Formation and Reactivity of Copper Acetylacetone Bis(Thiosemicarbazone) Complexes

Author Manuscript

Title: Formation and Reactivity of Copper Acetylacetone bis(Thiosemicarbazone) Complexes

Authors: Jessica K. Bilyj; Jeffery R. Harmer; Paul V. Bernhardt, PhD

This is the author manuscript accepted for publication and has undergone full peer review but has not been through the copyediting, typesetting, pagination and proofrea- ding process, which may lead to differences between this version and the Version of Record.

To be cited as: 10.1002/ejic.201801002

Link to VoR: https://doi.org/10.1002/ejic.201801002 Formation and Reactivity of Copper Acetylacetone bis(Thiosemicarbazone) Complexes

Jessica K. Bilyj,[a] Jeffery R. Harmer[b] and Paul V. Bernhardt*[a]

Abstract

The complexation reaction mechanism of acetylacetone bis-thiosemicarbazone ligands (H2acacR) with CuII is explored using a variety of physical methods. The complexes form via a complicated multistep mechanism that is initiated by ring opening of the pyrazoline form of the ligand and leads, ultimately in air, to an oxidised ketone form of the ligand. Tetradentate N2S2 coordinated forms of the intermediate [Cu(acacR)] are stable only under anaerobic conditions. Upon exposure to air these complexes are cleanly oxidized to the ketone complex [Cu(acacRO)] as shown by X-ray crystallography, electrochemistry, UV-Vis and EPR spectroscopy. The behavior of these complexes contrasts with those of closely related bis-dithiocarbazate Schiff bases which stabilize CuIII.

[a] Ms J.K. Bilyj, Prof. P.V. Bernhardt, School of Chemistry and Molecular Biosciences, University of Queensland, Brisbane 4072, Australia. Email: [email protected] [b] Assoc. Prof. J.R. Harmer, Centre for Advanced Imaging, University of Queensland, Brisbane 4072, Australia. 1

Bis(thiosemicarbazones) are versatile tetradentate N2S2 ligands that have been studied in detail and have found application as chelating agents for metals of interest in radiochemistry for imaging and therapy.[1] For complexes of late transition elements such as NiII and CuII they can effectively complete the coordination sphere of the metal through their planar N2S2 donor set.

Bis(thiosemicarbazones) based on acetylacetone (pentane-2,4-dione) are a particularly well known example. In their neutral state they are isolated in the pyrazoline form[2] which must then revert to the linear isomer in order to coordinate as a tetradentate ligand (Scheme 1). These ligands and their complexes are also associated with a range of tautomeric and acid-base equilibria. The central methylene C-atom gains considerable acidity and typically deprotonates on coordination in a similar fashion to coordinated β-ketiminate[3] and acetylacetonate[4] ligands. The N-atoms adjacent to the coordinated imine are also readily deprotonated when the ligand is coordinated.[5] Acetylacetonate bis(thiosemicarbazone) complexes are known where the ligand is present in its monoanionic[5] (a, A) and dianionic[2b] (b, B) forms, even the trianionic coordinated form (c) is possible (Scheme 1).

Scheme 1. Cyclic and acyclic isomers of symmetrical acetylacetone bis(thiosemicarbazones) and the various acid/base and tautomeric forms of their divalent metal complexes. Labile protons are shown in bold type.

This article is protected by copyright. All rights reserved Ligands from this class also exhibit an unusual N2S2 linkage isomerism as shown in Scheme 2, which is associated with either an asymmetric 5-7-4 or symmetric 5-6-5-chelate ring formation. Nickel(II) complexes exhibit both kinds of structures[2b, 5] and we recently reported a study of the structures and mechanism of formation of complexes from this class.[6]

An additional complexity arises from the tendency of these complexes under aerobic conditions to undergo oxidation at the central C-atom to yield a ketone (Scheme 2). In the case of NiII the involvement of an oxidized NiII ligand radical intermediate (not NiIII) was implicated based on electrochemistry and computational chemistry.[6] The asymmetric NiII complex was inert and neither isomerized nor underwent ligand oxidation. Bis(thiosemicarbazone) complexes of PdII,[7] ZnII [8] and GaIII [9] have also been structurally characterized in analogous asymmetric coordination modes of the type shown in Scheme 2.

Scheme 2. Known structures of Ni(II) complexes of acetylacetone bis(thiosemicarboazones).

