH)]2+ and the hydroxyl radical.See Table 1 for HO• rate constants and their sources. HCOOH is notincluded; see the text.

Scheme 2. Catalytic HAT from Substrate to Iron(IV)Oxidant and the Proposed Resultant Product FeIII-OHComplexa

aThe product organic radical will initiate a chain reaction, leadingultimately to CO2 (reaction not balanced). Protons generated arereduced at the cathode.

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14942

[FeIV(O)(tpenaH)]2+ oxidizes MeOH and THF at least >12and ∼17 times faster, respectively, than [FeIV(O)(N4Py)]2+.To further evaluate the potency of [FeIV(O)(tpenaH)]2+, we

carried out bulk electrolysis on the closely related system[FeII(Cl)(bztpen)](ClO4) (in which the carboxylate arm isreplaced by a benzyl arm; Scheme 1) at +1.42 V in water, with10% acetonitrile added for solubility reasons. We found that,even though a two-electron oxidation of the iron is required inthis case, [FeIV(O)(bztpen)]2+ could be generated. The CV isshown in Figure 5A. Replacement of the carboxylato donor by a

pyridine donor results in a slight blue shift of the d−dabsorption to λmax = 720 nm (ε ∼ 220 M−1 cm−1) compared tothat of [FeIV(O)(tpenaH)]2+. The Faradaic charge transferredindicates ∼91% conversion and a three-electron-transferprocess, with the last electron equivalent assumed to be dueto Cl− oxidation (eq 1).

+

→ + + +

+

+ + −

[Fe (CI)(bztpen)] H O

[Fe (O)(bztpen)] 2H12

CI 3e

II2

IV 22 (1)

The simplicity of the CVs recorded for [FeII(Cl)(bztpen)]+

is striking compared to those for Fe-tpena. This is notsurprising given the more complicated speciation and pHprofiles for Fe-tpena. The electrochemically generated[FeIV(O)(bztpen)]2+ is able to oxidize i-PrOH (Figures 5Band S16); however, the rate constant of 8.6 × 10−3 M−1 s−1 (atpH 4) is only half the rate constant obtained for the reactionwith [FeIV(O)(tpenaH)]2+.

Electrocatalytic Degradation of Cyclohexanol. [Fe-(tpena)]2+ catalyzes the degradation of cyclohexanol underelectroanodic conditions with turnover. Using a bulk graphitefelt electrode as the charge carrier, the oxidation of cyclo-hexanol (∼80 mM, emulsion) proceeds under both galvano-static and potentiostatic conditions. When electrolyzed at 1.51V in the presence of 0.8 mol % Fe-tpena, GC analysis showedthat 44% of cyclohexanol was oxidized after 5 h (Figures S17and S18). By comparison, only 18% was oxidized in the sametime period in the absence of a catalyst. After 16 h, ∼81% of theoriginal cyclohexanol content had disappeared from thecatalyzed sample. In these solutions, cyclohexanone amountingto 3−5% of the mass balance was detected by GC. Thecyclohexanone concentration reaches a steady state despiteongoing cyclohexanol loss, suggesting that cyclohexanone isfurther oxidized with the possibility of eventual totalmineralization, as depicted in eq 2.

+ → + ++ −C H OH 11H O 6CO 34H 24e6 11 2 2 (2)

This observation is consistent with other studies of chemicaland electrochemical cyclohexanol oxidations where the ring-opened carboxylic acids (6-hydroxyhexanoic acid, 6-oxohex-anoic acid, and adipic acid) have been identified as majorproducts of the oxidation of cyclohexanol.58−60 We assumetherefore that the missing mass balance in the gas chromato-grams is due to the formation of these GC-incompatiblecarboxylic acids or total mineralization (producing undetectedCO2). NMR analysis supports the notion of total mineralizationoccurring because the spectra of 24 h electrocatalysisexperiments show no other signals apart from those due towater (Figure S17). The charging curves show a significant 23%enhancement (after 5 h; Figure S18) in the charge transferredin the catalyzed experiments, compared to the control. Thecatalytic charge transferred after 16 h gives a turnover numberof 12 for Fe-tpena.The CV of cyclohexanol in water reveals an oxidation wave

with onset at E ≥ 1.32 V. Upon titration with increasingconcentrations of [Fe2(O)(tpenaH)2]

4+, a clear enhancementin this current is observed (Figure 6A). A plot of the catalyticcurrent (icat) at 1.50 V against the iron concentration gives alinear correlation (Figure 6A, inset). Likewise, the CVs of[Fe(tpena)]2+ with increasing concentrations of cyclohexanolreveal that icat increases linearly with the square root of thecyclohexanol concentration until the point of maximal solubilityis reached (Figure 6B). The concentration dependenceindicates that the reaction rate has the form given in eq 3and that it can be modeled as a kinetically controlled two-stepcatalytic process (EC′ mechanism).61−64 Essentially, the activeiron-based oxidant is generated in a diffusion-controlled(heterogeneous) electron transfer between the resting catalystand the electrode surface. The catalytically competent productspecies then diffuses into the solution and reacts with thesubstrate in a kinetically controlled homogeneous reaction. Inturn, the iron(III) resting catalyst is regenerated. The scan rate(v) dependence for an EC′ mechanism is given by eq 4.

