Accepted Article Title: Evident Improvement of Electrochemical Water Oxidation by Fine Tuning the Structure of Tetradentate N4 Ligands of Molecular Copper Catalysts Authors: Junyu Shen, Mei Wang, Jinsuo Gao, Hongxian Han, Hong Liu, and Licheng Sun This manuscript has been accepted after peer review and appears as an Accepted Article online prior to editing, proofing, and formal publication of the final Version of Record (VoR). This work is currently citable by using the Digital Object Identifier (DOI) given below. The VoR will be published online in Early View as soon as possible and may be different to this Accepted Article as a result of editing. Readers should obtain the VoR from the journal website shown below when it is published to ensure accuracy of information. The authors are responsible for the content of this Accepted Article. To be cited as: ChemSusChem 10.1002/cssc.201701458 Link to VoR: http://dx.doi.org/10.1002/cssc.201701458

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Evident Improvement of Electrochemical Water Oxidation by Fine Tuning the Structure of Tetradentate N4 Ligands of Molecular Copper Catalysts Junyu Shen,[a] Mei Wang,*[a] Jinsuo Gao,[b] Hongxian Han,[c] Hong Liu,[a] and Licheng Sun[a,d]

Two copper complexes, [L1Cu(OH2)](BF4)2 [1, L1 = N,N′-dimethyl-N,N′-bis(pyridin-2ylmethyl)-1,2-diaminoethane] and [L2Cu(OH2)](BF4)2 [2, L2 = 2,7-bis(2-pyridyl)-3,6-diaza2,6-octadiene], were prepared as molecular water oxidation catalysts. Complex 1 displayed an overpotential (η) of 1.07 V at 1 mA cm−2 and an observed rate constant (kobs) of 13.5 s−1 at η 1.0 V in pH 9.0 phosphate buffer solution, while 2 exhibited a significantly smaller η (0.70 V) to reach 1 mA cm−2 and a higher kobs (50.4 s−1) than 1 under identical test conditions. Additionally, 2 displayed a better stability than 1 in controlled potential electrolysis experiments in a Faradaic efficiency of 94% for O2 evolution at 1.58 V, when a casing tube was used for Pt cathode. The possible mechanism for 1- and 2-catalyzed O2 evolution reactions is discussed based on the experimental evidence. These comparative results indicate that fine tuning structures of tetradentate N4 ligands can bring about significant change in the performance of copper complexes for electrochemical water oxidation.

_________________________________________________________ [a] J. Shen, Prof. M. Wang, H. Liu, Prof. L. Sun State Key Laboratory of Fine Chemicals DUT-KTH Joint Education and Research Centre on Molecular Devices Dalian University of Technology (DUT), Dalian 116024 (PR China) E-mail: [email protected] (M. W.) [b] J. Gao School of Environmental Science and Technology Dalian University of Technology, Dalian 116024 (PR China) [c] Prof. H. Han State Key Laboratory of Catalysis Dalian Institute of Chemical Physics, Dalian 116023 (PR China) [d] Prof. L. Sun Department of Chemistry KTH Royal Institute of Technology, Stockholm 10044 (Sweden)  Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cssc2017xxxxx.

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Introduction The conversion of solar energy into storable chemical fuels with artificial photosynthetic systems is a holy cause pursued by many aspiring researchers in the field of catalysis, bioinorganic chemistry, and material science. Water oxidation (2H2O → O2 + 4H+ + 4e−) and proton reduction (2H+ + 2e− → H2) are two half reactions for the overall water splitting, of which the oxygen evolution reaction (OER) is more difficult than the hydrogen evolution reaction (HER) because the OER requires larger overpotential to transfer four protons and four electrons for the formation of an O−O bond. In Nature, the oxidation of water is highly efficiently catalyzed by a CaMn4O5 cluster in the photosystem II (PSII).[1] Although the research on electro- and photocatalytic water splitting has been continued for many decades, the water oxidation is still a bottleneck for large-scale production of H2 from water splitting with artificial photosynthetic systems. For constructing artificial systems, either one-pot photocatalyst

colloid

systems,

photoelectrochemical

(PEC)

cells,

or photovoltage

(PV)/electrolyzer apparatuses, one of the major challenges is developing highly efficient, robust, inexpensive, and scalable water oxidation catalysts (WOCs). In recent years, a large number of OER electrocatalysts, both heterogeneous and homogenous WOCs, have been reported and substantial progress in this subject area has been achieved. The attractive points of molecular WOCs are their tunable chemical structures and their availability for getting knowledge on OER mechanisms. More importantly, highly active and stable molecular WOCs are available for modifying the surface of semiconductor material photoanodes.[2] Many Ru and Ir noble metal-based molecular WOCs have been intensively studied in previous years.[3] In recent years, more attention has been attracted to developing first-raw metal-based WOCs,[3,4] such as manganese,[5] iron,[6] cobalt,[7,8] nickel,[9] and copper complexes.[10−18] Some of these earth-abundant molecular WOCs displayed very high catalytic activity, e.g., the rate constant (kobs) of 1400 s−1 was obtained for a molecular cobalt porphyrin 2

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WOC in pH 7 phosphate buffer solution (PBS).[8] Molecular copper catalysts have drawn particular attention for OER due to the diverse redox properties and the well-studied coordination chemistry of copper complexes, as well as their low cost and high earth reserve. Since the first molecular copper WOC was found by Mayer and co-workers in 2012,[10] a number of single-site and multinucleate copper WOCs have been successively reported in the following years.[11−18] Some studies have demonstrated that the catalytic activity and the overpotential of copper WOCs could be considerably altered with small structural tuning of ligands. For example, with changing one of the pyridyl unit in bipyridine (bpy) ligand to an imidazolyl, the onset overpotential of water oxidation was reduced by more than 100 mV, along with apparent decrease of OER catalytic activity from 100 s−1 for [(bpy)Cu(H2O)2]2+ to 35 s−1 for [Cu(pimH)(H2O)2]2+ [pimH = 2-(2'-pyridyl)-imidazole] under similar test conditions (pH 12‒12.5).[10,11] In another example, the overpotential of the copper(II) catalyst [(L)Cu]2− [L = N1,N1'-(1,2-phenylene)bis(N2-methyloxalamide)] for water oxidation was drastically reduced by 530 mV with introducing two electron-donating substituents (OMe) to the aromatic ring of ligand, while the kobs of the catalyst is decreased by a factor of 22.[12] Although simultaneous improvement of the OER activity and overpotential of catalysts is expected, diminishing in overpotential accompanied with decreasing in OER activity or vise verse is a common phenomenon resulting from fine tuning ligand structures of molecular catalysts. To develop OER molecular catalysts that have high activity, low overpotential, and good stability with designing of effective ligands, more studies are needed to accumulate the knowledge about the influence of electronic and conjugate effects of ligands on the activity and overpotential of molecular catalysts. Herein, we report the synthesis and performance of two well-characterized molecular copper catalysts, [L1Cu(OH2)](BF4)2 [1, L1 = N,N′-dimethyl-N,N′-bis(pyridin-2-ylmethyl)-1,2diaminoethane] and [L2Cu(OH2)](BF4)2 [2, L2 = 2,7-bis(2-pyridyl)-3,6-diaza-2,6-octadiene], 3