In contrast, no examples of asymmetrically coordinated CuII complexes of tetradentate bis(thiosemicarbazones), related bis-dithiocarbazate Schiff bases (N2S2), bis(isothiosemicarbazones) (N4) or bis(thiocarbohydrazone) (N2S2) relatives have been reported. Three crystal structures have been reported of copper complexes (Scheme 3, 1-3) that have an unsubstituted central C-atom within the six-membered chelate ring.[2b, 10] Two of these (2 and 3) are rare CuIII complexes stabilized by electronically delocalized trianionic ligands (analogous to structure c in Scheme 1) while compound 1 bears a methylene group within the chelate ring and was isolated as a CuII complex.[2b] There are a very few structurally characterised examples of CuII complexes bearing a 4-membered NS thiosemicarbazone chelate ring (e.g. compound 4)[11] and these are only found in the presence of co-ligands.

Oxidised ketone analogues of related ligands have been structurally characterized as CuII complexes [12] including acetylacetone bis(hydrazone) (N2O2, compound 5) and bis(aminoguanadine) Schiff [13] base (N4, compound 6) ligands.

Scheme 3. Structurally characterised Cu complexes of bis(thiosemicarbazone) and related ligands.

There are likely other examples of inadvertently oxidized ketone bis(thiosemicarbazone) complexes of Cu that have been formed from the parent bis(thiosemicarbazone) but not recognized in the absence of crystallographic characterisation.[14] Oxidation at the central C-atom can be blocked by substituents at this position.[15]

The objectives of this work were to understand the mechanism of formation and properties of acetylacetone bis(thiosemicarbazone) complexes of Cu, and in particular to examine the structural, spectroscopic and electrochemical properties of these complexes and their reaction with oxygen, based on our previous work with their NiII analogs.[6]

This article is protected by copyright. All rights reserved Results and Discussion Copper Complexation The use of the four R-substituted acetylacetonate bis(thiosemicarbazone) ligands (R = methyl, ethyl, allyl and phenyl) complexed with copper acetate has allowed the investigation of this family of copper complexes under anaerobic and aerobic conditions. One goal was to determine if the linkage isomers (symmetric and asymmetric, Scheme 2) identified for the analogous [Ni(acacR)] complexes[6] were also formed in complex with Cu and also to probe the susceptibility toward ligand oxidation at the central C-atom in the presence of dioxygen.

The complexation of CuII under strictly anaerobic conditions enabled the formation of the putative O complexes [Cu(acacR)]. Due to their oxygen sensitivity no attempts were made to isolate crystalline 2 samples of these compounds but solution characterization of the complexes formed in situ by combining equivalent amounts of Cu(OAc)2·H2O and ligand (H2acacR) was achieved. The acetate ions doubly deprotonate the ligand and assist complexation. As a background to this work, the crystal structure of [Cu(acacme)] has been reported[2b] in its unusual methylene tautomeric form (Scheme 3, compound 1), and this remains the only structurally characterised acetylacetone bis(thiosemicarbazone) complex of CuII in the literature. It is particularly interesting that both thioamide NH groups are deprotonated yet the methylene C-atom is not deprotonated. The complex [Cu(acacme)] was crystallised in an inert atmosphere and upon exposure to air appeared to form the ketone analogue [Cu(acacmeO)];[2b] corroborating previous work by others with the same compound.[2a] Crystallographic structural proof of the identity of compounds from the [Cu(acacRO)] family (Scheme 4, right hand side) is lacking until now.

Scheme 4. Relevant structures of symmetrically coordinated CuII complexes in this work.

Structural Characterization Crystals of the oxidised copper bis(thiosemicarbazone) complexes [Cu(acacRO)] were obtained by slow evaporation of DMF solutions in air and at room temperature. X-ray crystal structures of [Cu(acacetO)] and [Cu(acacpheO)] were determined.

Figure 1. ORTEP plots of (top) [Cu(acacetO)] (one of two independent molecules shown) Cu-N3 1.998(3), Cu-N4 2.002(3), Cu-S1 2.252(1) and Cu-S2 2.270(1) Å, N4-Cu-S1 179.5(1); N3-Cu-S2 165.8(1)° and (bottom) [Cu(acacpheO)]·DMF (DMF and phenyl ring disorder not shown); Cu- N3 1.981(2), Cu-N4 1.991(2), Cu-S1 2.2380(9), Cu-S2 2.2412(8) Å, N4-Cu-S1 177.56(8); N3-Cu- S2 175.95(8)°.