= krate [Fe (O)][ROH]catIV

(3)

=ii

nn

k RTn Fv0.4463

[ROH]cat

p

cat

p

cat

p (4)

The anodic peak potential (ip) for the reversible[FeIII(tpenaH)(OH)]2+/[FeII(tpenaH)(OH)]+ redox couple is

Figure 5. (A) CVs of solutions containing [FeII(Cl)(bztpen)](ClO4)(unbuffered, pH 4.7). (B) Time-resolved UV/visible spectra of thereaction of [FeIV(O)(bztpen)]2+ (0.8 mM) with i-PrOH (120 equiv)in a pH 4 acetate-buffered electrolyte.

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14943

diffusion controlled (Figure S19) and therefore a good choicefor normalizing icat according to eq 4. A fit of the values of icat/ip, extracted from the CVs of cyclohexanol in the presence ofhydrated [Fe(tpena)]2+ and recorded over a range of scan rates(Figure 6C), gives a turnover frequency of kcat = 3.1 s−1 for theinitial oxidation of cyclohexanol by [FeIV(O)(tpenaH)]2+.Stability of [FeIV(O)(tpenaH)]2+. The clear evidence of

turnover provided here indicates that the Fe-tpena system mustbe reasonably robust; however, to investigate this further, wesubjected a solution containing [Fe2(O)(tpenaH)2](ClO4)4 tosix cycles of repeated electrochemical oxidation and decay(Figure S20). The UV/visible spectra of the generated[FeIV(O)(tpenaH)]2+ did not change throughout the cycles,with all spectra consistent with [FeIV(O)(tpenaH)]2+ revertingback to an iron(III) resting state (with an isosbestic point at502 nm); however, the concentration diminished by an averageof 5% with each cycle. Importantly, however, crystalline[Fe2(O)(tpenaH)2](ClO4)4 can be reisolated from multiplyelectrolyzed solutions (purity checked by X-ray diffraction andESI-MS spectra; Figure S21). The ESI-MS spectra of thesupernatant solutions show no indication of demetalation orligand breakdown, and GC did not reveal peaks due to lightorganics. The absence of identifiable breakdown products orspectral changes does not, however, exclude the possibility thatslow catalyst decomposition occurs because any organicfragments are likely to be mineralized.

■ CONCLUSIONS

We have demonstrated that the nonheme iron(IV) oxocomplex acid−base congeners [FeIV(O)(tpenaH)]2+ and[FeIV(O)(tpena)]+ can be generated electrochemically inaqueous solution over a relatively large pH window (2−8).These complexes can oxidize a broad range of substratescontaining C−H and O−H bonds under both acidic andneutral conditions in water. The oxidation of aqueous solutionsof formate to CO2 and bleaching of rhodamine B suggests thatthe iron(IV) oxo species can oxidatively degrade otherwiserecalcitrant organic pollutants and, as such, the results arehighly pertinent for the development of electrocatalytic waterpurification. Under the conditions and potentials applied,

[Fe2(O)(tpenaH)2]4+ appears to be relatively robust, with

[FeIV(O)(tpenaH)]2+ electrochemically regenerated repeatedly.Catalytic activity in the electrooxidation of cyclohexanol wasobserved with at least 12 turnovers.The weak field carboxylate donor, positioned cis to the oxo

group, likely plays an important role in the enhanced reactivityobserved for [FeIV(O)(tpenaH)]2+/[FeIV(O)(tpena)]+ com-pared to the aminopolypyridyl [FeIV(O)(bztpen)]2+ and[FeIV(O)(N4Py)]2+ analogues. Interestingly, cis-carboxylatedonors are commonly found in the active sites of nonhemeiron enzymes involved in biological oxidation reactions. Despitethis fact, monocarboxylate ligands have rarely been employed inbiomimetic iron chemistry. In addition, the positively chargedFe-tpena complexes show remarkable flexibility with respect tometal geometry and ligand protonation and denticity. In water,a second coordination sphere base is enabled and catalyticactivity is observed over a wide pH range.We have noted a parallel in the oxidative activity trend of

[FeIV(O)(tpenaH)]2+ with that of the HO• radical. Notsurprisingly given their relative sizes, the rates of substrateoxidation by the iron-based system are orders of magnitudeslower. This might not be of practical disadvantage with respectto implementation in the electrocatalytic water purificationbecause [FeIV(O)(tpenaH)]2+ can be generated in concen-trations many orders of magnitude higher than HO•. Adrawback, however, is that the time required to generate[FeIV(O)(tpenaH)]2+ electrochemically is much longer than itschemical generation by cerium(IV) in water.26 The latter isinstantaneous, whereas it approaches 30 min for a steady-stateconcentration using electrochemical means (Figure 2A).Clearly, this will result in much slower catalysis, withreoxidation of FeIII to FeIV being rate-determining. Ultimately,for practical purposes, immobilization onto electrode materialsis likely to be necessary, and it can be expected to improve thereoxidation efficiency. Immobilization also offers the practicaladvantage of preventing discharge of the complex and/orligand, which is undesirable from both an efficiency standpointand the potential of trace amounts to act as a contaminant in itsown right, which would need to be established in future work,e.g., assessing for any endocrine-disrupting potential.65