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for electrocatalytic OER. Appealingly, considerable decrease of overpotential was observed along with evident enhancement of OER catalytic activity for the molecular copper catalysts when diamine-dipyridine was tuned to diimine-dipyridine ligand. Copper(II) complex 2 bearing a diimine-dipyridine ligand displayed a 427 mV lower onset overpotential than that of the analogous complex 1 containing a diamine-dipyridine ligand for OER at pH 9.0, while the kobs of 2 (50.4 s−1) is about 3.7 times as high as that of 1. Additionally, 2 exhibited a better durability compared to 1 for water oxidation in pH 9.0 PBS. These results reveal that the conjugate effect and the type of coordinate N atoms of a ligand in molecular catalysts play an important role in the performance of molecular OER catalysts, and that it is possible to simultaneously enhance the activity, reduce the overpotential, and improve the stability of molecular catalysts through fine tuning ligand structures.

Results and Discussion Synthesis and characterization of 1 and 2 Two copper(II) complexes, [L1Cu(OH2)](BF4)2 (1) and [L2Cu(OH2)](BF4)2 (2), were conveniently synthesized by the reaction of Cu(BF4)2·6H2O with an equivalent of L1 or L2 in ethanol at RT. Complex 1 was obtained as blue crystal and 2 as purple crystal in high yields (86‒94%). The products were characterized by HRMS and elemental analysis, and the results are in good agreement with the proposed compositions of 1 and 2. The molecular structures of 1 and 2 were further determined by single crystal X-ray diffraction (Figure 1). The selected bond lengths and angles for 1 and 2 as well as crystallographic data are given in Tables S1 and S2. Both of these copper(II) complexes have a distorted square pyramidal geometry configuration with a tetradentate N4 ligand in the basal square plane and an aqua molecule at the apical site, being similar to the structures of previously reported analogous complexes [N4Cu(OH2)]2+.[13] The most noticeable difference in the 4

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structures of 1 and 2 is that complex 1 bears a diamine-dipyridine ligand, while 2 has a diiminedipyridine ligand, in which two C=N bonds are conjugated with the pyridine moieties. The lengths of C=N bonds in 2 are 1.251‒1.254 Å and those of the corresponding C−N bonds in 1 are 1.469‒1.486 Å. The coordination of the diimine of L2 to the CuII ion of 2 gives considerably shorter lengths of Cu−N1 [1.971(7) Å] and Cu−N3 bond [1.956(5) Å] compared to the lengths of Cu−N bonds between the diamine moiety of L1 and the CuII ion of 1 [2.024(5) and 2.028(5) Å]. In contrast, the avarage lengths of Cu−N bonds between the CuII ion and the N atoms of pyridine units are almost the same for 1 (2.015 Å) and 2 (2.013 Å). More importantly, the bond length between the CuII ion and the O of coordinated water molecule in 2 is 0.165 Å longer than that in 1. We envisaged that these subtle discrepancies in the bond lengths of 1 and 2 could influence the redox property and the catalytic behavior of these copper complexes in OER.

Figure 1. Molecular structures of [1]2+ (left) and [2]2+ (right) as ball-and-stick drawings. Counter-ions and hydrogen atoms are omitted for clarity.

The electronic absorbance spectrum of 1 in aqueous solution showed two intense π→π* absorptions arising from the pyridine units of L1 at λmax = 259 and 287 nm (Figure S1a), together with a broad weak absorption arising from d–d transition of CuII at λmax = 631 nm (Inset of Figure S1a), while the absorptions of 2 appeared at λmax = 283 nm with a shoulder peak at about 298 nm for the pyridine units of L2 and at 585 nm for the d–d transition of CuII. In comparison, under the same conditions free L1 and L2 each showed an intense band at 261– 5

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263 nm for π→π* absorptions of pyridine units in the UV-vis spectra (Figure S1b). Additionally, L2 displayed an extra broad band at λmax = 366 nm, which is ascribed to the π→π* absorption of the C=N double bonds of the diimine conjugated with pyridine units. To explore the stability of the aqua-coordinated copper complexes in the phosphate electrolyte used for electrochemical studies, the UV-vis spectra of 1 and 2 were measured in 0.1 M PBSs at pH 9.0, which were exactly identical to their spectra measured in pure aqueous solutions (Figure S2a). These observations suggested that the water molecule coordinating to the CuII ion in the copper complexes has not been replaced by the phosphate of electrolyte. Moreover, 1 and 2 both displayed the same EPR spectra in pure water and in 0.1 M pH 9.0 PBS (Figure S3), with the isotropic g values of 2.10 and 2.130 for 1 and 2, respectively, calculated by simulation (Gauss line shape).[19] This evidence provided further proofs for the intactness of these copper complexes in the initial test electrolyte.[10] What’s more, the UV-vis spectra that were measured before and after the basic solutions (pH 9) of 1 and 2 stood for a week under air did not show observable difference (Figure S2b), which indicated that 1 and 2 were stable in pH 9 solutions in the presence of O2. Besides, the UV-vis spectra (Figures S4a and S5a) measured with titration of 5 M NaOH solution to the PBSs of 1 and 2, respectively, in the pH range of 7 to 12 were used to evaluate the pKa values of these copper complexes. The plots of absorbances of 1 and 2 as a function of pH are shown in Figures S4b and S5b, in which the inflection points of the curves appear at pH 10.5 for 1 and 10.7 for 2. Taken together, all of the evidence showed that 1 and 2 existed as their original complexes with no dissociation nor deprotonation of the coordinated aqua molecule in PBSs at pH 9.0. Electrochemical properties of 1 and 2 Initially, the electrochemical properties of 1 and 2 were studied by cyclic voltammetry and differential pulse voltammetry in 0.1 M PBSs at pH 9.0 with a glassy carbon working electrode (0.071 cm2), a Ag/AgCl reference electrode, and a platinum wire counter electrode. All 6