The structure of [Cu(acacetO)] (Figure 1) comprises two neutral complexes in the asymmetric unit. The oxidised ketone ligand is dianionic (deprotonated at N2 and N5) which balances the charge of the divalent Cu ion. The complex exhibits an approximate square planar coordination geometry with very distant contacts with the S atoms of an adjacent complex (Cu…S > 2.96 Å). Interestingly these weak intermolecular interactions are responsible for a large out-of-plane distortion of this S- atom of more than 0.5 Å toward the distant Cu ion while the remaining donor atoms and the Cu are coplanar. This is also evident in a contraction of the ideally linear trans N3-Cu-S2 coordinate angle (165.8(1)°) by ~14 degrees relative to the other pair of trans donor atoms (179.5(1)°). This occurs in both independent molecules which form a polymeric structure (Supplementary Information (SI),

The structure of the phenyl analogue [Cu(acacpheO)]·DMF shows a similar symmetrical planar cis-

CuN2S2 coordination environment. In contrast to the structure of [Cu(acacetO)], there are no axial contacts with the Cu centre within 3.16 Å and both S-donors lie within 0.15 Å of the plane defined by the Cu atom and both N-donors. One of the N-phenyl rings is disordered and this is correlated with positional disorder of the DMF solvent molecule.

The Cu-N bond lengths of [Cu(acacetO)] and [Cu(acacpheO)] are within the range seen for square II II planar tetradentate-coordinated thiosemicarbazone complexes of Cu in a cis-Cu N2S2 configuration.[2b, 7, 16] Bidentate N,S ligands from this family can coordinate in either a cis or trans configuration. In a cis configuration, repulsion between adjacent substituents on the imine N-donors [17] distort the N2S2 donor atoms well away from planarity (toward tetrahedral). The trans-CuN2S2 complexes of bidentate thiosemicarbazones and dithiocarbazate Schiff bases are typically planar[18] and their Cu-N and Cu-S bonds are similar to those found here in [Cu(acacetO)] and [Cu(acacpheO)].

Figure 2. Time resolved UV-Vis spectra of the complexation reaction of Cu(OAc)2 (330 μM) with H2acacR (330 μM) in DMF under anaerobic conditions: (A) [Cu(acacme)]; (B) [Cu(acacet)]; (C) [Cu(acacal)] and (D) [Cuacacphe)]. The arrows highlight spectral changes with time. 7

Progress of the Cu complexation reactions with the cyclic H2acacR (R = Me, Et, Al, Ph) ligands was followed with UV-Vis spectroscopy and the full time-dependent spectral data were modelled with Reactlab Kinetics, which generated rate constants and spectra for the initial, intermediate and final complexes.[19] The overall complexation and oxidation reaction (leading ultimately to the ketone complex [Cu(acacRO)]) may be separated into two stages; an initial complexation reaction under anaerobic conditions (generating [Cu(acacR)]) followed by a ligand oxidation process under aerobic conditions to yield [Cu(acacRO)] (Scheme 4).

Upon combining Cu(OAc)2 with the cyclic pyrazoline ligand (H2acacR, Scheme 1) under anaerobic conditions, an initial complex occurs during mixing, as evident by an immediate colour change (broad absorption maximum ~600-800 nm, Figure 2); formation of this intermediate was too fast to be measured and this species was taken as the nominal starting complex. Two subsequent slower first order steps were monitored by UV-Vis spectroscopy; one with a half-life of ~5-25 min (k1) followed by a second somewhat slower reaction (k2). The reaction was considered complete after about 5 h when no significant variations to the UV-Vis spectrum were apparent. The spectra of the initial complex, an intermediate and the final product are given in Figure 3.

Scheme 5. Proposed mechanism for Cu complex formation during the anaerobic period of the reaction. The rate constants appear in Table 1.

The final spectra of the CuII complexes formed under anaerobic conditions ([Cu(acacR)], Figure 3) were all similar and featured a pair of prominent maxima around 550-600 and 750-800 nm. The intensities of these bands (ε ~ 2000 M-1cm-1) suggest that they are of (ligand-to-metal) charge transfer origin. These spectra match the published electronic spectral data for structurally 8

This article is protected by copyright. All rights reserved characterised [Cu(acacme)].[2b] The exact tautomeric form of [Cu(acacR)] in solution is not known and even an equilibrium between the two forms shown in Scheme 5 is possible. This is discussed in more detail in light of the EPR spectroscopy results.

