Cu complexes that catalyze the oxygen reduction reaction

Coordination Chemistry Reviews 257 (2013) 130–139

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Coordination Chemistry Reviews journal homepage: www.elsevier.com/locate/ccr

Review

Cu complexes that catalyze the oxygen reduction reaction Matthew A. Thorseth, Claire E. Tornow, Edmund C.M. Tse, Andrew A. Gewirth ∗ Department of Chemistry, University of Illinois at Urbana-Champaign, 600 S. Matthews Avenue, Urbana, IL 61801, United States

Contents 1. 2. 3. 4. 5. 6. 7. 8.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Copper complexes with porphyrins and phthalocyanines . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Copper complexes with amino-alkyl ligands . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Copper complexes with substituted 1,10-phenanthrolines . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Copper complexes with other aromatic N-donor ligands . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Copper complexes with tris(2-pyridylmethyl)amine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Copper complexes with substituted triazoles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

a r t i c l e

i n f o

Article history: Received 26 December 2011 Received in revised form 2 March 2012 Accepted 30 March 2012 Available online 5 April 2012

130 131 132 133 133 135 135 136 138

a b s t r a c t The oxygen reduction reaction (ORR) is employed in a large number of systems such as fuel cells and air batteries. Currently, the catalyst with the lowest overpotential for the ORR is the enzyme laccase. Laccase only functions at a very narrow pH range, and its large size prevents high current densities. Using copper based catalysts to mimic the ORR activity is an area with many spectroscopic results, but relatively few electrochemical studies. This review catalogs the various copper based ORR catalysts and their activities.

Keywords: Oxygen reduction reaction (ORR) Cu catalysts Electrochemistry

© 2012 Elsevier B.V. All rights reserved.

1. Introduction

Abbreviations: ads, adsorbed; AFC, alkaline fuel cell; apy(4-imidazolyl)ethylene-2-amino-1-ethylpyridine; baEtO, hist, 2-(bis(2-aminoethyl)amino)ethanol; bistripic, 1,2-bis(6-(bis(6-methylpyridin2-yl)methyl)pyridin-2-yl)ethane); C–E, chemical–electrochemical; 5-Cl-phen, 5-chloro-1,10-phenanthroline; DMP, 2,9-dimethyl-1,10-phenanthroline; DPP, 2,4bis(2-pyridyl)pyrimidine; 4,7-dppds, 4,7-diphenyl-l,l0-phenanthrolinedisulfonate; ESR, electron spin resonance; GC, glassy carbon; Hdatrz, 3,5-diamino-1,2,4-triazole; Im, imidazole; Me3 TACN, 1,4,7-trimethyl-1,4,7-triazacyclononane; Me6 tren, N1 ,N1 bis(2-(dimethylamino)ethyl)-N2 ,N2 -dimethylethane-1,2-diamine; NHE, normal hydrogen electrode; NPM, non-platinum metal; ORR, oxygen reduction reaction; PBS, phosphate buffered saline; PDT, 3-(2-pyridyl)-5,6-diphenyl-1,2,4-triazine; phen, 1,10-phenanthroline; Piv, pivalamido; PMAP, (2-(pyridine-2-yl)PMEA, N-(2-(pyridine-2-yl)ethyl)-N-(pyridine-2-ylmethyl)ethanamine); poly-his, poly(2-(pyridine-2-yl)-N,N-bis(pyridine-2-ylmethyl)ethanamine); l-histidine; POM, polyoxometalate; RHE, reversible hydrogen electrode; RRDE, rotating ring-disk electrode; SCE, saturated calomel electrode; TEPA, TPA, tris(2-pyridylmethyl)amine; TPT, tris(2-(pyridine-2-yl)ethyl)amine; 2,4,6-tris(2-pyridyl)-1,3,5-triazine; tren, tris(2-aminoethyl)amine; trpn, tris(3aminopropyl)amine. ∗ Corresponding author. Tel.: +1 217 333 8329; fax: +1 217 244 3186. E-mail address: [email protected] (A.A. Gewirth). 0010-8545/$ – see front matter © 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.ccr.2012.03.033

The oxygen reduction reaction (ORR) is a four-electron reduction of dioxygen to water (Eq. (1)). The standard reduction potential for the reaction is 1.23 V, meaning that dioxygen is a powerful oxidizing agent [1]. The triplet ground state of dioxygen often prevents direct reaction with many molecules, which typically have a singlet ground state, and thus prevents rapid reactivity of dioxygen despite the high oxidation potential. In acid, the ORR can proceed by a direct four-electron and four-proton reaction with O2 to yield H2 O, or can proceed by consecutive two-electron and two-proton steps to yield H2 O2 and then H2 O. In base, the four-electron direct reduction results in the formation of four equivalents of hydroxide, while the two-electron reduced product is one equivalent of hydroxide and one hydroperoxyl anion. The hydroperoxyl anion can then be further reduced by two electrons to three hydroxide ions. Acid form O2 + 4e− + 4H+  2H2 O

E 0 = 1.229 V

(1)

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O2 + 2e− + 2H+  H2 O2

E 0 = 0.695 V

H2 O2 + 2e− + 2H+  2H2 O

131

(2)

E 0 = 1.763 V

(3)

E 0 = 0.401 V

(4)

Base form O2 + 2H2 O + 4e−  4OH−

O2 + H2 O + 2e−  HO2 − + OH− HO2 − + H2 O + 2e−  3OH−

E 0 = −0.076 V E 0 = 0.878 V

(5) (6)