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potentials given in this paper are versus normal hydrogen electrode [NHE, E(NHE) = E(Ag/AgCl) + 0.197 V]. As shown in Figure 2a, the cyclic voltammogram (CV) of 1 displayed the current with the onset potential at 1.43 V (defined as the potential at a current density of 0.1 mA cm−2). Although the CuII/CuIII wave of 1 could not be observed from CV, it was observed in differential pulse voltammogram (DPV) at Epa = 1.58 V, together with a weak catalytic peak at Epa = 1.84 V (Figure 2b). In comparison, the CuII core of 2 is easily oxidized to higher oxidation states than that of 1. The event with an onset potential of 1.0 V in the CV of 2 is ascribed to the oxidation process of CuII/CuIII couple. At potential more positive than 1.4 V, the current density of the PBS containing 2 rose rapidly and oxygen bubbles appeared on the surface of working electrode. Accordingly, the DPV of 2 with anodic scanning showed the oxidation peak of CuII/CuIII couple at 1.34 V together with an extra peak at 1.62 V (Figure 2b), which is tentatively assigned to the further oxidation process of CuIII(OH)/CuIII(O∙) or CuIII/CuIV couple accompanied with catalytic oxidation of water. Noticeably, the two oxidative events of 2 are negatively shifted by 220‒240 mV as compared to the corresponding oxidative events of 1. More importantly, the catalytic current enhancement for 2 was much more pronounced compared to that for 1. The catalytic current of 2 was increased by factors of 7.0‒4.4 at applied potentials of 1.5‒1.7 V as compared to that of 1. For 2, the overpotential required to reach a catalytic current density of 1 mA cm−2 is about 700 mV, which is 370 mV smaller than that required by catalyst 1; it is also smaller than the overpotentials at 1 mA cm−1 observed for most previously reported Cu-based molecular WOCs measured in buffer solutions at pH 7‒11 (Table S3),[14] except for [(TGG4−)Cu(H2O)]2− (TGG4− = triglycylglycine macrocyclic ligand),[15] [(Py3P)Cu] (Py3P = N,N-bis(2-(2-pyridyl)ethyl)pyridine-2,6-dicarboxamidate) and [Cu(en)2]2+ (en = 1,2-ethylenediamine).[16] These molecular copper catalysts display overpotentials of 620, 544 and 604 mV at 1 mA cm−2, respectively, in pH 8‒11 PBSs. It is appealing that the little change in the structure of tetradentate N4 ligand L1 brings about an evident enhancement of 7

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catalytic activity together with an apparent decrease of overpotential of the copper catalyst for OER. The distinctively different performances of 1 and 2 in electrocatalytic water oxidation should be related to the subtle discrepancy in the structures of 1 and 2, viz., the longer Cu−O bond of 2 relative to that of 1 and the shorter σ-coordination Cu−N bonds between the Cu core and the diimine moiety of L2 compared with those between the Cu core and the diamine moiety of L1.

(b)

(a) blank 1 2

4

0.8

3 j / mA cm

2

j / mA cm

1.62 V

2

0.6

blank 1 2

2 1

0.4 1.34 V

0

1.84 V 1.58 V

0.2

0.0

0.4

0.8

1.2 E / V vs. NHE

1.6

0.8

2.0

1.0

1.2

1.4

1.6

1.8

E / V vs. NHE

Figure 2. (a) CVs of 1 and 2 (both in 1.0 mM) as well as the blank CV of a glassy carbon electrode in 0.1 M PBSs at pH 9.0 at a scan rate of 100 mV s−1. (b) DPVs in the absence and presence of 1 and 2 at a scan rate of 50 mV s−1.

To the best of our knowledge, only a few non-noble metal Schiff base complexes have been used as WOCs.[20] To understand whether the catalytic current of 2 is ligand-based or metal center-based, the CVs of L2 and a zinc complex [L2Zn(OH2)](BF4)2, an analogue of 2, were studied. Neither L2 nor [L2Zn(OH2)](BF4)2 showed oxidation event in the CVs within a scanning range of 0.4‒2.0 V (Figure S6). It indicated that both L2 and [L2Zn(OH2)]2+ were no OER catalytic activity under test conditions. This evidence demonstrated that the catalytic activity of 2 for water oxidation was as the result of oxidation of CuII center, and that the C=N 8

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bond-containing N4 ligand was innocent in the catalytic redox process and only played a role in stabilization of high oxidation states of the active intermediates. The Pourbaix diagrams of 1 and 2 showed that their CuII/CuIII oxidation peak potentials were pH-dependent in the pH range of 8.0–10.5 (Figure S7), shifting negatively by 59.8 and 57.2 mV per pH, which are consistent with the decrease of about 59 mV per pH unit predicted by the Nernst equation for a 1e−/1H+ redox couple. These observations suggested that the oxidation of the CuII forms of 1 and 2 to CuIII species took place by a proton coupled electron transfer (PCET) process in moderately basic solutions. To have an overall understanding of the electrochemical properties of 1 and 2 in water, the CVs and DPVs of these copper complexes were also measured in the range of 0.5 V to −0.5 V in 0.1 M PBSs at pH 9. Complex 1 exhibited a quasi-reversible redox event with E1/2 at −0.09 V (Figure S8a) at a scan rate of 100 mV s−1. In contrast, 2 displayed an event with Epa at about 0.33 V together with an irreversible reduction peak at Epc = −0.22 V in its CV (Figure S8b). The DPV of 2 clearly showed two peaks at 0.30 and −0.25 V, the latter of which was in a similar potential to that observed for the reduction of free L2 (Figure S8c). In comparison, the DPV of 1 showed only a reduction peaks of CuII/CuI couple at −0.12 V, while the DPV of [L2Zn(OH2)](BF4)2 showed a reduction peak of ligand at about –0.26 V. These observations verified that the first reduction peak in the CV (Figure S8b) of 2 was attributed to the CuII/CuI couple and the second one to the reduction of coordinating ligand, which began to be reduced when the applied potential was more negative than 0 V. These results showed that the structural tune of ligands is an effective strategy to considerably adjust the redox properties of molecular catalysts. Kinetics of OER catalyzed by 1 and 2 The kinetics of OER catalyzed by 1 and 2 was studied by measuring the CVs with different concentrations of catalysts and by varying scan rates in pH 9.0, 0.1 M PBSs at 23 °C. Figures 9