Given the lability of CuII, the two slow steps seen in the anaerobic complexation cannot be limited by ligand exchange. Pyrazoline ring opening (C-N bond breaking) is a precondition for tetradentate coordination (Scheme 1) so this seems likely to be the first phase (k1 in Scheme 5). The nature of the other intramolecular reaction is less certain. The ligand bears two imine C=N groups that may adopt E or Z isomers but only one conformation will allow tetradentate N2S2 coordination. On this basis, the second step (with rate constant k2) is assigned tentatively to Z-E isomerisation of the C=N

(imine) groups (k2) to enable tetradentate N2S2 coordination (Scheme 5). The paramagnetic nature of the complexes studied here precludes NMR analysis which might otherwise shed more light on this mechanism. Indeed NMR was most useful in studying the related diamagnetic NiII systems.[6]

Figure 3. Reactlab Kinetics calculated spectra for the initial (CuL’), intermediate (CuL”) and final [Cu(acacR)] complexes observed during the complexation reaction; a) [Cu(acacme)], b) [Cu(acacet)], c) [Cu(acacal)], [Cu(acacphe)].

Complexation (anaerobic) Oxidation (aerobic) -1 -1 -1 -1 Complex k1 (s ) k2 (s ) k3 (s ) k4 (s ) [Cu(acacme)] 3.07(3) × 10-3 2.67(1) × 10-4 8.50(4) × 10-5 1.69(2) × 10-5 [Cu(acacet)] 4.682(9) × 10-4 4.8(1) × 10-5 1.03(1) × 10-4 6.61(6) × 10-5 [Cu(acacal)] 9.06(2) × 10-4 3.56(4) × 10-4 5.20(4) × 10-5 1.73(8) × 10-5 [Cu(acacphe)] 3.976(4) × 10-3 2.81(2) × 10-3 1.342(9) × 10-4 1.440(5) × 10-5

Figure 4. Time resolved UV-Vis spectra of the aerobic oxidation reaction of (a) [Cu(acacme)]; (b) [Cu(acacet)]; (c) [Cu(acacal)] and (d) [Cuacacphe)]. Arrows highlight the important spectral changes with time.

Oxidation Kinetics

The anaerobic solutions of [Cu(acacR)] were then exposed to oxygen (the aerobic period) and monitored spectrophotometrically for an additional 20 h (Figure 4). The initial electronic absorption maxima in the regions 550-600 and 750-800 nm, characteristic of [Cu(acacR)], vanished and were replaced ultimately by a spectrum comprising a single maximum around 500 nm (ε ~ 1000-2000 M-

This article is protected by copyright. All rights reserved -1cm-1). However, an intermediate species was also identified (orange curves in Figure 5) with enhanced absorption around 400 nm so the overall aerobic reaction comprised two first order processes (k3 and k4 in Table 1). In all cases the Reactlab Kinetics calculated final UV-Vis spectra (Figure 5) were consistent with that of [Cu(acacRO)] by comparison with an isolated sample of crystalline [Cu(acacalO)] (SI, Fig. S2). They also match the published UV-Vis spectra of the compound assigned as [Cu(acacmeO)] and prepared under aerobic conditions.[2]

Figure 5. Reactlab Kinetics calculated spectra for the initial ([Cu(acacR)]), intermediate ([Cu(acacR)]) and final ([Cu(acacRO)]) complexes observed during the aerobic ligand oxidation reaction: starting complexes (a) [Cu(acacme)]; (b) [Cu(acacet)]; (c) [Cu(acacal)] and (d) [Cu(acacphe)].

EPR Spectroscopy

The d9 complexes [Cu(acacR)] and [Cu(acacRO)] were investigated by EPR spectroscopy as frozen DMF solutions. The spectrum of [Cu(acacphe)], as a representative example, prepared by stoichiometric combination of Cu(OAc)2 and H2acacphe under anaerobic conditions, is shown in the supporting information (Fig. S3). It is apparent that the spectrum of neutral [Cu(acacphe)] is a composite of two species (labelled A and B where the hyperfine lines in the parallel region are indicated). Species A exhibits a lower g|| (~2.11) and larger A|| values (~207 G) than species B (g|| ~

This article is protected by copyright. All rights reserved 2.13 and A|| ~176 G). Due to the complexity of the spectrum, overlap and the low signal to noise ratio, accurate spin Hamiltonian parameters were not pursued for the two components.