The ORR is important in systems ranging from fuel cells, to airbatteries, to corrosion. In fuel cells, the ORR forms one of the half reactions responsible for extracting energy from a fuel; the other is the oxidation half reaction. Many fuels are utilized for fuel cells, but the most common one for low temperature (<100 ◦ C) is hydrogen [2–4]. Fuel cells are interesting because they have the potential to extract most of the energy available in a fuel, without the Carnot cycle limitations that attend combustion as would occur, for example, in an internal combustion engine or a gas turbine. Thus, the thermodynamic efficiency of a fuel cell could be quite high. Although fuel cells were invented some 170 years ago, and modern implementations have been in use since the 1960s, there is still a variety of factors which inhibits widespread use of low temperature fuel cells. One of the most compelling challenges is the slow kinetics and substantial overpotentials which attend the ORR. The consequence of this overpotential means that fuel cells are unable to deliver the full 1.23 V that a H2 /O2 fuel cell should deliver. Rather, the fuel cell potential is only ca. 40–60% of the thermodynamic value when practical current densities are reached. This drop in potential means that the full thermodynamic efficiency of the system is not achieved, which obviates the major advantage of the fuel cell in the first place. The origin of the overpotential for the ORR is the high bond strength of the dioxygen double bond (498 kJ) [5]. Pt or its alloys are the most commonly used cathode catalysts, and these exhibit an overpotential of ∼300 mV for the ORR [5–7]. Unfortunately, the large amount of Pt found in typical implementations makes the fuel cell cost prohibitive. There are also issues with catalyst stability and susceptibility toward poisoning. Catalyst cost is prompting a search for non-precious metal containing materials that might function as ORR catalysts. Many different non-Pt containing materials have been examined for their ORR activity [8–12]. One area of renewed focus is copper-containing proteins, including laccase. Many different methods have been utilized to connect these proteins to electrode surfaces for interrogation, including adsorption, covalent attachment, and by using mediators, both attached to the electrode surface, or otherwise available [5,13,14]. Laccase has an almost nonexistent overpotential (20 mV), a turn-over rate of 2.1O2 per laccase per second, as well as reaching diffusion limited behavior by 70 mV of overpotential (Fig. 1) [15,16]. Unfortunately, laccase only functions in a narrow pH range, and its size makes it difficult to achieve the high packing densities that would lead to high current densities with good mass transport. Additionally, the stability of these systems can be limited, particularly in the aggressive environment of a fuel cell [17–19]. One way to achieve the activity of a copper enzyme might be through the agency of inorganic complexes. While many copper compounds exhibit reactivity with dioxygen and have been studied spectroscopically, there have been relatively few attempts to understand their reactivity for the ORR electrochemically (Fig. 2). This review summarizes key results from copper ORR electrocatalysts and develops future directions for research.

Fig. 1. Polarization diagram of dioxygen electroreduced at a carbon fiber electrode modified with laccase wired with polymers (curves 1 and 2) vs. a 6-␮m diameter platinum electrode (curve 3) in pH 5 citrate buffer and Pt fiber electrode in 0.5 M H2 SO4 in air. Reprinted with permission from the American Chemical Society [15].

2. Copper complexes with porphyrins and phthalocyanines Jasinski’s early discovery, that cobalt phthalocyanine could catalyze the ORR, has laid the foundation for subsequent research on non-platinum metal (NPM)-based ORR catalysts [20]. Shortly following this discovery, Collman and Anson reported a more active dicobalt porphyrin dimer system [21–23]. Although a great amount of work has since contributed to the initial successes of cobalt complexes, limited success using copper as the metal center has been achieved using both the phthalocyanine and porphyrin sub-structures. Studies using metal phthalocyanines for O2 reduction report an activity trend following the order of: Fe(II) > Co(II) > Ni(II) > Cu(II) [24,25]. These differences in electrocatalytic activity have been clarified using MO theory. Bonding of O2 to the metal orbitals depends largely on vacancies in the d orbitals of the metal center [26]. Fe(II), Co(II), Ni(II), and Cu(II) each have filled dxz and dyz orbitals; however, Fe(II) and Co(II) possess either an empty or half-filled dz 2 orbital, which can accept electrons from O2 . Following these observed trends, a study by Itaya et al. demonstrated that while Co phthalocyanine adsorbed on a Au(1 1 1) surface shows significant activity toward the ORR, a Cu phthalocyanine-modified electrode exhibits decreased O2 reduction current [27]. More recently, Bekaro˘glu et al. reported the O2 reduction activity of ball-type supramolecular metallophthalocyanines using both Co and Cu [28]. In acidic solution, O2 reduction with the Co complex occurs at a significantly lower overpotential than the Cu complex, a feature that was again attributed to the enhanced binding ability of O2 by the Co structure. The copper–porphyrin systems also exhibit limited activity toward O2 reduction as well as poor catalyst stability. Qui et al. reported an insoluble copper(II) triphenylporphyrin complex deposited on a glassy carbon (GC) electrode that exhibits no catalytic activity for the reduction of O2 in acidic conditions [29]. By modifying the porphyrin with a hexadecyl-pyridiniumyl group, a modest onset of 0.13 V vs. RHE was observed, yet the electrocatalyst appears to be stable for only a few scans. Similarly, Takehira and coworkers discovered that a water-soluble complex of copper and a sulfonated tetraphenylporphyrin does not reduce O2 , unlike the analogous Co, Fe, and Mn porphyrin complexes [30]. Zhuang et al., however, demonstrated an onset of ORR at 0.44 V

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Fig. 2. Diagram of some of the reported Cu ORR catalysts.

vs. RHE with a copper(II) tetraphenylporphyrin complex at pH 4.0 using a paraffin-impregnated graphite electrode [31]. Swavey et al. attempted to coat edge-plane pyrolytic graphite with a copper(II) tetraphenylporphyrin, however, due to limited stability of the copper(II) porphyrin in acidic conditions, it was not possible to evaluate the electrocatalytic reduction of O2 [32]. 3. Copper complexes with amino-alkyl ligands A copper complex with tris(3-aminopropyl)amine (trpn) and imidazole (Im) was examined by Cai et al. [33]. [Cu(trpn)(Im)](ClO4 )2 was adsorbed onto a pyrolytic graphite electrode by soaking the electrode in a solution containing the