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S9 and S10 show that the current maxima of the CuI/II waves of 1 and 2 vary linearly with the square root of the scan rate, indicative of a diffusion-controlled behavior of the copper catalysts under test conditions.[21] In addition, the catalytic current (ic) and the peak current (id) of noncatalytic CuII reduction both displayed linear relations with [1] and [2] from 0.25 to 1.25 mM (Figures S11 and S12), indicating the first-order dependence of id and ic against the concentration of catalyst. In general, the intrinsic rate constant (kcat) of a molecular catalyst for the first order or pseudo-first order OER in aqueous solutions at 23 °C can be estimated by the following simplified equation (eq. 1) when the jc‒E plot shows an “S-shaped” catalytic response, viz., a potential-independent plateau, at high applied potentials.[22] ic/id = 1.38(kcat/ν)1/2

(eq. 1)

Where ic is the limiting catalytic peak current of the “S-shaped” catalytic response, id is the peak current of a reversible or quasi-reversible noncatalytic wave, and ν is the scan rate. In practice, the plateau is often not achieved for some homogeneous electrocatalytic systems due to obvious water oxidation current from working electrode at forcing positive potentials required to reach the plateau, which is just the case for catalysts 1 and 2. In such situation, the observed rate constant (kobs) can be estimated according to eq. 2.[23,24] ic'/id = 1.38(kobs/ν)1/2

(eq. 2)

Where ic' is the highest achievable current obtained at the most positive applied potential before an apparent background current appearing from the working electrode, and it should be measured at scan rates where the catalytic current is invariable; kobs is a practical, overall rate constant at a particularly applied potential, which is a lower limit of kcat (kobs < kcat; only in the plateau region, kobs = kcat).[23] Figure 3 shows that the constant current responses are independent of scan rate when ν is larger than 500 mV s−1 for 1 and 400 mV s−1 for 2. So the ic' values of 1 and 2 were obtained 10

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from the CVs measured at the scan rate of 500 and 400 mV s−1 (Insets in Figure 3), respectively, with deduction of background current. As the CuII/CuIII waves of 1 and 2 cannot be clearly distinguished from the catalytic events in their CVs, the non-catalytic currents of CuI/CuII waves of 1 and 2 were used to estimate the diffusive peak currents (id). From the obtained values of ic', id, and the minimum ν to get constant catalytic current responses, the kobs calculated on the basis of eq. 2 is 13.5 s−1 for 1 and 50.4 s−1 for 2 at applied potential of 1.70 V (corresponding to an overpotential of 1.0 V). Although the kobs values reported here underestimate the intrinsic rate constants of these copper catalysts, the kobs value of 2 is comparable to the highest rate constant values reported to date for molecular copper OER catalysts working in neutral or moderate basic solutions (pH 7–11, Table S3).[13−17]

4

2 1

3 2 1

blank 1 at 1.70 V jc'/jd = 7.18 jd

(b)

600 mV s1

6

jc'

0

4

2

-1 -0.5 0.0 0.5 1.0 1.5 2.0 E / V vs. NHE

blank 2 at 1.70 V jc'/jd = 15.49

4 2

1

300 mV s 1 400 mV s 1 500 mV s

 jc'

jd

0 0.0

0 0.4

1

50 mV s 1 100 mV s 1 200 mV s

6

 2

5

3 j / mA cm 2

j / mA cm

2

4

300 mV s1 400 mV s1 500 mV s1

j / mA cm2

50 mV s1 100 mV s1 200 mV s1

j / mA cm

(a) 5

0.4 0.8 1.2 1.6 E / V vs. NHE

0 0.8

1.2 E / V vs. NHE

1.6

0.4

2.0

0.8 1.2 E / V vs. NHE

1.6

Figure 3. Linear sweep voltammograms of (a) 1 and (b) 2 (both in 1.0 mM) in 0.1 M PBSs at pH 9.0 with varying scan rates. Insets: cyclic voltammograms of (a) 1 at a scan rate of 500 mV s−1 and (b) 2 at a scan rate of 400 mV s−1.

Stability of 1 and 2 as molecular OER electrocatalysts To explore the stability of 1 and 2 under test conditions, controlled potential electrolysis (CPE) experiments were carried out with fluorine-doped tin oxide (FTO, 1 cm2) as a working electrode 11

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in pH 9.0, 0.1 M PBSs at 1.58 V in a home-made H-type electrochemical cell. Figure 4a shows that the current densities of 1 and 2 are decreased by about 54% and 41%, respectively, during 1 h of CPE experiments. It was observed that the surface of Pt counter electrode turned dark brown and the intensity of UV-vis absorption of the electrolyte containing catalyst apparently attenuated after CPE experiments (Figure S13). To avoid the influence of counter electrode and the H2 evolved from the cathode on copper catalysts, the Pt cathode was put in a casing tube fitted with a porous glass frit bottom (G3, 0.071 cm2) to isolate catalyst from the Pt wire and the generated H2 (Figure S14). The casing tube contained 0.1 M PBS at pH 9.0 without catalyst. When the CPE experiment was carried out with Pt cathode in a casing tube, the initial current density is only 66%‒70% of that attained from CPE experiments without using casing tube at the same potential, primarily due to the increase of system resistance. With adoption of a casing tube for Pt cathode, 1 and 2 both displayed better stability than in the case of one-pot electrolysis. The current density was attenuated by about 44% for 1 and 26% for 2 during 1 h of CPE experiments (Figure 4a). Complex 2 displayed a better stability than 1. The possible reason is that the predominant degradation path for catalyst 2, viz., degradation through reduction and hydrogenation of the C=N bonds in the ligand, was avoided when the working electrode was isolated from Pt cathode. In the case of adoption of a casing tube for Pt wire, the slow attenuation in current in the first hour of CPE experiments was predominantly attributed to pH changes of the electrolytes during electrolysis, as the UV-vis spectra of electrolytes containing 1 or 2 before and after electrolysis exhibited only very slight attenuation in the absorption intensity (Figure S15). The pH values of electrolytes decreased to about 8.8 in the anodic chamber and increased to 9.3 in the casing tube of Pt cathode after 1 h of electrolysis. It is worthy of note that the current density of 2 was almost completely recovered when the pH was readjusted to 9.0 after 1 h of electrolysis (Figure 4b), indicating that 2 remained its catalytic activity over 2 h under test conditions. When such pH adjustment of the resulting electrolyte 12

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was repeated for the third time, the activity of 2 began to decrease apparently. With further addition of NaOH solution to adjust the pH of electrolyte back to 9.0, the current density could not recover, most possibly due to the decomposition of the catalyst after long-time CPE experiments.