Figure 6. (A) CVs of cyclohexanol (0.04 M) titrated with increasing concentrations of [FeIII2(μ-O)(tpenaH)2]4+ (0−1 mM per Fe, 50 mV s−1, and

0.1 M NaClO4). Inset: Plot of the catalytic current for increasing concentrations of Fe-tpena, extracted from the CVs. (B) CVs of[FeIII2(O)(tpenaH)2](ClO4)4 (0.8 mM per Fe, 0.1 M NaClO4) titrated with increasing concentrations of cyclohexanol (0−0.18 M). Inset: Plot ofthe catalytic current for increasing concentrations of cyclohexanol, extracted from the CVs. (C) CVs of mixtures of [FeIII2(μ-O)(tpenaH)2]

4+ andcyclohexanol (0.8 mM Fe, 40 mM cyclohexanol, and 0.1 M NaClO4) recorded at different scan rates (2.5−200 mV s−1). Inset: Plot of icat/ipeak, takenat 0.6 and 1.50 V, respectively, against the inverse square root of the scan rate.

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14944

Although immobilization is a large and unpredictable hurdle,the discovery of a candidate molecular iron-based oxidant thatcan be electrochemically activated and that can operate underboth neutral and acidic conditions, with insignificant back-ground oxidation of water, represents a significant step towardfundamentally new and potentially sustainable methods forwater remediation.

■ EXPERIMENTAL SECTIONMaterials and Preparations. Commercially available reagents

were purchased from Sigma-Aldrich and used without furtherpurification (unless otherwise noted). 14C-labeled sodium formatewas purchased from American Radiolabeled Chemicals. Rhodamine B(laser grade, 99%) was purchased from Acros. [Fe2O(tpenaH)2]-(ClO4)4(H2O)2 and [Fe(bztpen)(Cl)](ClO4)2(H2O) were preparedas described previously.31,66 Standard commercial electrodes andelectrochemical cells were purchased from ALS Japan or BASi, eitherdirectly or through retailer IJ Cambria UK. Graphite felt (99%, G100PAN-based battery felt, AVCarb) and Nafion-117 sheets (Dupont)were purchased from FuelCellStore.com. Sealed nafion bags wereprepared by folding sheets together and heat-fusing the edges using animpulse sealer, with one side left open to allow insertion of theelectrode mesh and electrolyte. Graphite rods (1 mm, 99.95%),reticulated vitreous carbon (100 ppi glassy carbon), and platinummesh (1 × 1 cm2, 99.9%) were purchased from GoodFellow. Platinumand titanium coil electrodes were made by manually coiling metal wire(∼20 cm, Alfa Aesar) around a 0.8 cm template. Micro Ag/AgClreference (“leak-free”) electrodes were made by casing a AgCl wire in atube filled with KCl electrolyte. The housing consisted either of a 0.3mm poly(tetrafluoroethylene) tube fitted with a microporous frit or a0.3-mm-diameter glass tube fitted with a heat-sealed graphite pin (bothcustom-made).Instrumentation. UV/visible spectra were recorded on an Agilent

Cary 60 or an Agilent 8453 spectrophotometer using 1 cm quartzcuvettes. Variable-temperature spectra were recorded using a UnisokuScientific Instruments adaptor. EPR spectra were recorded on a BrukerEMX Plus X-band continuous-wave spectrometer using an ER 4103TM110 cavity. ESI-MS was performed on a nanospray Bruker micro-OTOF-Q II spectrometer. The pH was measured using a MetrohmPHM240 or a Hanna HI2000 pH meter calibrated with IUPACbuffers. GC analyses of reaction solutions were performed on a HP6890 gas chromatograph equipped with a flame ionization detectorand a Zebron ZB-WAX column. Products formed in the stoichiometricreactions were identified by comparing the chromatograms ofcommercial standards.Electrochemistry. Electrochemical experiments were performed

on either a CH Instruments 650D or an Eco Chemie AutolabPGSTAT10 electrochemical analyzer. Experiments were generallyexposed to the ambient atmosphere and performed using solutionvolumes in the range of 3−10 mL. Solutions were degassed by brief N2or argon bubbling prior to commencing experiments. Cyclicvoltammetry was performed using a standard three-electrode setup,with a glassy carbon disk working electrode (0.28 cm2), a platinumwire counter electrode, and a Ag/AgCl (3 M KCl) reference electrode.The working electrode was cleaned by polishing to a shine with 0.05μm alumina followed by sonication in acetone and water, followed byrepeated cyclic voltammetry sweeps in a blank electrolyte solution.Extraction of the peak potentials was aided by inspection of the filteredfirst- and second-order derivatives of the anodic and cathodic sweeps.This is especially helpful for identifying oxidation/reduction waves ofweak intensity that may otherwise be disguised under catalytic orcapacitive currents. Bulk electrolyses (galvanostatic or potentiostatic)were performed under stirring using primarily a graphite felt WE (3.0× 1.5 × 0.4 cm3, Aproj = 12.6 cm2) pierced with a graphite rod or aplatinum wire as the connector. A platinum or titanium coil electrodewas used as the counter electrode. The counter electrode was housedin a separate fritted glass chamber filled with ∼400 μL of electrolyte,submerged in the bulk electrolysis cell. For electrocatalytic experi-ments, a larger 6.0 × 1.5 × 0.4 cm3 (Aproj = 24 cm2) felt electrode was