Cu complex. The complex exhibited a reversible CuI/II couple at −0.23 V vs. SCE in pH 6.4 Britton–Robinson buffer with 0.1 mM imidazole. Removal of the imidazole resulted in a decrease in the CuI/II couple current upon subsequent scans, but this current could be revived with the addition of imidazole to the electrolyte. Upon addition of O2 to the solution, increasing cathodic current was observed with an onset of 0.58 V vs. RHE with n = 4e− by Koutecky–Levich analysis. Reduction of H2 O2 occurs at the same potential with an n = 2e− , and at similar rates to O2 reduction. A similar complex, [Cu(baEtO)(Im)](ClO4 )2 [baEtO = 2-(bis(2aminoethyl)amino)ethanol], has an E1/2 (CuI/II ) = −0.248 V vs. SCE in pH 7 phosphate buffer also adsorbed to pyrolytic graphite [34]. The CuI/II couple is observed even in the absence of imidazole in

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the electrolyte and appears to be stable for 48 h of potential cycling. The onset of O2 reduction occurs at 0.57 V vs. RHE with an n = 4e− as determined from the Cottrell equation, while H2 O2 reduction was not studied. The complex exhibits two linear regions in a plot of peak current for O2 reduction vs. pH: the region of pH 4–6 where the potential is independent of the pH, and pH 6–10, where the peak current varies by −56 mV/pH. Our group tested a few copper(II) complexes with tris(2-aminoethyl)amine (tren), N1 ,N1 -bis(2(dimethylamino)ethyl)-N2 ,N2 -dimethylethane-1,2-diamine and 1,4,7-trimethyl-1,4,7-triazacyclononane (Me6 tren), (Me3 TACN) [35]. [Cu(tren)(H2 O)](ClO4 )2 and [Cu(Me6 tren)(H2 O)](ClO4 )2 showed similar voltammetry to [Cu(trpn)(Im)](ClO4 )2 under Ar; no CuI/II couple was observed. Onsets of O2 reduction were observed for the [Cu(tren)(H2 O)](ClO4 )2 at 0.43 V vs. RHE and [Cu(Me6 tren)(H2 O)](ClO4 )2 at 0.33 V vs. RHE, but both only exhibited the 2e− reduction of O2 to H2 O2 . [Cu(Me3 TACN)(H2 O)](ClO4 )2 does have a CuI/II couple at 0.18 V vs. RHE and an onset of ORR at 0.30 V vs. RHE, much more negative than other reported complexes. We attributed the poor activity of the alkyl complexes to slow electron transfer to the Cu center. The overlap between the ␴* orbitals on the alkyl ligand and the ␲* orbitals of the graphitic carbon is small, leading to poor electron conduction from the support to the catalyst.

4. Copper complexes with substituted 1,10-phenanthrolines A series of copper(II) complexes with 1,10-phenanthroline (phen), 5-chloro-1,10-phenanthroline (5-Cl-phen), 4,7diphenyl-l,l0-phenanthrolinedisulfonate (4,7-dppds), 2,9-dimethyl-1,10-phenanthroline (DMP), which may form a mixture of the [CuL]2+ and [CuL2 ]2+ complexes, adsorbed onto edge plane graphite electrodes were prepared by Anson and coworkers [36–40]. Attaching chloro groups to the 5 position resulted in a positive shift in E1/2 (CuI/II ) by 75 mV compared to the unsubstituted phenanthroline. The CuI/II couple of the Cu complex with DMP shifts further positive by 300 mV vs. copper phen. The onset potentials for ORR for the phen, 5-Cl-phen, and 4,7-dppds complexes are similar at ∼0.5 V vs. RHE. The addition of sterically hindering methyl groups resulted in a much larger positive shift in the onset of ORR to 0.69 V vs. RHE. The shift in ORR onset potential is attributed to a simple outer sphere electron transfer to O2 being the rate-determining step. All the catalysts examined also showed activity for H2 O2 reduction, although at much slower rates than for O2 reduction. In 2007, a more comprehensive library of copper complexes with substituted 1,10-phenanthrolines was reported by Chidsey, Stack, and coworkers [41]. The potential of the CuI/II couple can be shifted more positive by 125 mV, as compared to phen, by substitution with carboethoxy and nitro electron withdrawing groups. A more significant increase (∼300 mV) of the CuI/II couple was observed by substitution of sterically hindering methyl and ethyl groups at the 2,9 positions (Fig. 3). Increasing the CuI/II couple increased the peak potential for ORR, as seen in Anson’s work, up to 0.59 V vs. RHE for Cu complex with 2,9-diethyl-phen. The proposed mechanism of the ORR, as proposed by Anson, was modified. The CuII center is first reduced to CuI , followed shortly by the binding of oxygen. The Cu–O2 complex is then reduced and protonated in subsequent steps, an inner-sphere mechanism. The authors assumed a mono-Cu reaction mechanism. Chidsey later covalently attached 3-ethynyl-phenanthroline to azide-modified GC surfaces via the click reaction [42]. By varying the copper coverage on the electrode, a second-order dependence

Fig. 3. Cyclic voltammograms of Cu complexes with substituted 1,10phenanthrolines adsorbed onto edge plane graphite in N2 (dotted) and air (solid) saturated solutions at pH 4.8. Reprinted with permission copyright American Chemical Society [41].

on the quantity of immobilized copper complexes was observed and the electroreduction of O2 proceeded via a binuclear [Cu2 O2 ] intermediate. In Anson’s case, since the copper complexes were physisorbed onto the carbon electrodes, after binding one equivalent of O2 with [CuI (phen)]ads , a second [CuI (phen)]ads could rapidly migrate to form the [Cu2 O2 ] adduct, which resulted in the observed first order dependence on Cu. 5. Copper complexes with other aromatic N-donor ligands [LCu2 (CH3 CO2 )2 ](ClO4 )2 ·5H2 O [L = 3,7,11,18,22,26hexaazatricyclo-[26.2.2.213,16 ]tetratriaconta1(31),13(33),14,16(34),28(32),29-hexaene)] was prepared and tested for its ORR activity [43,44]. The di-Cu hexaaza-xylyl macrocyclic complex exhibits a reversible 2e− reduction wave at −0.1 V vs. SCE under Ar. ESR spectroscopy does not exhibit exchange