(b) 0.7

(a) 0.7

0.6

0.5 j / mA cm

j / mA cm

2

0.5 0.4 0.3

0.4 0.3

0.2

0.2

0.1

0.1

0.0

0.0

0

1000

2000 Time / s

blank 2 with a casing tube adjust pH of electrolyte to 9.0, the 1st time adjust pH of electrolyte to 9.0, the 2nd time adjust pH of electrolyte to 9.0, the 3rd time

0.6

2

blank 1 with no casing tube 1 with a casing tube 2 with no casing tube 2 with a casing tube

3000

0

1000

2000 Time / s

3000

Figure 4. (a) Catalytic currents obtained over 1 h of CPE experiments without (blank) and with 1 or 2 (both in 1.0 mM) on an FTO electrode (1.0 cm2) at 1.58 V in 0.1 M PBSs at pH 9.0. (b) Catalytic currents obtained over 4 h of CPE experiment of 2 with adjusting pH back to 9.0 by addition of 1 M NaOH to the electrolyte after each hour.

Faradaic efficiency experiments were carried out in a custom built gas-tight electrochemical cell fitted with a casing tube for Pt cathode. GC analysis showed that about 2.2 μmol of O2 was evolved in the gas phase of the system during the electrolysis of 2 at 1.58 V for 1 h (The oxygen dissolved in the electrolyte was neglected), corresponding to a Faradaic efficiency of ~94%, while 1 displayed a lower Faradaic efficiency of ~69% during CPE experiment under identical conditions (Figure S16). To explore whether complex 2 functioned as a molecular catalyst for OER or only a precursor for electrodeposition of copper oxides under test conditions, the FTO working 13

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electrode was rinsed with deionized water after 3 h of CPE experiment with a casing tube for Pt counter electrode, and reused in a fresh 0.1 M PBS at pH 9.0 in the absence of catalyst. It was found that the CV of rinsed FTO electrode was identical to the blank CV of the unused FTO (Figure S17). The SEM (Figure S18), EDX (Figure S19), and XPS (Figure S20a) showed no evidence for deposition of copper oxide on the surface of the used FTO electrode. According to dynamic light scattering (DLS) spectra (Figure S20b), no nanoparticles were formed in the electrolyte after 3 h of CPE experiment at 1.58 V. All results from post-analyses after 3 h of electrolysis with a casing tube for Pt cathode supported that complex 2 acted as a molecular catalyst for OER under test conditions over 3 h. In strong alkaline buffer solution (pH > 11), 2 gradually decomposed to form a Cu-Pi film on the surface of FTO electrode at the applied potential higher than 1.2 V, as evidenced by SEM images (Figure S21a,b), EDX (Figure S21c), and XPS (Figure S22). Plausible mechanism for water oxidation catalyzed by 1 and 2 The mechanism for water oxidation catalyzed by 1 and 2 is presumed according to the available experimental evidence and referring to the mechanisms proposed for other single-site molecular copper catalysts in the literatures (Scheme 1).[15] The UV-vis and EPR spectroscopic studies (Figures S2 and S3) revealed that both 1 and 2 maintained their electronic structures in pH 9 PBSs. The initial step of water oxidation catalyzed by 1 and 2 is assumed to be the oxidation of [LCuII(OH2)]2+ (L = L1, L2) to [LCuIII(OH)]2+, on the basis of the experimental facts that the coordinating aqua molecule in 1 and 2 was not deprotonated in PBS at pH 9.0 and the oxidation of the CuII forms of 1 and 2 to CuIII species took place by a PCET process in the pH of 8.0– 10.5. The kinetic studies revealed that the id and ic were both in the first-order dependence on [1] and [2], which suggests that the OER catalyzed by 1 and 2 should occur through a mononuclear mechanism. Therefore, it is assumed that [LCuIII(OH)]2+ is further oxidized to [LCuIV=O]2+ or [LCuIII(O∙)]2+ intermediate via a PCET process.[15] Once formed, the CuIV or 14

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CuIII intermediate could be attacked by H2O to generate [LCuII(OOH)]+. Considering that the catalytic water oxidation reaction is first order in [HPO42−] (Figure S23) and the kinetic isotope effects (KIE) of 2.25 and 2.23 for 1 and 2 (Figure S24), respectively, the [LCuII(OOH)]+ intermediate is most possibly formed through a rate-limiting atom-proton transfer (APT) process as proposed for single-site RuII,[25] CoII,[8,26] and CuII catalysts (For more explanation, see the caption of Figure S23).[15] The formation of copper(II) hydroperoxyl intermediates in 1and 2-catalyzed OERs were tested by using horseradish peroxidase (HRP, a special catalyst for hemolysis of the peroxide bond of H2O2 to form •OH radicals) and Ampliflu red (AR, a reliable titrant for •OH) reagents.[27] The resulting blue electrolytes after 3 h of electrolysis of 1 and 2, respectively, turned pink upon successive addition of HRP and AR (Figure S25). This observation together with the results of control experiments supports the existence of peroxide intermediates in the electrolytes. The further oxidation of [LCuII(OOH)]+ would result in O2 release, accompanied with deprotonation. In the meantime, coordination of H2O to the CuII center leads to the regeneration of the initial catalyst, [LCuII(OH2)]2+ to complete the catalytic cycle.

Scheme 1. Plausible mechanism for the OER catalyzed by copper complexes 1 and 2.