used as the charge carrier, with a platinum mesh (1 × 1 cm2) counterelectrode housed in a sealed nafion bag. The reference electrode for allelectrolysis experiments was a micro Ag/AgCl electrode (3 M KCl).All potentials have been given versus the normal hydrogen electrode(NHE). Experimental potentials measured versus the micro Ag/AgClelectrode (3 M KCl) were referenced to the NHE by adding +0.21 V.Potentials measured versus the micro Ag/AgCl electrode (3 M KCl)were referenced to the NHE by adding +0.25 V. These calibrationvalues were determined using the redox potential of the[FeIII(CN)6]

3−/[FeII(CN)6]4− couple (10 mM, 0.5 M KCl, E1/2 =

0.456 V vs NHE)67 as the standard.Kinetic Measurements. [Fe2O(tpenaH)2](ClO4)4 (0.8 mM per

Fe) in a pH 4 buffered electrolyte (0.05 M acetate and 0.1 M NaClO4)was electrolyzed at a constant potential (1.42 V vs NHE) using thesetup described above until the measured current had dropped from∼10 to ∼0.5 mA. The resulting pale-green solutions were quicklyfiltered, and 3000 μL was transferred to a cuvette. The kinetics of thereaction was determined by monitoring the loss of the iron(IV) oxocomplex at its near-IR absorbance peak at 730 nm. The UV/visiblespectrum was first monitored for 120−200 s to determine thebackground decay rate, after which an aliquot of substrate (3−400equiv) from a concentrated aqueous stock solution was added and thereaction rate measured. The background decay rate (kdecay) and thesubstrate reaction rate (kobs) were determined by a least-squares fittingof the absorbance at 730 nm to an expression appropriate for(pseudo)-first-order exponential decay; i.e.

= + − −k tAbs Abs (Abs Abs ) exp( )t f 0 f obs

Second-order rate constants (k2) were extracted from linearregressions of plots of kobs′ = kobs − kdecay against the substrateconcentrations. For the pH-dependent formate/formic acid reactivitystudies, unbuffered media were used, and the pH was adjustedmanually with dilute NaOH prior to starting the UV/visiblemeasurements. The second-order rate constant for the oxidation ofi-PrOH by [FeIV(O)(bztpen)]2+ was determined in pH 4 acetate-buffered electrolytes containing 10 vol % acetonitrile added forsolubility reasons. The degradation of rhodamine B by [FeIV(O)-(tpenaH)]2+ was semiquantitatively followed by UV/visible spectros-copy using the approach described by Miller et al.68 with molarabsorption coefficients taken from Watanabe et al.69

Radiolabeling Experiments. 14C-labeled sodium formate wasspiked at 40 nM in experimental solutions, with the remaining desiredformate concentration made up with unlabeled sodium formate.Experiments were performed in an enclosed apparatus with theexhaust gas passing through a 1 M NaOH trap to recover 14CO2 asNa2

14CO3. The concentrations of 14C-formate and Na214CO3 were

determined by scintillation counting using an Ultima Gold scintillationcocktail (PerkinElmer) and a PerkinElmer Tri-Carb 2910TR liquidscintillation analyzer.

Caution! Perchlorate salts of metal complexes are potentially explosiveand should be handled with caution in small quantities. 14C is a β-emittingisotope that should be handled in a way minimizing this potentialradiation risk and in accordance with any applicable legal requirements.

■ ASSOCIATED CONTENT*S Supporting InformationThe Supporting Information is available free of charge on theACS Publications website at DOI: 10.1021/acs.inorg-chem.7b02208.

Further experimental details, EPR spectra and CVs ofaqueous FeIII-tpena as a function of the pH, an extensivedescription of the construction of the Pourbaix diagramin Figure 2C, kinetic analysis for the oxidation ofphenols, benzyl alcohol, and isopropyl alcohol, GC tracesfrom stoichiometric ethanol, isopropyl alcohol, 1-butanol, 2-butanol, and benzyl alcohol oxidations by[FeIV(O)(tpenaH)]2+, GC and 1H NMR spectra fromanodic oxidation of cyclohexanol catalyzed by Fe-tpena,

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14945

Randles−Sevic plot of the [FeIII/II(tpenaH)(OH)]+/2+

peak currents, ESI-MS spectrum of recovered catalyst,and data for repeat electrolysis and decay cycles of thecatalyst (PDF)

■ AUTHOR INFORMATIONCorresponding Author*E-mail: mckenzie@sdu.dk.ORCIDChristopher J. Miller: 0000-0003-3898-9734T. David Waite: 0000-0002-5411-3233Christine J. McKenzie: 0000-0001-5587-0626NotesThe authors declare no competing financial interest.