The EPR spectrum clarified when the same CuII complex was formed in the presence of three equivalents of Et3N, which indicates that the mixture of species seen in Figure S3 comprises different tautomeric or acid/base forms of the complex. The spectrum of the fully deprotonated [Cu(acacphe-H)]- shows axial symmetry and is consistent with an essentially square planar coordination geometry (Figure 6). Deprotonation of the six-membered chelate ring necessarily leads to a delocalized electronic structure (structure c in Scheme 1).

The highly symmetrical (axial) EPR spectrum of [Cu(acacphe-H)]- and its resemblance to species A in Figure S3, suggests that species A is the neutral conjugate acid [Cu(acacphe)] in the same delocalized tautomeric form. The spin Hamiltonian parameters for [Cu(acacphe-H)]- are essentially the same as those reported for the CuII analogue of compound 2 in Scheme 3.[10a] By elimination this leaves species B as the alternative localized tautomer where the methylene group is present

(Scheme 1). The spectra are qualitatively similar and consistent with axial symmetry gz > gx = gy and a dx2-y2 ground state. This is a new finding in the chemistry of these complexes i.e. that their solution structures do not necessarily match what has been identified in the solid state and that tautomeric equilibria may be present (see Scheme 5).[2b]

In contrast with the NiII chemistry, for CuII we can conclusively rule out an asymmetric (4-7-5 chelate ring) binding mode for the ligand as angular distortions due to the enforced 4-7-5 chelate ring formation would lead to a spectrum with rhombic symmetry (gz > gy > gx).

Figure 6. Experimental (black) and simulated (red) X-band (9.3711 GHz) EPR spectra of - [Cu(acacphe-H)] (1 mM) prepared in situ under anaerobic conditions with 3 equivalents of Et3N -4 -1 (130 K DMF). Spin Hamiltonian parameters: gx 2.0240 (Ax 34.9 × 10 cm ), gy 2.0259 (Ay 35.3 × -4 -1 -4 -1 -4 -1 -4 -1 10 cm G), gz 2.1049 (Az 185.8 × 10 cm ); AN,x 14.9 × 10 cm , AN,y 15.9 × 10 cm , AN,z 14.3 × 10-4 cm-1.

Upon exposure of a DMF solution of [Cu(acacal)] (as a representative example) to oxygen at room temperature the EPR signature of [Cu(acacal)] vanishes and is replaced by a single species over the course of a few hours (Figure 7, black curve). This final spectrum matches that of crystalline [Cu(acacalO)] (Figure 7, blue curve). Simulation of the spectrum (Figure 7, red curve) gave spin Hamiltonian parameters that match the previous work of Davies et al. on the product obtained from [2a] aerial oxidation of [Cu(acacet)] i.e. [Cu(acacetO)]. In passing, the increase in gz and decrease in

Az in the spectrum of [Cu(acacRO)] (from [Cu(acacR)]) upon oxidation are both reminiscent of species B in the spectra of the initial mixture of tautomers (supporting information Fig. S3) adding further weight to the assignment of the mixture seen in this solution as being due to a tautomeric

This article is protected by copyright. All rights reserved equilibrium; the structural similarity between [Cu(acacalO)] and the electronically localized tautomer of [Cu(acacal)] (species B) would imply a similar EPR spectrum.

Figure 7. X-band EPR spectra of (black) [Cu(acacal)] (9.301 GHz, 130 K) exposed to air for 2 d, (blue) crystalline [Cu(acacalO)] in DMF (9.374 GHz, 130 K) and (red) simulation with the spin -4 -1 -4 -1 Hamiltonian parameters: gx 2.0259 (Ax 26.6 × 10 cm ), gy 2.0356 (Ay 29.1 × 10 cm ), gz 2.1330 -4 -1 -4 -1 -4 -1 -4 -1 (Az 176.4 × 10 cm ); AN,x 11.0 × 10 cm , AN,y 14.1 × 10 cm , AN,z 12.3 × 10 cm . Solutions were 1-2 mM in DMF.

Cyclic Voltammetry Cyclic voltammetry of the anaerobically prepared [Cu(acacR-H)]- complexes was performed in

DMF (Figure 8) with three equivalents of Et3N added to ensure full deprotonation. Without excess

Figure 8. Cyclic voltammetry of 1.2 mM solutions of [CuII(acacR-H)]- in DMF. Sweep rate 200 -1 mV s . Solutions prepared anaerobically by addition of 3 equivalents of Et3N to each complex to ensure full deprotonation.

The main feature in each case was a quasi-reversible CuII/I redox response in the range -1000 to - 600 mV vs Fc+/0. The redox potential varied according to the inductive effect of the N-substituent with the electron withdrawing phenyl analogue appearing at the highest potential. Scanning to higher potentials identified poorly defined, irreversible oxidation processes around +300 mV vs Fc+/0.