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Fig. 4. (a) Current–potential curves for reduction of O2 electrocatalyzed by a Cu complex with TPT adsorbed onto graphite electrode, recorded at various electrode rotating rates. (b) Koutecky–Levich plots for reduction of O2 . O2 concentration: 2.4 × 10−4 mol/dm3 . Britton–Robinson buffer: pH 5.3. Scan rate: 100 mV/s. TPT surface concentration: 9.97 × 10−10 mol/cm2 . Reprinted with permission copyright Elsevier [46].

coupling between the two Cu atoms, likely due to the large spacing ˚ between them. Introduction of O2 increases the cathodic (8.40 A) current at potentials similar to the CuI/II couple, with an onset of ORR at −0.1 V vs. RHE. At low loading of the Cu complex, increased H2 O2 production was detected. Increasing the loading eliminated the peroxide oxidation peak due to further reduction of H2 O2 by the excess Cu complex. The mechanism of O2 reduction was proposed to be a 2e− reduction of O2 to H2 O2 and then subsequent reduction of the H2 O2 to H2 O, at a slower rate. [Cu2 (apyhist)2 Cl2 ]2+ [apyhist = (4Treating imidazolyl)ethylene-2-amino-1-ethylpyridine] with pH 9 buffer resulted in the formation of a tetracopper complex from the deprotonation of the imidazole groups [45]. The Cu4 complex was cast directly onto a GC electrode, where poor electron transfer rates (1.8 × 10−11 cm2 /s) to the Cu centers were observed. Introduction of O2 resulted in a modest increase of the onset potential of 350 mV over the GC electrode. The onset was 0.53 V vs. RHE in pH 9 solutions and chronoamperometric experiments revealed a 4e− transfer process. Cu complexes with 2,4-bis(2-pyridyl)pyrimidine (DPP), 3(2-pyridyl)-5,6-diphenyl-1,2,4-triazine (PDT), and 2,4,6-tris(2pyridyl)-1,3,5-triazine (TPT) exhibit similar redox waves at potentials more negative than −0.5 V vs. SCE [39,46]. Addition of Cu to the ligand adsorbed onto pyrolytic graphite electrodes exhibit CuI/II couples at 0.0 V vs. SCE. The Cu complex with PDT does not exhibit any reactivity with O2 , which is attributed to the tetrahedral coordination of the CuI complex. Both Cu(I) complexes with DPP and TPT as supporting ligands do catalyze the ORR with onsets at 0.56 V and 0.51 V vs. RHE respectively with n = 4e− (Fig. 4). Both Cu(I) complexes also catalyze H2 O2 reduction to H2 O at similar potentials, again at slower overall rates than O2 reduction. While the ligands exhibit reversible couples, their potentials are too negative to play a large role in the ORR catalysis. Beyer et al. reported the use of polymeric film-modified Au electrodes of copper(II) Schiff-base complexes possessing pyrrole groups for use as O2 reduction catalysts [47]. Specifically, [CuII (LH−1 )2 (L )] (L = 2-(3-pyrrol-1-yl-propylimino-methyl)phenol, L = H2 O or a neutral 2e− donor) achieves electrochemical O2 reduction at a potential 400 mV more positive than the bare

electrode in basic solution. RRDE measurements reveal the production of a small amount of H2 O2 , demonstrating that a parallel 2e− pathway is followed, where O2 is reduced to H2 O2 rather than directly to H2 O. Studies in both basic and acidic media show that the potentials at which the reduction of O2 takes place are similar, suggesting that the rate-determining step is the pH independent formation of superoxide. Na14 [SiW9 O34 Cu3 (N3 )2 (OH)(H2 O)]2 ·24H2 O, the first multidimensional, antiferromagnetic, hexacopper complex based on azido polyoxometalate (POM) units, was synthesized by Nadjo and coworkers [48]. Cyclic voltammetry exhibits two Cu reduction peaks at 0.38 V and 0.25 V vs. RHE, indicating two reduction processes, CuI/II , and Cu0/I . The multiple one-electron reduction steps were suggested to be the key to trigger the reduction of O2 . The compound reduces O2 at a high rate via an overall four-electron process; however, the onset for the ORR is fairly negative at 0.33 V vs. RHE. In effort to synthesize an enzyme biomimetic, based on the framework of laccase, Bard and coworkers reported the use of polyl-histidine as a matrix and ligand to complex CuII [49]. Using a Cu complex-coated thin film GC electrode, it was observed that O2 reduction started at ∼0.07 V vs. Ag/AgCl, a slightly more positive onset than that achieved using a bare electrode (Fig. 5). At −0.3 V vs. Ag/AgCl, the current for the CuII -poly-his-modified electrode was 2.3-fold larger as compared to the bare GC electrode, demonstrating electrocatalytic success with this proof of principle system. However, further performance enhancements are expected using other polypeptide mixtures in hopes of better mimicking laccase. Pap and coworkers synthesized [CuII (LH−4 )]−2 , (L = 8,17-dioxa-1,2,5,6,10,11,14,15-octaaza-tricyclo[13.3.1] eicosane-3,4,12,13-tetrone), which catalyzed the 4e− reduction of oxygen and the simultaneous 2e− oxidation of ascorbic acid in basic condition (pH 8) [50]. Due to the rigidness of the planar, tetradentate ligand, a square pyramidal intermediate is anticipated when an additional ligand is coordinated to the apical position. In contrast to the biological system, which usually involves the CuI/II redox couple, Cu(III) and Cu(II) are the active species in Pap’s system, with a E1/2 of 0.64 V vs. RHE. According to the observed rate expression, the Pap group proposed a peroxodicopper(III)

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Fig. 5. Cyclic voltammograms of CuII -poly-his-modified GC electrode in 0.2 M PBS (pH 7). The condition of modified films: 8 mM histidine residues, 2 mM Cu2+ . Scan rate = 50 mV/s. (1) Bare GC, no O2 ; (2) bare GC, with O2 ; (3) modified GC, no O2 ; (4) modified GC, with O2 .