■ ACKNOWLEDGMENTSThis work was supported by the Danish Council forIndependent Research, Natural Sciences (Grant 4181-00329to C.J.McK.), and the Australian Research Council, DiscoveryProgram (Grant DP150102248 to T.D.W.).

■ REFERENCES(1) World Health Organization. UN-Water Global Analysis andAssessment of Sanitation and Drinking Water (GLAAS) Report, 2014.(2) Jacobsen, C. S.; Sørensen, S. R.; Juhler, R. K.; Aamand, J. GEUSreport 49: Emerging contaminants in Danish groundwater; GeologicalSurvey of Denmark and Greenland, 2005.(3) Brusch, W. M.; Rosenbom, A. E.; Juhler, R. K.; Gudmundsson, L.;Nielsen, C. B.; Plauborg, F.; Olsen, P. The Danish Pesticide LeachingAssessment Programme: Monitoring results May 1999−June 2012;Geological Survey of Denmark and Greenland, 2013.(4) Daughton, C. G.; Ternes, T. A. Pharmaceuticals and personalcare products in the environment: agents of subtle change? Environ.Health Perspect. 1999, 107, 907−938.(5) Toppari, J.; Larsen, J. C.; Christiansen, P.; Giwercman, A.;Grandjean, P.; Guillette, L. J.; Jegou, B.; Jensen, T. K.; Jouannet, P.;Keiding, N.; Leffers, H.; McLachlan, J. A.; Meyer, O.; Muller, J.; Meyts,E. R.-D.; Scheike, T.; Sharpe, R.; Sumpter, J.; Skakkebaek, N. E. MaleReproductive Health and Environmental Xenoestrogens. Environ.Health Perspect. 1996, 104, 741−803.(6) Boyle, B.; McConkey, R.; Garne, E.; Loane, M.; Addor, M.;Bakker, M.; Boyd, P.; Gatt, M.; Greenlees, R.; Haeusler, M.;Klungsøyr, K.; Latos-Bielenska, A.; Lelong, N.; McDonnell, R.;Metneki, J.; Mullaney, C.; Nelen, V.; O’Mahony, M.; Pierini, A.;Rankin, J.; Rissmann, A.; Tucker, D.; Wellesley, D.; Dolk, H. Trends inthe prevalence, risk and pregnancy outcome of multiple births withcongenital anomaly: a registry-based study in 14 European countries1984−2007. BJOG 2013, 120, 707−716.(7) Bolong, N.; Ismail, A. F.; Salim, M. R.; Matsuura, T. A review ofthe effects of emerging contaminants in wastewater and options fortheir removal. Desalination 2009, 239, 229−246.(8) von Sonntag, C. Advanced oxidation processes: mechanisticaspects. Water Sci. Technol. 2008, 58, 1015−1021.(9) Duesterberg, C. K.; Waite, T. D. Process Optimization of FentonOxidation Using Kinetic Modeling. Environ. Sci. Technol. 2006, 40,4189−4195.(10) Duesterberg, C. K.; Mylon, S. E.; Waite, T. D. pH Effects onIron-Catalyzed Oxidation using Fenton’s Reagent. Environ. Sci.Technol. 2008, 42, 8522−8527.(11) Bataineh, H.; Pestovsky, O.; Bakac, A. pH-induced mechanisticchangeover from hydroxyl radicals to iron(IV) in the Fenton reaction.Chem. Sci. 2012, 3, 1594−1599.(12) Miller, C. J.; Wadley, S.; Waite, T. D. Advanced OxidationProcesses for Water Treatment: Fundamentals and Applications; IWAPublishing, 2017.