Figure 9. Cyclic voltammetry of 1.2 mM solutions of [CuII(acacRO)] in DMF formed after aerobic oxidation. Sweep rate 200 mV s-1. The responses at -600 and -50 mV in the CV of [Cu(acacpheO)] are due to minor impurities.

Electrochemistry of the fully oxidized DMF solutions of [Cu(acacRO)] (Figure 9) gave a significantly different outcome. Two reversible one-electron redox couples were identified that 15

This article is protected by copyright. All rights reserved were separated by almost 1.5 V. Crystals of [Cu(acacetO)] characterised crystallography (Figure 1) dissolved in DMF and analysed by CV (supporting information Fig. S4) presented redox couples at 190 mV and -1066 mV, which match the CV in Figure 9. The nature of these couples is investigated in greater detail through spectroelectrochemistry below.

Spectroelectrochemistry

Spectroelectrochemistry of the four [Cu(acacR-H)]- complexes was conducted in situ under anaerobic conditions in DMF. Due to the possibility of different tautomers being present in solution (see EPR section), three equivalents of triethylamine were again added to ensure formation of a single species ([Cu(acacR-H)]-). Spectra were acquired at 50 mV intervals in the vicinity of the reversible redox responses determined from cyclic voltammetry (Figures 8 and 9). All potential dependent spectra are shown in the supporting information (Fig. S5). Data modelling using the Nernst and Beer-Lambert equations on the entire potential dependent spectra was undertaken with Reactlab Redox[20] which generated the spectra of the fully oxidized and fully reduced species (Figure 10) along with the redox potential, which was found to align with the redox potential obtained from the cyclic voltammetry (Figure 8).

In all cases the fully resting form of the [Cu(acacR)] complex exhibited maxima around 780 and 550 nm. These spectra match those of the species formed after complexation of the free ligand with II Cu(OAc)2 under anaerobic conditions (Figure 4) and EPR spectroscopy showed this to be a Cu complex. Upon electrochemical reduction, both of these electronic absorption maxima in the visible region vanish and the spectrum of the putative CuI form is featureless except for a maximum around 400 nm that is also present in the CuII complex (Figure 10).

Figure 10. Spectroelectrochemical analysis of the [Cu(acacR-H)]- complexes; a) [Cu(acacme-H)]-, b) [Cu(acacet-H)]-, c) [Cu(acacal-H)]-, d) [Cu(acacphe-H)]-. The spectra of the fully oxidized CuII (red) and reduced CuI (purple) forms are shown.

Spectroelectrochemistry of the corresponding [Cu(acacRO)] solutions can also be understood by comparison with the time-resolved UV-Vis data in Figures 4 and 5 and the EPR data. The complete data are included in the supporting information (Figs S6-S8). Spectra measured in the vicinity of the lower potential redox couple (see Figure 11) revealed that the oxidised form of this couple matched the species formed after aerobic oxidation of the complex i.e. the CuII complex [Cu(acacRO)] with a single visible absorption maximum around 500 nm (supporting information Fig. S2). Upon reduction a broad band appeared around 820 nm but there was little change to the rest of the spectrum; the invariance of the maximum around 500 nm is of note. The origin of the maximum above 800 nm is unknown and is in distinct contrast with the featureless spectra of reduced [CuI(acacR-H)]2- (Figure 10). The electronic structure of reduced [Cu(acacRO)]- remains unknown which may be either a CuII-ligand-centred radical or a CuI complex but our data cannot discriminate between these possibilities. Other studies have shown that ligands of this type may be redox non- innocent.[21]

Spectroelectrochemical oxidation of crystalline [Cu(acacetO)] (supporting information Fig. S8) in the region of the higher potential redox couple (200-300 mV vs Fc+/0) led to a reversible spectral 17

This article is protected by copyright. All rights reserved change but the spectral changes were minor. Again two possibilities emerge for the product of oxidation; a CuIII complex or a CuII ligand centred radical and our data do not discriminate the two. Stabilisation of CuIII in the related dithiocarbazate Schiff base 2 (Scheme 3) has been established crystallographically and spectroscopically and the CuIII/II redox potential was -126 mV vs Fc+/0, which is about 300-400 mV lower than the high potential couple in [Cu(acacetO)] (Figure 9). This shows that the adjacent terminal substituents (NHR or SR) on bis(thiosemicarbazone) or dithiocarbazate Schiff base ligands have distinctly different capacities to stabilize higher oxidation states. Similar experiments could not be pursued for the complexes prepared in situ due to the presence of excess Et3N which is electroactive in this region.