Fig. 6. Graph of the onset of O2 (red squares) and H2 O2 (black triangles) for [CuTPA](ClO4 )2 in Britton–Robinson buffers at various pHs.

Reprinted with permission copyright American Chemical Society [49].

Reprinted with permission copyright American Chemical Society [51].

intermediate, which dissociated into two copper(III)-oxyl radical anions in the rate-determining step. Despite a high rate of O2 reduction, the onset for the ORR is 0.55 V vs. RHE. To date, this is one of the rare Cu(III) complexes that carries out ORR.

(2-(pyridine-2-yl)-N,N-bis(pyridine-2-ylmethyl)ethanamine) (2-(pyridine-2-yl)-N-(2-(pyridine-2-yl)ethyl)(PMEA), N-(pyridine-2-ylmethyl)ethanamine) (PMAP), and (tris(2-(pyridine-2-yl)ethyl)amine) (TEPA) [35]. The CuI/II couple potential increases as the lengths of the linkers between the amine and the pyridine groups are increased, however, the onset for the ORR is the same for all four complexes at 0.69 V vs. RHE at pH 7. All four complexes also exhibit similar Tafel slopes. This implies that the rate-determining step for TPA based complexes is not the CuI/II couple reduction. Addition of H-bonding groups has been shown in the past to greatly affect the stability and reactivity of O2 complexes. The first crystal structure of a Cu–OOH species was obtained with a bispivalamido complex that H-bonded to a H2 O2 reactant and was stable for over a month at a time, unlike previous reports [54]. Adding H-bonding groups to [CuTPA]2+ does not lower the overpotential, however. [Cu((NH2 )2 -TPA)](NO3 )2 has the same ORR onset potential of 0.53 V vs. RHE as TPA at pH 1, while [Cu(Piv2 -TPA)](NO3 )2 has a more negative onset of 0.40 V vs. RHE [35]. The lower onset for the bispivalamido complex is due to the steric repulsions of the tert-butyl groups, preventing a di-Cu–O2 interaction. The similarity of the voltammetry between [CuTPA]2+ and [Cu(Piv2 -TPA)]2+ leads to the conclusion that the mechanism is the same between the two complexes and that protonation is likely not a part of the rate-determining step, which is in agreement with the mechanism suggested by the Tafel slope.

6. Copper complexes with tris(2-pyridylmethyl)amine The mode of O2 binding to a Cu complex plays a large role in the catalyst’s activity. O2 binds to di-Cu centers in two fashions, an endon trans ␮-1,2 coordination, or a side-on ␮-␩2 :␩2 mode. The initial coordination is an oxidative addition that reduces O2 to either the peroxo or oxo state, depending on the complex. [CuTPA]2+ binds O2 in the end-on mode while [Cu2 bistripic(L )]2+ (L = a neutral 2e− donor) binds O2 side-on. In acidic solutions, [CuTPA]2+ has an overpotential for the ORR of 700 mV (pH 1), while [Cu2 bistripic(L )]2+ has an overpotential of 890 mV (pH 2) [51]. The two complexes also exhibited different mechanisms based on the Tafel slopes: [CuTPA]2+ exhibits a 2e− reduction step (slope = 70 mV/dec) while [Cu2 bistripic(L )]2+ likely goes through a chemical–electrochemical (C–E) step (slope = 167–680 mV/dec). The 2e− rate-determining step for [CuTPA]2+ is corroborated by the 30 mV/pH slope of the onset potential between pH 4 and 10 (Fig. 6). We speculated that the differences between the two complexes were due to the nucleophilicity of the end-on complex (the TPA system) which leads to easy protonation of the peroxo intermediate, allowing further reduction of a neutral hydroperoxo intermediate. The side-on peroxo species (the bistripic system) is electrophilic, and the protonation of the species is slow and leads to greater overpotentials. By changing the loading of [CuTPA]2+ , we determined that the ORR mechanism must be at least a di-Cu mechanism [51]. Decreasing the partial pressure of O2 over the solution results in a linear decrease of the limiting current density [35]. The current density is proportional to the rate of reaction, so the reaction is first order in O2 . The rate-determining reaction then must consist of two Cu and one O2 , as is seen in solution [52,53]. The electroreduction of O2 involves a number of electron and proton transfer events. As was seen with Cu complexes with substituted 1,10-phenanthrolines, the CuI/II couple usually plays a large role in the ORR overpotential. Complexes based off of [CuTPA]2+ do not exhibit a change in overpotential with a change in the CuI/II couple. This is seen by examining the Cu complexes with

7. Copper complexes with substituted triazoles Inspired by the activity of the multicopper oxidase active site of laccase, Thorum et al. have introduced one of the most efficient synthetic copper electrocatalysts for the ORR to date [55]. The carbon-supported, insoluble copper coordination complex of copper with 3,5-diamino-1,2,4-triazole (Hdatrz) has a reported onset of 0.73 V vs. RHE at pH 7, and the highest overall onset of 0.86 V vs. RHE at pH 13 (Fig. 7). The onset potential vs. pH has a slope of 30 mV/pH. It is also reported to be stable at pH 7 over a testing period of 24 h. The crystal structure of [Cu2 (Hdatrz)2 (␮OH2 )(H2 O)4 (SO4 )](SO4 )·3.5H2 O reveals a binuclear complex with two copper centers, bridged by neutral guanazole ligands and one water molecule, with Cu· · ·Cu spacing comparable to Cu spacing