(13) Huang, W.; Brigante, M.; Wu, F.; Mousty, C.; Hanna, K.;Mailhot, G. Assessment of the Fe(III)-EDDS Complex in Fenton-LikeProcesses: From the Radical Formation to the Degradation ofBisphenol A. Environ. Sci. Technol. 2013, 47, 1952−1959.(14) Miller, C. J.; Rose, A. L.; Waite, T. D. Importance of IronComplexation for Fenton-Mediated Hydroxyl Radical Production atCircumneutral pH. Front. Mar. Sci. 2016, 3, Article 134.(15) Chahbane, N.; Lenoir, D.; Souabi, S.; Collins, T. J.; Schramm,K.-W. FeIII-TAML-Catalyzed Green Oxidative Decolorization ofTextile Dyes in Wastewater. Clean: Soil, Air, Water 2007, 35, 459−464.(16) Shen, L. Q.; Beach, E. S.; Xiang, Y.; Tshudy, D. J.; Khanina, N.;Horwitz, C. P.; Bier, M. E.; Collins, T. J. Rapid, BiomimeticDegradation in Water of the Persistent Drug Sertraline by TAMLCatalysts and Hydrogen Peroxide. Environ. Sci. Technol. 2011, 45,7882−7887.(17) Chanda, A.; Khetan, S. K.; Banerjee, D.; Ghosh, A.; Collins, T. J.Total Degradation of Fenitrothion and Other OrganophosphorusPesticides by Catalytic Oxidation Employing Fe-TAML PeroxideActivators. J. Am. Chem. Soc. 2006, 128, 12058−12059.(18) Kundu, S.; Chanda, A.; Thompson, J. V. K.; Diabes, G.; Khetan,S. K.; Ryabov, A. D.; Collins, T. J. Rapid degradation of oxidationresistant nitrophenols by TAML activator and H2O2. Catal. Sci.Technol. 2015, 5, 1775−1782.(19) Tang, L. L.; DeNardo, M. A.; Schuler, C. J.; Mills, M. R.;Gayathri, C.; Gil, R. R.; Kanda, R.; Collins, T. J. HomogeneousCatalysis Under Ultradilute Conditions: TAML/NaClO Oxidation ofPersistent Metaldehyde. J. Am. Chem. Soc. 2017, 139, 879−887.(20) Tang, L. L.; DeNardo, M. A.; Gayathri, C.; Gil, R. R.; Kanda, R.;Collins, T. J. TAML/H2O2 Oxidative Degradation of Metaldehyde:Pursuing Better Water Treatment for the Most Persistent Pollutants.Environ. Sci. Technol. 2016, 50, 5261−5268.(21) Demeter, E. L.; Hilburg, S. L.; Washburn, N. R.; Collins, T. J.;Kitchin, J. R. Electrocatalytic Oxygen Evolution with an ImmobilizedTAML Activator. J. Am. Chem. Soc. 2014, 136, 5603−5606.(22) Rohde, J.-U.; Torelli, S.; Shan, X.; Lim, M. H.; Klinker, E. J.;Kaizer, J.; Chen, K.; Nam, W.; Que, L., Jr Structural Insights intoNonheme Alkylperoxoiron(III) and Oxoiron(IV) Intermediates by X-ray Absorption Spectroscopy. J. Am. Chem. Soc. 2004, 126, 16750−16761.(23) Lubben, M.; Meetsma, A.; Wilkinson, E. C.; Feringa, B.; Que, L.,Jr Nonheme Iron Centers in Oxygen Activation: Characterization ofan Iron(III) Hydroperoxide Intermediate. Angew. Chem., Int. Ed. Engl.1995, 34, 1512−1514.(24) Codola, Z.; Gomez, L.; Kleespies, S. T.; Que, L., Jr; Costas, M.;Lloret-Fillol, J. Evidence for an oxygen evolving iron−oxo−ceriumintermediate in iron-catalysed water oxidation. Nat. Commun. 2015, 6,6405.(25) Wang, D.; Ray, K.; Collins, M. J.; Farquhar, E. R.; Frisch, J. R.;Gomez, L.; Jackson, T. A.; Kerscher, M.; Waleska, A.; Comba, P.;Costas, M.; Que, L. Nonheme oxoiron(IV) complexes of pentadentateN5 ligands: spectroscopy, electrochemistry, and oxidative reactivity.Chem. Sci. 2013, 4, 282−291.(26) Vad, M. S.; Lennartson, A.; Nielsen, A.; Harmer, J.; McGrady, J.E.; Frandsen, C.; Morup, S.; McKenzie, C. J. An aqueous non-hemeFe(IV)oxo complex with a basic group in the second coordinationsphere. Chem. Commun. 2012, 48, 10880−10882.(27) Fillol, J. L.; Codola, Z.; Garcia-Bosch, I.; Gomez, L.; Pla, J. J.;Costas, M. Efficient water oxidation catalysts based on readily availableiron coordination complexes. Nat. Chem. 2011, 3, 807−813.(28) de Sousa, D. P.; Bigelow, J. O.; Sundberg, J.; Que, L., Jr;McKenzie, C. J. Caught! Crystal trapping of a side-on peroxo bound toCr(IV). Chem. Commun. 2015, 51, 2802−2805.(29) Vad, M. S.; Nielsen, A.; Lennartson, A.; Bond, A. D.; McGrady,J. E.; McKenzie, C. J. Switching on oxygen activation by cobaltcomplexes of pentadentate ligands. Dalton Trans. 2011, 40, 10698−10707.(30) de Sousa, D. P.; Wegeberg, C.; Vad, M. S.; Mørup, S.; Frandsen,C.; Donald, W. A.; McKenzie, C. J. Halogen-Bonding-Assisted