Conclusions In summary, two water soluble copper(II) complexes, one bearing a diamine-dipyridine ligand 15

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(1) and the other containing a diimine-dipyridine ligand (2), were prepared and well characterized. Although X-ray single crystal diffraction analyses showed that 1 and 2 had very similar structures, they displayed quite different electrochemical redox properties in moderately basic solutions. It is really attractive that considerable decrease in overpotential was obtained along with evident enhancement of OER catalytic activity for the molecular copper catalysts when diamine-dipyridine was tuned to diimine-dipyridine ligand. Specifically, complex 2 exhibited an overpotential of 700 mV to reach 1 mA cm−2 current density in 0.1 M PBS at pH 9.0, which is 370 mV lower than that required by the analogous complex 1 and also lower than the overpotentials at 1 mA cm−2 observed for most previously reported Cu-based molecular WOCs measured in solutions at pH 7‒11. More importantly, the observed rate constant (kobs) of 2 (50.4 s−1) at an overpotential of 1.0 V is about 3.7 times as high as that of 1, which is comparable to the highest rate constant values reported to date for molecular copper OER catalysts working in neutral or moderately basic solutions. In addition, 2 displayed a better catalytic durability compared to 1 for water oxidation in pH 9.0 PBS. These results reveal that fine tuning the conjugate effect and the type of coordinate N atoms of a N4 ligand in molecular catalysts has an important influence on the performance of catalysts in electrochemical water oxidation. The comparative study of these two structurally similar molecular copper(II) catalysts for OER shows that it is possible to simultaneously enhance the activity, reduce the overpotential, and improve the stability of molecular catalysts through tuning the structure of ligands.

Experimental Section Materials. Manipulations for preparation of the copper complexes were carried out under pure N2 by using standard Schlenk techniques. Commercially available chemicals, Cu(BF4)2·6H2O, 2-acetylpyridine,

1,2-ethanediamine,

N,N'-dimethylethylenediamine, 16

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and

2-

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(chloromethyl)pyridine hydrochloride, were purchased from local suppliers and used as received. Ligands L1 and L2 were prepared according to the literature procedures.[28] Glassy carbon electrode, fluorine-doped tin oxide (FTO) glass plate, and platinum wire were purchased from Tianjin Gaoss Union for the electrochemical studies. All buffers were prepared with deionized water (18 MΩ-cm resistivity). Instruments. NMR Spectra were collected with a Varian INOVA 500 NMR spectrometer. Mass spectra were recorded with HP 1100 HPL/ESI-DAD-MS and Waters/Micromass LC/QTOF-MS instruments. Elemental analyses were performed with a Thermoquest-Flash EA 1112 elemental analyzer. UV-Vis absorption measurements were carried out on an Agilent 8453 spectrophotometer. SEM images and EDX spectra were obtained with a FEI Nova NanoSEM 450 instrument equipped with an EDX detector. XPS surveys were acquired with a ThermoFisher ESCALAB 250Xi surface analysis system. The measurement of dynamic light scattering (DLS) spectra were measured with a Zetasizer Nano ZS90 instrument. EPR spectra were collected on a Bruker electron paramagnetic resonance spectrometer (A200) with microwave frequency of 9.538 GHz at RT. For the simulation of these EPR spectra, the ratio of theoretical gyromagneticratio [n(63Cu)/n(65Cu) = 0.935] was used for the set of the Cu isotope parameters.[29] The experimental EPR spectra of 1 and 2 were roughly simulated using following EPR parameters: A(63Cu) = 63 G，A(65Cu) = 67 G, A(2N) = 11.5 G for 1 and A(63Cu) = 77 G, A(65Cu) = 80 G, A(4N) = 11.4 G for 2. The simulations gave giso = 2.10 with A(N) = 14 G and g = 2.045 for 1 and giso = 2.130 with A(N) = 14 G and g = 2.030 for 2. Preparation of [L1Cu(OH2)](BF4)2 (1). The mixture of Cu(BF4)2·6H2O (0.345 g, 1.0 mmol) and L1 (0.270 g, 1.0 mmol) in ethanol (40 mL) was stirred at RT for 8 h. The solution was concentrated to about 20 mL by evaporation under vacuum, and the blue crystals of 1 were obtained by diffusing diethyl ether into the resulting solution. Yield: 0.45 g (86%); elemental analysis calcd for C16H24N4OB2F8Cu (%): C 36.57, H 4.60, N 10.66; found: C 36.59, H 4.50, 17

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N 10.68; TOF-MS: calcd for [M − H2O − 2BF4]2+ (C16H22N4Cu): m/z 166.5570; found: 166.5574. Preparation of [L2Cu(OH2)](BF4)2 (2). Purple crystals of 2 were prepared with L2 as a ligand in an essentially identical protocol as that used for the preparation of 1. Yield: 0.49 g (94%); elemental analysis calcd for C16H20N4OB2F8Cu (%): C 36.85, H 3.87, N 10.74; found: C 37.12, H 3.67, N 11.00; TOF-MS: calcd for [M − H2O − 2BF4]2+ (C16H18N4Cu): m/z 164.5415; found: 164.5405. Preparation of [L2Zn(OH2)](BF4)2. Colorless crystals of [L2Zn(OH2)](BF4)2 were prepared by the reaction of Zn(BF4)2·H2O and L2 in ethanol in an identical protocol as that used for the preparation of 1. Yield: 0.37 g (71%); elemental analysis calcd for C16H20N4OB2F8Zn (%): C 36.72, H 3.85, N 10.71; found: C 36.89, H 3.78, N 10.76; TOF-MS: calcd for [M − H2O − 2BF4]2+ (C16H18N4Zn): m/z 165.0411; found: 165.0389. Crystallographic structure determinations. The single-crystal X-ray diffraction data were collected on a Bruker Smart Apex II CCD diffractometer with a graphite-monochromated MoK radiation ( = 0.071073 Å) at 296 K using the -2 scan mode. Data processing was accomplished with the SAINT processing program. Intensity data were corrected for absorption by the SADABS program. All structures were solved by direct methods and refined on F2 against full-matrix least-squares methods. Non-hydrogen atoms were refined anisotropically. Hydrogen atoms were located by geometrical calculation. Crystallographic data and selected bond lengths and angles for 1 and 2 are given in Tables S1 and S2 (CCDC-1526273 for 1 and -1526272 for 2). Electrochemistry studies. All electrochemical measurements were performed with a model CHI660E electrochemical workstation (CH instruments). Cyclic voltammetry experiments were carried out in a three-electrode cell under argon. The working electrode was a glassy carbon electrode disc (0.071 cm2), the reference electrode was an aqueous Ag/AgCl electrode, 18