136

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Fig. 7. Graph of the onset potentials for the ORR vs. pH for [Cu(Hdatrz)(H2 O)2 ]2+ in Britton–Robinson buffers. Reprinted with permission copyright Angewandte [55].

in laccase, approximately 3.5 A˚ [56]. Magnetic susceptibility measurements of the carbon-supported complex demonstrate that spin pairing between Cu centers occurs, illustrating the presence of on electrode synthetic multi-copper sites. Studies of the O2 reduction activity of [Cu(Hdatrz)(H2 O)2 ]2+ in the presence of several anions and poisons probed whether or not these multi-copper sites are the active site for the ORR [57]. Using poisons that have demonstrated coordination to Cu complexes, a significant decrease in O2 reduction activity (up to ∼200 mV decrease in the onset potential) is observed in the presence of sodium fluoride, potassium thiocyanate, and ethanethiol. A lack of poisoning by sodium azide suggests the existence of an active site containing a neutral copper(II) complex with Hdatrz and SO4 2− ligands. [Cu(Hdatrz)(H2 O)2 ]2+ is also a promising candidate for future application as a cathode catalyst in alkaline fuel cells (AFC). This is the first reported synthetic multi-copper complex for use in an AFC environment. Employing an alkaline microfluidic H2 /O2 fuel cell platform, Brushett et al. demonstrated that on a per metal basis, the copper complex with triazole ligand outperforms both Pt/C and Ag/C cathode catalysts [58]. However, the power density achieved using the copper triazole complex still falls below that achieved with Pt/C electrocatalysts. Based on the initial success of copper coordinated with an amine-substituted triazole, it was expected that further increases of ORR onset potentials may come as a result of tailoring the triazole ligand with different functionalities. Since the triazole motif is expected to be essential for the ligand to coordinate the active copper site, various substitutions at the triazole’s 3,4, and 5 positions have exposed possible correlations between structure and activity [59]. Further substitution of the 3,5-diamino-1,2,4-triazole ligand at the 4 position with an amino group produces a linear trinuclear unit with a sulfate group bridging the copper centers. This complex has an onset of oxygen reduction at 0.67 V vs. RHE and reaches a steady-state diffusion-limited current of approximately −5.0 mA/cm2 . In contrast, the copper(II) complex with 3,5-dimethyl-4-amino-1,2,4-triazole produces a tricopper triangular cluster framework, possessing a ␮-3-OH group bridging the trinuclear center. The onset of O2 reduction for this complex occurs at 0.58 V vs. RHE and reaches diffusion-limited current densities only slightly better than the Vulcan carbon support. Removal of all substitutions produces a copper(II) complex of 1,2,4-triazole, which results in a triangular Cu(II) cluster, with an OH− group bridging the tricopper center. In this case, O2 reduction commences at 0.70 V vs. RHE, but displays irreproducibility between anodic and cathodic

Fig. 8. Graph of the onset potentials for the ORR for the reported Cu catalysts normalized to the RHE as listed in Table 1 vs. the reported pH.

scans, specifically in the mixed kinetic diffusion-control region, reaching current densities of approximately −5.4 mA/cm2 . 8. Conclusions While laccase remains the single best ORR catalyst to date, replicating its activity with synthetic compounds remains elusive. To date, only a very limited number of Cu catalysts have been examined for their ORR activity when compared to the number of Cu compounds that exhibit O2 reactivity. To summarize results based on the catalysis of the ORR with Cu complexes, Table 1 and Fig. 8 show the onset potentials of the catalysts discussed in this publication. In Table 1, the reported onset potentials are converted to RHE, and the pH at which the measurement was made is noted. Of the compounds studied, a few trends can be established. Firstly, most of the compounds examined likely react in a di-Cu fashion, as elegantly shown by Chidsey’s click chemistry study. Most of the observed Cu complexes catalyze both O2 and H2 O2 reduction, but the peroxide reduction occurs at a much slower rate. The ORR consists of 4H+ transfers, and as a result, a large dependence on the pH is observed for many Cu complexes. The overall trend is a 30 mV/pH increase in the onset potential for ORR, seen in the copper Hdatrz and copper TPA systems. The 30 mV/pH vs. RHE is a −30 mV/pH dependence vs. NHE, which implies that the rate-determining step for most catalysts is a 2e− transfer for every proton transferred, half the −60 mV/pH predicted for a 1e− transfer by the Nernst equation. Anson also observed a similar −30 mV/pH trend in the CuI/II couple potential in copper complexes with substituted 1,10-phenanthrolines. Finding a catalyst with a smaller pH dependence on the ORR activity would be a large step toward developing a better ORR catalyst. The ligand choice plays a very large role in the efficacy of the ORR. Alkyl ligands seem to perform quite poorly, possibly due to poor electronic conductivity. Most of the porphyrin and phthalocyanine complexes also exhibit low onset potentials for the ORR. The cause of the poor reactivity may be due to the forced square planar geometry around the Cu center. Since open coordination sites are only available above and below the square planar plane, only the filled dz 2 orbital is available to interact with the oxygen rather than the half-occupied dx 2 –y 2 orbital. Thus, O2 is not interacting with the orbital which forms the ground state of the complex. Ligands with pyridine, pyrrole, imidazole, and triazole structures appear to have the best overall activity. Future studies of Cu based ORR catalysts should focus on further understanding the fundamental reaction mechanisms, since it has

M.A. Thorseth et al. / Coordination Chemistry Reviews 257 (2013) 130–139

137

Table 1 Listing of the Cu catalysts studied for their ORR activity. Catalyst 2+

[Cu(phthalocyanine)]