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14946

Iodosylbenzene Activation by a Homogenous Iron Catalyst. Chem. -Eur. J. 2016, 22, 3810−3820.(31) Lennartson, A.; McKenzie, C. J. An Iron(III) IodosylbenzeneComplex: A Masked Non-Heme FeVO. Angew. Chem., Int. Ed. 2012,51, 6767−6770.(32) Deville, C.; Finsel, M.; de Sousa, D. P.; Szafranowska, B.;Behnken, J.; Svane, S.; Bond, A. D.; Seidler-Egdal, R. K.; McKenzie, C.J. Too Many Cooks Spoil the Broth - Variable Potencies of OxidizingMn Complexes of a Hexadentate Carboxylato Ligand. Eur. J. Inorg.Chem. 2015, 18, 3543−3549.(33) de Sousa, D. P.; McKenzie, C. J. In Iron Catalysis II; Bauer, E.,Ed.; Springer International Publishing, 2015.(34) McDonald, A. R.; Que, L., Jr High-valent nonheme iron-oxocomplexes: Synthesis, structure, and spectroscopy. Coord. Chem. Rev.2013, 257, 414−428.(35) Hill, E. A.; Weitz, A. C.; Onderko, E.; Romero-Rivera, A.; Guo,Y.; Swart, M.; Bominaar, E. L.; Green, M. T.; Hendrich, M. P.; Lacy, D.C.; Borovik, A. S. Reactivity of an FeIV-Oxo Complex with Protons andOxidants. J. Am. Chem. Soc. 2016, 138, 13143−13146.(36) Luo, Y.-R. Comprehensive handbook of chemical bond energies;CRC Press, 2007.(37) Blanksby, S. J.; Ellison, G. B. Bond dissociation energies oforganic molecules. Acc. Chem. Res. 2003, 36, 255−263.(38) Ruscic, B.; Feller, D.; Dixon, D. A.; Peterson, K. A.; Harding, L.B.; Asher, R. L.; Wagner, A. F. Evidence for a lower enthalpy offormation of hydroxyl radical and a lower gas-phase bond dissociationenergy of water. J. Phys. Chem. A 2001, 105, 1−4.(39) Buxton, G. V. Pulse radiolysis of aqueous solutions. Rate ofreaction of OH with OH−. Trans. Faraday Soc. 1970, 66, 1656−1660.(40) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B.Critical Review of rate constants for reactions of hydrated electrons,hydrogen atoms and hydroxyl radicals (·OH/·O−) in AqueousSolution. J. Phys. Chem. Ref. Data 1988, 17, 513−886.(41) Land, E. J.; Ebert, M. Pulse radiolysis studies of aqueous phenol.Water elimination from dihydroxycyclohexadienyl radicals to formphenoxyl. Trans. Faraday Soc. 1967, 63, 1181−1190.(42) Stafford, U.; Gray, K. A.; Kamat, P. V. Radiolytic and TiO2-assisted photocatalytic degradation of 4-chlorophenol. A comparativestudy. J. Phys. Chem. 1994, 98, 6343−6351.(43) Prutz, W.; Sommermeyer, K. Role of the hydrated electron inproducing chemiluminescence in aqueous dye solutions by roentgenrays and electron impulses. Radiat. Environ. Biophys. 1967, 4, 48−62.(44) Kucherenko, E. A. Pulse-Radiolysis of Rhodamine DyeSolutions. 3. Aqueous and Ethanolic Solutions of Rhodamine-B.High Energy Chem. 1982, 168−173.(45) Eigen, M.; Wilkins, R. G. In Mechanisms of Inorganic Reactions;Advances in Chemistry Series 49; Marmann, R. K., Fraser, R. T.,Bouman, Eds.; American Chemical Society, 1965.(46) Rana, S.; Dey, A.; Maiti, D. Mechanistic elucidation of C-Hoxidation by electron rich non-heme iron(iv)-oxo at room temper-ature. Chem. Commun. 2015, 51, 14469−14472.(47) Planas, O.; Clemancey, M.; Latour, J.-M.; Company, A.; Costas,M. Structural modeling of iron halogenases: synthesis and reactivity ofhalide-iron(IV)-oxo compounds. Chem. Commun. 2014, 50, 10887−10890.(48) Hong, S.; So, H.; Yoon, H.; Cho, K.-B.; Lee, Y.-M.; Fukuzumi,S.; Nam, W. Reactivity comparison of high-valent iron(IV)-oxocomplexes bearing N-tetramethylated cyclam ligands with differentring size. Dalton Trans. 2013, 42, 7842−7845.(49) Sastri, C. V.; Lee, J.; Oh, K.; Lee, Y. J.; Lee, J.; Jackson, T. A.;Ray, K.; Hirao, H.; Shin, W.; Halfen, J. A.; Kim, J.; Que, L., Jr; Shaik,S.; Nam, W. Axial ligand tuning of a nonheme iron (IV)-oxo unit forhydrogen atom abstraction. Proc. Natl. Acad. Sci. U. S. A. 2007, 104,19181−19186.(50) Siewert, I. Proton-Coupled Electron Transfer ReactionsCatalysed by 3 d Metal Complexes. Chem. - Eur. J. 2015, 21,15078−15091.