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and the counter electrode was a platinum wire. The solution of 0.1 M phosphate buffer was used as supporting electrolyte, which was degassed by bubbling with argon for 15 min before measurement. All potentials are reported versus the normal hydrogen electrode (NHE) by addition of 0.197 V to the experimentally measured values. CPE experiments were carried out in a home-made H-type electrochemical cell with an FTO (1.0 cm2) glass slide as working electrode. The auxiliary electrode was a platinum wire which was protected by a casing pipe and the reference electrode was a commercially available aqueous Ag/AgCl electrode. The sample was bubbled with argon for 20 min before measurement and the CPE experiments were carried out under argon with constantly stirring. The Faradaic efficiencies were determined from CPE experiments of the solutions of 1 and 2 in 0.1 M phosphate buffer at pH 9.0 in a custom built gas-tight electrochemical cell at an applied potential of 1.58 V vs NHE for 60 min. The gas in the headspace of the cell was analyzed by CEAULIGHT GC-7920 gas chromatograph equipped with a 5 Å molecular sieve column (2 mm × 2 m) during the electrolysis and the oxygen dissolved in the solution was neglected.

Acknowledgements We are grateful to the Basic Research Program of China (No. 2014CB239402), the National Natural Science Foundation of China (Grant Nos. 21373040 and 21673028), the Swedish Energy Agency, the Swedish Research Council and the K & A Wallenberg Foundation for financial support of this work.

Keywords: Amine-pyridine ligand, Copper complex, Electrocatalysis, Imine-pyridine ligand, Water oxidation

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[1] a) J. P. McEvoy, G. W. Brudvig, Chem. Rev. 2006, 106, 4455−4483; b) N. Cox, D. A. Pantazis, F. Neese, W. Lubitz, Acc. Chem. Res. 2013, 46, 1588−1596. [2] M. Wang, Y. Yang, J. Shen, J. Jiang, L. Sun, Sustainable Energy & Fuels 2017, DOI: 10.1039/c7se00222j. [3] a) B. Limburg, E. Bouwman, S. Bonnet, Coord. Chem. Rev. 2012, 256, 1451−1467; b) X. Liu, F. Wang, Coord. Chem. Rev. 2012, 256, 1115−1136; c) D. G. H. Hetterscheid, J. N. H. Reek, Angew. Chem. Int. Ed. 2012, 51, 9740−9747; d) M. D. Kärkäs, O. Verho, E. V. Johnston, B. Åkermark, Chem. Rev. 2014, 114, 11863−12001; e) J. D. Blakemore, R. H. Crabtree, G. W. Brudvig, Chem. Rev. 2015, 115, 12974−13005. [4] M. D. Kärkäs and B. Åkermark, Dalton Trans. 2016, 45, 14421−14461. [5] a) R. Brimblecombe, A. Koo, G. C. Dismukes, G. F. Swiegers, L. Spiccia. J. Am. Chem. Soc. 2010, 132, 2892−2894; b) E. A. Karlsson, B.-L. Lee, T. Åkermark, E.V. Johnston, M. D. Kärkäs, J. Sun, Ö. Hansson, J.-E. Bäckvall, B. Åkermark, Angew. Chem., Int. Ed. 2011, 50, 11715−11718; c) K. J. Young, M. K. Takase, G. W. Brudvig, Inorg. Chem. 2013, 52, 7615−7622. [6] a) J. L. Fillol, Z. Codola, I. Garcia-Bosch, L. Gomez, J. J. Pla, M. Costas, Nat. Chem. 2011, 3, 807−813; b) Z. Codolà, L. Gómez, S. T. Kleespies, L. Q. Jr, M. Costas, J. Lloret-Fillol, Nat. Commun. 2015, 6, 5865−5875; c) M. Okamura, M. Kondo, R. Kuga, Y. Kurashige, T. Yanai, S. Hayami, V. K. K. Praneeth, M. Yoshida, K. Yoneda, S. Kawata, S. Masaoka, Nature, 2016, 530, 465−468. [7] a) Q. Yin, J. M. Tan, C. Besson, Y. V. Geletii, D. G. Musaev, A. E. Kuznetsov, Z. Luo, K. I. Hardcastle, C. L. Hill. Science 2010, 328, 342−345; b) D. K. Dogutan, R. McGuire, D. G. Nocera, J. Am. Chem. Soc. 2011, 133, 9178−9180. [8] D. Wang, J. T. Groves, Proc. Natl. Acad. Sci. 2013, 110, 15579−15584. [9] a) M. Zhang, M.-T. Zhang, C. Hou, Z.-F. Ke, T.-B. Lu, Angew. Chem., Int. Ed. 2014, 53, 20

This article is protected by copyright. All rights reserved.