[Cu(heptadecafluorodecyl substituted ball-type metallophthalocyanine)]2+ (BTMPcs) [Cu(5-(4-pyridyl)-10,15,20triphenylporphyrin)]2+ [Cu(5-(4-N-hexadecylpyridiniumyl)-10,15,20triphenylporphyrin bromide)]2+ [Cu(meso-tetrakis(p-sulfonatophenyl)porphyrins)]2+ [Cu(5,10,15,20-tetraphenyl21H,23H-porphyrin)]2+ [Cu(5,10,15,20-tetrakis(4hydroxy-3methoxyphenyl)porphyrin)]2+ [Cu(3,5-diamino-1,2,4triazole)]2+ [Cu(1,2,4-triazole)]2+ [Cu(3,5-dimethyl-4-amino1,2,4-triazole)]2+ [Cu(3,4,5-triamino-1,2,4triazole)]2+ [Cu(phen)]2+ 2+

[Cu(phen)]

[Cu(5-Cl-phen)]2+

[Cu(5-Cl-phen)]2+ 2+

[Cu(2,9-Me2 -phen)]

[Cu(2,9-Me2 -phen)]2+ [Cu(4,7-diphenyl-phendisulfonate)]2+ [Cu(5-NH2 -phen)]2+ [Cu(5-NO2 -phen)]2+ [Cu(3-CO2 Et-4-Cl-phen)]2+ [Cu(3,8-(CO2 Et)2 -4,7-Cl2 phen)]2+ [Cu(2-Me-phen)]2+ [Cu(5-NH2 -2,9-Me2 -phen)]2+ [Cu(2,9-Et2 -phen)]2+ 2+

[Cu(2,9-nBu2 -phen)]

[Cu(5-NO2 -2,9-Me2 -phen)]2+ [Cu(3-ethynylphenanthroline)]2+ [Cu(trpn)(Im)](ClO4 )2 [Cu(baEtO)(Im)](ClO4 )2 [Cu(tren)](ClO4 )2 [Cu(Me6 tren)](ClO4 )2 [Cu(Me3 TACN)](ClO4 )2 [Cu(TPA)](ClO4 )2 [Cu(TPA)](ClO4 )2 [Cu(tripic)(NCMe)]PF6 [Cu2 (bistripic)(NCMe)2 ](PF6 )2

Cond (pH/electrode)

E1/2 CuI/II

E1/2 CuI/II vs. RHEa

EORR vs. Ref

EORR vs. RHEa

Ref.

Adlayer on Au (111), hanging meniscus, 0.1 M HClO4 GC disk, complex supported on Vulcan, 0.5 M H2 SO4 ,

–

–

0 V vs. RHE

0V

[27]

–

–

−0.1 V vs. SCE

0.14 V

[28]

Complex adsorbed on GC disk, 0.05 M H2 SO4 Complex adsorbed on GC disk, 0.05 M H2 SO4

–

–

–

–

[29]

–

–

−0.15 V vs. Ag/AgCl

0.13 V

[29]

Adsorbed on Ag electrode, 0.05 M H2 SO4 Paraffin-impregnated graphite electrode (PIGE), 0.1 M KCl (pH 4) Edge-plane pyrolytic graphite electrode, 0.5 M H2 SO4

–

–

−0.3 V vs. SCE

0.02 V

[30]

–

–

∼ 0 V vs. Ag/AgCl

0.44 V

[31]

0.10 V vs. Ag/AgCl

0.28 V

–

–

[32]

GC disk, supported on Vulcan, pH 7 Britton–Robinson GC disk, supported on Vulcan, pH 7 Britton–Robinson GC disk, supported on Vulcan, pH 7 Britton–Robinson GC disk, supported on Vulcan, pH 7 Britton–Robinson Edge-plane pyrolytic graphite electrode, pH 5.2, Britton–Robinson pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH Edge-plane pyrolytic graphite electrode, pH 5.2 Britton–Robinson buffer pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH Edge-plane pyrolytic graphite electrode, pH 5 Britton–Robinson buffer pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 5.3, Britton Robinson,

–

–

0.73 V vs. RHE

0.73 V

[55]

–

–

0.70 V vs. RHE

0.70 V

[59]

–

–

0.58 V vs. RHE

0.58 V

[59]

–

–

0.67 V vs. RHE

0.67 V

[59]

−0.19 V vs. SCE

0.37 V

−0.1 V vs. SCE

0.46 V

[37]

0. 025 V vs. NHE

0.313 V

0.01 V vs. NHE

0.30 V

[41]

−0.17 V vs. SCE

0.39 V

−0.07 V vs. SCE

0.49 V

[37]

0.05 V vs. NHE

0.34 V

0.04 V vs. NHE

0.33 V

[41]

0.05 V vs. SCE

0.59 V

0.05 V vs. SCE

0.59 V

[36]

0.31 mV vs. NHE

0.60 V

0.29 V vs. NHE

0.58 V

[41]

−0.18 V vs. SCE

0.38 V

−0.15 V vs. SCE

0.41 V

[40]

pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 100 mM NaClO4 , 20 mM NaAcO, 20 mM AcOH pH 4.8, 50 mM NaAcO, 50 mM AcOH, 1 M NaClO4 pH 6.4 Britton–Robinson buffer Pyrolytic graphite electrode, pH 7 phosphate buffer pH 1, 0.1 M HClO4 , on Vulcan C pH 1, 0.1 M HClO4 , on Vulcan C pH 1, 0.1 M HClO4 , on Vulcan C pH 1, 0.1 M HClO4 , on Vulcan C pH 7 Britton–Robinson buffer, on Vulcan C pH 2 Britton–Robinson buffer, on Vulcan C pH 2 Britton–Robinson buffer, on Vulcan C

0.02 V vs. NHE

0.31 V

0.01 V vs. NHE

0.30 V

[41]

0.075 V vs. NHE

0.37 V

0.040 V vs. NHE

0.33 V

[41]

0.090 V vs. NHE

0.38 V

0.065 V vs. NHE

0.36 V

[41]

0.15 V vs. NHE

0.44 V

0.130 V vs. NHE

0.42 V

[41]

0.215 V vs. NHE

0.51 V

0.205 V vs. NHE

0.50 V

[41]