(51) Warren, J. J.; Tronic, T. A.; Mayer, J. M. Thermochemistry ofProton-Coupled Electron Transfer Reagents and its Implications.Chem. Rev. 2010, 110, 6961−7001.(52) Mayer, J. M. Understanding Hydrogen Atom Transfer: FromBond Strengths to Marcus Theory. Acc. Chem. Res. 2011, 44, 36−46.(53) Warren, J. J.; Mayer, J. M. In RSC Catalysis Series; Formosinho,S., Barroso, M., Eds.; Royal Society of Chemistry: Cambridge, U.K.,2011.(54) Mader, E. A.; Davidson, E. R.; Mayer, J. M. Large Ground-StateEntropy Changes for Hydrogen Atom Transfer Reactions of IronComplexes. J. Am. Chem. Soc. 2007, 129, 5153−5166.(55) Mitroka, S.; Zimmeck, S.; Troya, D.; Tanko, J. M. How SolventModulates Hydroxyl Radical Reactivity in Hydrogen Atom Abstrac-tions. J. Am. Chem. Soc. 2010, 132, 2907−2913.(56) Anglada, J. M.; Gonzalez, J. Different catalytic effects of a singlewater molecule: The gas-phase reaction of formic acid with hydroxylradical in water vapor. ChemPhysChem 2009, 10, 3034−3045.(57) Wang, D.; Zhang, M.; Buhlmann, P.; Que, L., Jr Redox Potentialand C-H Bond Cleaving Properties of a Nonheme Fe(IV)OComplex in Aqueous Solution. J. Am. Chem. Soc. 2010, 132, 7638−7644.(58) Usui, Y.; Sato, K. A green method of adipic acid synthesis:organic solvent- and halide-free oxidation of cycloalkanones with 30%hydrogen peroxide. Green Chem. 2003, 5, 373−375.(59) Hasanzadeh, M.; Karim-Nezhad, G.; Mahjani, M. G.; Jafarian,M.; Shadjou, N.; Khalilzadeh, B.; Saghatforoush, L. A. A study of theelectrocatalytic oxidation of cyclohexanol on copper electrode. Catal.Commun. 2008, 10, 295−299.(60) Pal, N.; Bhaumik, A. Mesoporous materials: versatile supports inheterogeneous catalysis for liquid phase catalytic transformations. RSCAdv. 2015, 5, 24363−24391.(61) Fang, T.; Fu, L.-Z.; Zhou, L.-L.; Zhan, S.-Z. A water-solubledinuclear copper electrocatalyst, [Cu(oxpn)Cu(OH)2] for both waterreduction and oxidation. Electrochim. Acta 2015, 161, 388−394.(62) Liu, T.; DuBois, D. L.; Bullock, R. M. An iron complex withpendent amines as a molecular electrocatalyst for oxidation ofhydrogen. Nat. Chem. 2013, 5, 228−233.(63) McCrory, C. C. L.; Uyeda, C.; Peters, J. C. ElectrocatalyticHydrogen Evolution in Acidic Water with Molecular CobaltTetraazamacrocycles. J. Am. Chem. Soc. 2012, 134, 3164−3170.(64) Nicholson, R. S.; Shain, I. Theory of stationary electrodepolarography. Single scan and cyclic methods applied to reversible,irreversible, and kinetic systems. Anal. Chem. 1964, 36, 706−723.(65) Schug, T. T.; Abagyan, R.; Blumberg, B.; Collins, T. J.; Crews,D.; DeFur, P. L.; Dickerson, S. M.; Edwards, T. M.; Gore, A. C.;Guillette, L. J.; Hayes, T.; Heindel, J. J.; Moores, A.; Patisaul, H. B.;Tal, T. L.; Thayer, K. A.; Vandenberg, L. N.; Warner, J. C.; Watson, C.S.; vom Saal, F. S.; Zoeller, R. T.; O’Brien, K. P.; Myers, J. P. Designingendocrine disruption out of the next generation of chemicals. GreenChem. 2013, 15, 181−198.(66) Duelund, L.; Hazell, R.; McKenzie, C. J.; Preuss Nielsen, L.;Toftlund, H. Solid and solution state structures of mono- and di-nuclear iron(III) complexes of related hexadentate and pentadentateaminopyridyl ligands. J. Chem. Soc., Dalton Trans. 2001, 152−156.(67) Hanania, G. I. H.; Irvine, D. H.; Eaton, W. A.; George, P.Thermodynamic aspects of the potassium hexacyano-ferrate(III)-(II)system. II. Reduction potential. J. Phys. Chem. 1967, 71, 2022−2030.(68) Miller, C. J.; Yu, H.; Waite, T. D. Degradation of rhodamine Bduring visible light photocatalysis employing Ag@AgCl embedded onreduced graphene oxide. Colloids Surf., A 2013, 435, 147−153.(69) Watanabe, T.; Takizawa, T.; Honda, K. Photocatalysis throughexcitation of adsorbates. 1. Highly efficient N-deethylation ofrhodamine B adsorbed to cadmium sulfide. J. Phys. Chem. 1977, 81,1845−1851.

Inorganic Chemistry Article

DOI: 10.1021/acs.inorgchem.7b02208Inorg. Chem. 2017, 56, 14936−14947

14947