Figure 11. Spectroelectrochemical analysis of the [Cu(acacRO)]0/- complexes; a) [Cu(acacmeO)], b) [Cu(acacetO)], c) [Cu(acacalO)], d) [Cu(acacpheO)]. The spectra of the fully oxidised (red) and reduced (purple) forms calculated with Reactlab Redox are shown.

Conclusions Complexation of the cyclic pyrazoline forms of acetylacetone bis-thiosemicarbazone ligands with CuII is a complicated process that proceeds via a partially coordinated intermediate then to a tetradentate N2S2 form [Cu(acacR)] that is stable under anaerobic conditions but exists in two different tautomeric forms. Upon exposure to air oxidation to the ketone complex [Cu(acacRO)] occurs cleanly as shown by X-ray crystallography, electrochemistry, UV-Vis and EPR 18

This article is protected by copyright. All rights reserved spectroscopy. The mechanism of this oxidation remains unknown. The sensitivity of these complexes to ligand oxidation contrasts with those of closely related bis-dithiocarbazate Schiff bases which stabilise CuIII.[10a]

By comparison with the previously studied NiII systems,[6] there was no evidence here of the formation of asymmetrically coordinated complexes with CuII (Figure 2) at any stage. It has been shown before that these asymmetric NiII thiosemicarbazone complexes are inert and do not undergo ligand oxidation to the ketone analog. In this case all CuII complexes of the form [Cu(acacR)] are very air sensitive and convert cleanly to the ketone complex [Cu(acacRO)] over the course of a few hours in air.

This study has shown that the copper coordination chemistry of acetylacetone bis- thiosemicarbazone N2S2 ligands present several challenges and potential applications of these ligands in areas such as Cu-radioimaging need to take these factors into consideration. The sluggish complexation kinetics are a consequence of significant ligand rearrangement originating from the cyclic pyrazoline. Furthermore the ensuing CuII complexes are unstable in air and are quantitatively oxidized to their ketone analogues. Electrochemistry and EPR spectroscopy of the CuII complexes has shown that their structures and redox properties are significantly altered by this ligand oxidation process. Also the redox non-innocence of these ligands in complex with Cu has been shown by optical spectroelectrochemistry.

Experimental Section Synthesis

All reagents were obtained commercially and used as received. The ligands H2acacme, H2acacet,

H2acacal and H2acacphe were synthesized as previously reported and are isolated in their ring- closed pyrazoline forms (Scheme 1).[6]

[Cu(acacR)] (R = Me, Et, Allyl, Ph) (prepared in situ but not isolated)

In an anaerobic glovebox (Belle Technology, O2 < 20 ppm), solutions of Cu(OAc)2·H2O (240 L,

250 mM) and ligand H2acacR (5 mL, 1.2 mM) in DMF were mixed and allowed to equilibrate for 24 h to give approximately 1.14 mM solutions of [Cu(acacR)] which were stable in the absence of oxygen.

[Cu(acacRO)] (R = Me, Et, Allyl, Ph) Reactions were typically carried out on a ~500 μmol scale. The free bis-thiosemicarbazone ligand

H2acacR was combined with an equimolar amount of Cu(OAc)2.H2O dissolved in ethanol (30 mL) and stirred at 70°C for 1 h. The mixture was left to cool and the crude product was filtered off and

This article is protected by copyright. All rights reserved washed with cold ethanol. The product was recrystallized by dissolving the solid in a minimum amount of DMF then allowing the solution to slowly evaporate on a watch glass to afford dark red crystals of the corresponding ligand-oxidized compounds [Cu(acacRO)]. The crystals served only as material for X-ray analysis or as standards to identify species formed in situ and studied by UV- Vis, EPR or cyclic voltammetry.

Physical Methods Cyclic Voltammetry Cyclic voltammetry (CV) was performed with a BAS100B/W potentiostat using a glassy carbon working electrode, platinum auxiliary electrode and a non-aqueous Ag/Ag+ reference electrode in DMF. Ferrocene was used as an external standard and all potentials are cited versus Fc+/0. Approximately 1.2 mM solutions of [Cu(acacR)] (R = Me, Et, Al, Ph) were prepared in an anaerobic glovebox as described above except the dry, distilled DMF solvent also contained 0.1 M tetraethylammonium perchlorate (Et4NClO4) as supporting electrolyte. To eliminate any alternative protonated forms (see Scheme 1), 3 molar equivalents of Et3N were added to the solutions before measurement to ensure complete deprotonation to the anion [Cu(acacR-H)]-.