0.285 V vs. NHE

0.58 V

0.275 V vs. NHE

0.57 V

[41]

0.335 V vs. NHE

0.63 V

0.305 V vs. NHE

0.59 V

[41]

0.340 V vs. NHE

0.63 V

0.260 V vs. NHE

0.55 V

[41]

0.390 V vs. NHE

0.68 V

0.080 V vs. NHE

0.37 V

[41]

0.28 V vs. NHE

0.56 V

0.10 V vs. NHE

0.39 V

[42]

−0.23 V vs. SCE −0.248 V vs. SCE

0.40 V 0.42 V

−0.05 V vs. SCE −0.10 V vs. SCE

0.58 V 0.57 V

[33] [34]

– – 0.18 V vs. RHE – 0.23 V vs. RHE

– – 0.18 V vs. RHE – 0.23 V

0.43 V vs. RHE 0.33 V vs. RHE 0.30 V vs. RHE 0.53 V vs. RHE 0.69 V vs. RHE

0.43 V 0.33 V 0.30 V 0.53 V 0.69 V

[35] [35] [35] [51] [51]

–

–

0.34 V vs. RHE

0.34 V

[51]

–

–

0.40 V vs. RHE

0.40 V

[51]

138

M.A. Thorseth et al. / Coordination Chemistry Reviews 257 (2013) 130–139

Table 1 (Continued) Catalyst

Cond (pH/electrode)

E1/2 CuI/II

E1/2 CuI/II vs. RHEa

EORR vs. Ref

EORR vs. RHEa

Ref.

[Cu(PMEA)](ClO4 )2

pH 7 Britton–Robinson buffer, on Vulcan C pH 7 Britton–Robinson buffer, on Vulcan C pH 7 Britton–Robinson buffer, on Vulcan C pH 1 0.1 M HClO4 , on Vulcan C pH 1 0.1 M HClO4 , on Vulcan C pH 7.3 borate buffer in solution at GC electrode pH 5.3 Britton–Robinson buffer, PG electrode pH 5.3 Britton–Robinson buffer, PG electrode pH 5.3 Britton–Robinson buffer, PG electrode pH 9 phosphate buffer, GC electrode Complex-coated thin film on GC electrode, 0.2 M PBS (pH 7) Complex-coated Au electrode, 0.1 M H2 SO4

0.37 V vs. RHE

0.37 V

0.69 V vs. RHE

0.69 V

[35]

0.42 V vs. RHE

0.42 V

0.69 V vs. RHE

0.69 V

[51]

0.52 V vs. RHE

0.52 V

0.69 V vs. RHE

0.69 V

[35]

0.42 V vs. RHE 0.22 V vs. RHE −0.1 V vs. SCE

0.42 V 0.22 V 0.58 V

0.53 V vs. RHE 0.40 V vs. RHE 0.0 V vs. SCE

0.53 V 0.40 V 0.68 V

[35] [35] [44]

0.0 V vs. SCE

0.56 V

0.0 V vs. SCE

0.56 V

[39]

0.0 V vs. SCE

0.56 V

–

–

[39]

−0.17 V vs. SCE

0.39 V

−0.05 V vs. SCE

0.51 V

[46]

−0.35 V vs. SCE –

0.43 V –

−0.25 V vs. SCE −0.07 V vs. Ag/AgCl

0.53 V 0.54 V vs. RHE

[45] [49]

–

–

∼ 0 V vs. SCE

0.30 V vs. RHE

[47]

−0.17 V vs. SCE −0.09 V vs. 3 M CE

0.38 V 0.64 V

−0.22 V vs. SCE −0.18 V vs. 3 M CE

0.33 V 0.55 V

[48] [50]

[Cu(PMAP)](ClO4 )2 [Cu(TEPA)](ClO4 )2 [Cu((NH2 )2 -TPA)](NO3 )2 [Cu(Piv2 -TPA)](NO3 )2 Cu with hexaaza-xylyl macrocycle [Cu(DPP)(L )]2+ [Cu(PDT)(L )]2+ [Cu(TPT)(L )]2+ 2+

[Cu2 (apyhist)2 Cl2 ] Cu with poly-l-histidine matrix

[Cu(2-(3-pyrrol-1-ylpropylimino-methyl)phenol)] ·24H 5,21OM (CH3 COOLi + CH3 COOH) Na14 [SiW9 O34 Cu3 (N3 )2 (OH)(H2 O)]2pH [CuIII (LH−4 )]− (L = 8,17-dioxapH 8, 0.1 M NaClO4 1,2,5,6,10,11,14,15-octaazatricyclo[13.3.1] eicosane-3,4,12,13-tetrone)

E1/2 of CuI/II couple and EORR were estimated from the data provided in references. When no specific number was quoted in the text, potentials were estimated from the graphs provided in the reference. Potentials were then referenced to the RHE by conversion of the reported reference potential to the NHE (0.197 V for Ag/AgCl, and 0.244 V for SCE). NHE potentials were then converted to RHE by assuming a 60 mV/pH positive shift by the Nernst equation.

been observed that the O2 binding plays a critical role in how the catalyst behaves. Progress toward this goal can be seen in recent work by Tahsini et al. [60]. In attempts to elucidate the reactive intermediate of O2 reduction, which could exist as a peroxo or bis␮-oxo species, both binuclear and mononuclear copper(II) systems were used to catalyze the reduction of O2 by decamethylferrocene. By comparison of the activation entropies of electron transfer from ferrocene derivatives to either the binuclear copper complex or the peroxo dicopper(II) intermediate, it was possible, for the first time, to determine what intermediates were produced and what roles they played. With greater knowledge of the mechanisms, Cu–ligand systems can be tailor made to mimic the environment in the active site of laccase. As seen in Fig. 8, all of the reported Cu complexes to date, which are likely di-Cu, have very poor activity compared to laccase. The first priority then is to obtain a functional tricopper model with activity comparable to laccase.

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