10.1002/cssc.201701458

ChemSusChem

13042−13048; b) Y. Han, Y. Wu, W. Lai, R. Cao, Inorg. Chem. 2015, 54, 5604−5613. [10] S. M. Barnett, K. I. Goldberg, J. M. Mayer, Nat. Chem. 2012, 4, 498−502. [11] L. A. Stott, K. E. Prosser, E. K. Berdichevsky, C. J. Walsby, J. J. Warren, Chem. Commun. 2017, 53, 651−654. [12] P. Garrido-Barros, I. Funes-Ardoiz, S. Drouet, J. Benet-Buchholz, F. Maseras, A. Llobet, J. Am. Chem. Soc. 2015, 137, 6758−6761. [13] a) J. S. Pap, L. Szyrwiel, D. Sranko, Z. Kerner, B. Setner, Z. Szewczuk, W. Malinka, Chem. Commun. 2015, 51, 6322−6324; b) F. Yu, F. Li, J. Hu, L. Bai, Y. Zhu, L. Sun, Chem. Commun. 2016, 52, 10377−10380; c) J. Shen, M. Wang, P. Zhang, J. Jiang, L. Sun, Chem. Commun. 2017, 53, 4374−4377. [14] a) X.-J. Su, M. Gao, L. Jiao, R.-Z. Liao, P. E. M. Siegbahn, J.-P. Cheng, M.-T. Zhang, Angew. Chem. Int. Ed. 2015, 54, 4909−4914; b) L.-L. Zhou, T. Fang, J.-P. Cao, Z.-H. Zhu, X.-T. Su, S.-Z. Zhan, J. Power Sources 2015, 273, 298−304; c) R. Terao, T. Nakazono, A. R. Parent, K. Sakai, ChemPlusChem 2016, 81, 1064−1067; d) R.-J. Xiang, H.-Y. Wang, Z.J. Xin, C.-B. Li, Y.-X. Lu, X.-W. Gao, H.-M. Sun, R. Cao, Chem. Eur. J. 2016, 22, 1602−1067; e) A. Prevedello, I. Bazzan, N. D. Carbonare, A. Giuliani, S. Bhardwaj, C. Africh, C. Cepek, R. Argazzi, M. Bonchio, S. Caramori, M. Robert, A. Sartorel, Chem. Asian J. 2016, 11, 1281−1287; f) T.-T. Li, Y.-Q. Zheng, Dalton Trans. 2016, 45, 12685−12690. [15] M.-T. Zhang, Z. Chen, P. Kang, T. J. Meyer, J. Am. Chem. Soc. 2013, 135, 2048−2051. [16] a) M. K. Coggins, M. T. Zhang, N. Song, T. J. Meyer, Angew. Chem. Int. Ed. 2014, 53, 12226−12230; b) C. Lu, J. Du, X.-J. Su, M.-T. Zhang, X. Xu, T. J. Meyer, Z. Chen, ACS Catal. 2016, 6, 77−83. [17] T. Fang, L.-Z. Fu, L.-L. Zhou, S.-Z. Zhan, Electrochim. Acta. 2015, 161, 388−394. [18] a) T. Zhang, C. Wang, S. Liu, J.-L. Wang, W. Lin, J. Am. Chem. Soc. 2014, 136, 273−281; 21

This article is protected by copyright. All rights reserved.

10.1002/cssc.201701458

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b) D. L. Gerlach, S. Bhagan, A. A. Cruce, D. B. Burks, I. Nieto, H. T. Truong, S. P. Kelley, C. J. Herbst-Gervasoni, K. L. Jernigan, M. K. Bowman, S. Pan, M. Zeller, E. T. Papish, Inorg. Chem. 2014, 53, 12689−12698; c) V. K. K. Praneeth, M. Kondo, P. M. Woi, M. Okamura, S. Masaoka, ChemPlusChem 2016, 81, 1123−1128; d) K. J. Fisher, K. L. Materna, B. Q. Mercado, R. H. Crabtree, G. W. Brudvig, ACS Catal. 2017, 7, 3384−3387; e) S. J. Koepke, K. M. Light, P. E. VanNatta, K. M. Wiley, M. T. Kieber-Emmons, J. Am. Chem. Soc. 2017, 139, 8586−8600; f) J. Wang, H. Huang, T. Lu, Chin. J. Chem. 2017, 35, 586−590. [19] a) R. Pogni, M. C. Baratto, E. Busi, R. Basosi, J. Inorg. Biochem. 1999, 73, 157−165; b) I. Gamba, I. Mutikainen, E. Bouwman, J. Reedijk, S. Bonnet, Eur. J. Inorg. Chem. 2013, 115−123. [20] a) Z. Jin, P. Li, D. Xiao, J. Mater. Chem. A 2016, 4, 11228−11233; b) T.-T. Li, Y. Chen, F.-M. Li, W.-L. Zhao, C.-J. Wang, X.-J. Lv, Q.-Q. Xu, W.-F. Fu, Chem. Eur. J. 2014, 20, 8054−8061; c) C. Wang, Y. Chen, W.-F. Fu, Dalton Trans. 2015, 44, 14483−14493; d) E. Pizzolato, M. Natali, B. Posocco, A. M. López, I. Bazzan, M. D. Valentin, P. Galloni, V. Conte, M. Bonchio, F. Scandola, A. Sartorel, Chem. Commun. 2013, 49, 9941−9943. [21] A. J. Bard, L. R Faulkner, Electrochemical methods: Fundamentals and Applications; Wiley: New York, 2001. [22] J. M. Savéant, E. Vianello, Electrochim. Acta 1965, 10, 905−920. [23] A. G. Walden, A. J. M. Miller, Chem. Sci. 2015, 6, 2405−2410. [24] E. S. Rountree, B. D. McCarthy, T. T. Eisenhart, J. L. Dempsey, Inorg. Chem. 2014, 53, 9983−10002. [25] a) X. Yang, M. B. Hall, J. Am. Chem. Soc. 2010, 132, 120−130; b) Z. Chen, J. J. Concepcion, X. Hu, W. Yang, P. G. Hoertz, T. J. Meyer, Proc. Natl. Acad. Sci. U.S.A. 2010, 107, 7225−7229; c) Z. Chen, J. J. Concepcion, H. Luo, J. F. Hull, A. Paul, T. J. Meyer, J. Am. Chem. Soc. 2010, 132, 17670−17673. 22

This article is protected by copyright. All rights reserved.

10.1002/cssc.201701458

ChemSusChem

[26] a) D. J. Wasylenko, C. Ganesamoorthy, J. Borau-Garcia, C. P. Berlinguette, Chem. Commun. 2011, 47, 4249−4251; b) L. Xu, H. Lei, Z. Zhang, Z. Yao, J. Li, Z. Yu, R. Cao, Phys. Chem. Chem. Phys. 2017, 19, 9755−9761. [27] a) V. Mishin, J. P. Gray, D. E. Heck, D. L. Laskin and J. D. Laskin, Free Radical Biology & Medicine 2010, 48, 1485−1491; b) U. S. Akhtar, E. L. Tae, Y. S. Chun, I. C. Hwang and K. B. Yoon, ACS Catal. 2016, 6, 8361−8369. [28] a) T. Pandiyan, H. J. Guadalupe, J. Cruz, S. Bernès, V. M. Ugalde-Salvdivar, I. González, Eur. J. Inorg. Chem. 2008, 3274−3285; b) F. Baig, R. Kant, V. K. Gupta, M. Sarkar, RSC Adv. 2015, 5, 51220−51232. [29] H. Huang, M. Ata, Y. Yoshimoto, Chem. Commun. 2004, 1206−1207.

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Graphic abstract:

Evident enhancement of activity and significant decrease of overpotential for electrochemical water oxidation were simultaneously attained for molecular copper catalysts by tuning the tetradentate N4 ligand from diamine-dipyridine to diimine-dipyridine.

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