Optical Spectroelectrochemistry Experiments were conducted with a Pine Instruments quartz spectroelectrochemical cell (1.7 mm optical path length) using a Pt ‘honeycomb’ working electrode, a Pt auxiliary electrode and a non- aqueous Ag/Ag+ reference electrode in DMF. The solutions for analysis were prepared by diluting

0.2 mL of the above ~ 1.2 mM voltammetry solution with 0.4 mL of a 0.1 M Et4NClO4 DMF solution to make a ~400 M solution which was suitable for spectroelectrochemistry. Spectra were acquired within the anaerobic glovebox with an Ocean Optics USB2000 fibre optic spectrophotometer and a DT-MINI-2-GS miniature deuterium/tungsten/halogen UV-Vis-NIR light source. Potentials were set with a BAS100B/W potentiostat operating in chronocoulometry mode and UV-Vis spectra were taken when dynamic equilibrium was established and all absorbance changes ceased (usually within 5 min). Spectra were taken at potentials below and above the redox potentials determined by CV (above) and reversibility was established in each case by stepping the potential sequentially in positive and negative directions.

Time Resolved UV-Vis Spectroscopy Approximately 300 M anaerobic solutions of [Cu(acacR)] were prepared by mixing stoichiometric amounts of Cu(OAc)2·H2O and H2acacR in a glovebox as described above (Synthesis section). Upon mixing, the solutions were immediately transferred to a quartz cuvette and sealed with a Teflon stopper coated with silicone grease. The cuvettes were promptly removed from the glovebox

This article is protected by copyright. All rights reserved and UV-Vis spectra were acquired with an Agilent 8453 UV-Visible spectrophotometer equipped with a multi-cell holder (enabling parallel measurements of several samples). Spectra were measured every 30 s for the first 4 h and then the time intervals increased by 10% sequentially for the remaining reaction (~2d). Solutions were maintained at 25°C by a Huber Ministat recirculating bath.

Measurements under aerobic conditions were performed similarly but starting with the anaerobic equilibrated [Cu(acacR)] solution in the cuvette (above). The Teflon stopper was removed which initiated the reaction with oxygen (no mixing was necessary). Spectra were measured every 30 s for the first 3h then again in 10% increasing intervals for the remaining 17 h.

EPR Spectroscopy Continuous-wave (CW) X-band (ca. 9 GHz) electron paramagnetic resonance spectra of all compounds were acquired as frozen 1 mM solutions of compound in DMF using a Bruker Elexsys

E540 spectrometer equipped with an ElexSys Super High Sensitivity Probehead and liquid N2 cooling. The magnetic field was calibrated with 2,2-diphenyl-1-picrylhydrazyl (g = 2.0036) and measurements were carried out at 130 K using a modulation amplitude of 0.2 mT and a modulation frequency of 100 kHz. Solutions of [Cu(acacR)] were prepared anaerobically as described above (Synthesis section) then transferred to an EPR tube and sealed with a rubber septum and frozen immediately. An additional spectrum of [Cu(acacphe)] was acquired in its neutral state and also - after the addition of 3 equivalents of Et3N to give the anion [Cu(acacphe-H)] (see Scheme 1). After measurement, all samples were thawed then exposed to air for 2 h which demonstrated oxidation of the [Cu(acacR)] parent to its ketone analogue [Cu(acacRO)]. A solution of crystalline [Cu(acacalO)] in DMF (~0.5 mg in 300μL) was used as a standard.

Crystallography

Crystallographic data were collected on an Oxford Diffraction Gemini CCD X-ray diffractometer using Cu-Kα (1.54184 Å) radiation. The samples were cooled to 190 K with an Oxford Cryosystems Desktop Cooler. The structures were solved with SHELXS and refined with SHELXL[22] within the WinGX package.[23] The thermal ellipsoid diagrams were produced with ORTEP3[24] and rendered with POVRay. Crystallographic data in CIF format have been deposited with the Cambridge Crystallographic Data Centre with numbers 1855931 ([Cu(acacetO)] and 1855932 ([Cu(acacpheO)]·DMF).

Acknowledgements We gratefully acknowledge financial support from the Australian Research Council and the Uinversity of Queensland. 21

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