Reversible dioxygen binding and arene hydroxylation reactions: Kinetic and thermodynamic studies involving ligand electronic and structural variations

Abstract

Copper-dioxygen interactions are of intrinsic importance in a wide range of biological and industrial processes. Here, we present detailed kinetic/thermodynamic studies on the O₂-binding and arene hydroxylation reactions of a series of xylyl-bridged binuclear copper(I) complexes, where the effects of ligand electronic and structural elements on these reactions are investigated. Ligand 4-pyridyl substituents influence the reversible formation of side-on bound μ-η²:η²-peroxodicopper(II) complexes, with stronger donors leading to more rapid formation and greater thermodynamic stability of product complexes [Cu^II₂(^RXYL)(O₂²⁻)]²⁺. An interaction of the latter with the xylyl π-system is indicated. Subsequent peroxo electrophilic attack on the arene leads to C–H activation and oxygenation with hydroxylated products [Cu^II₂(^RXYLO²⁻)(⁻OH)]²⁺ being formed. A related unsymmetrical binucleating ligand was also employed. Its corresponding O₂-adduct [Cu^II₂(UN)(O₂²⁻)]²⁺ is more stable, but primarily because the subsequent decay by hydroxylation is in a relative sense slower. The study emphasizes how ligand electronic effects can and do influence and tune copper(I)–dioxygen complex formation and subsequent reactivity.

1. Introduction

Interests in copper(I)-dioxygen reactivity of coordination complexes derive from the need to establish fundamental properties and due to the occurrence of such chemistry in a variety of copper proteins which are essential in O₂-processing in aerobic organisms [1]. Notable examples of the latter include hemocyanins [2] (dioxygen carriers in arthropods and mollusks) and enzymes involved in O₂-activation towards substrate oxidations such as tyrosinase (Tyr) [3], dopamine β-monooxygenase (DβM) [4], peptidyl α-amidating monooxygenase (PHM) [4] and particulate (membrane bound) copper-dependent methane monooxygenase (pMMO) [1c,5]. The monooxygenase Tyr employs a coupled binuclear copper site similar to that in hemocyanin, where the reduced dicopper(I) site reacts with O₂ giving a side-on (or μ-η²:η²) peroxodicopper(II) species capable of phenol oxygenation, producing o-catechol or o-quinone directly.

With aromatic ring hydroxylation via Cu^I/O₂ chemistry as a goal in Tyr modeling, we earlier on discovered and detailed the chemistry with a binuclear copper(I) complex [Cu^I₂(H-XYL)]²⁺ which quantitatively reacts with resulting in the hydroxylation of the 2-position of the arene spacer (Scheme 1) [6]. A Cu₂O₂ reactive intermediate, [Cu₂(XYL)(O₂²⁻)]²⁺ (Scheme 1) with a side-on bridging peroxide moiety similar to that for oxy-hemocyanin and oxy-tyrosinase was proven based on its UV–Vis spectroscopic features (λ_max = 360, 435 nm) and resonance Raman (rR) spectroscopic properties (for R = NO₂) [6b,c]. The reaction mechanism deduced for this hydroxylation chemistry involves an electrophilic attack of the arene by the side-on peroxo moiety. Theoretical studies supported this conclusion [6b], however this was first shown by substituent effect studies where for R = MeO, t-Bu, F, CN, and NO₂ at the 5-position of the arene (Scheme 1), the rate constants (k₂) for arene hydroxylation increased with R electron-donating ability[6c,7]. Further support for the electrophilic mechanism came from studies where a 2-Me substituted xylyl ligand was employed, as a hydroxylation-induced methyl-migration (i.e. “N.I.H.” shift) occurred [8]. As an interesting aside, a substituent effect on the dioxygen binding to [Cu^I₂(R-XYL)]²⁺ where electron donating R-groups increased the copper-dioxygen binding strength was rationalized by the suggestion that an interaction between the xylyl π-system and the peroxo group occurs [6b,7].

Scheme 1.

Dicopper mediated hydroxylation of arenes as copper monooxygenase models systems.

While a good number of other binuclear copper(I) complexes with m-xylyl linker were since shown to also exhibit arene hydroxylation chemistry [9], considerable attention has also focused on dicopper systems effecting phenol o-hydroxylation. In this context, elegant studies by Itoh and coworkers [3b] revealed that such chemistry occurs efficiently, not only with certain pyridylalkylamine chelates in synthetic models, but also with the enzyme Tyr itself. In both cases, the reactive intermediate is a [Cu^II₂ (μ-η²:η²-O₂²⁻)]²⁺ species which attacks the phenol substrate as an electrophile in a manner very similar to that of our [Cu₂(XYL)]²⁺ system (Scheme 1). Yet, the Tolman and Stack research groups showed that Cu^II₂(μ-η²:η²-O₂²⁻) are often in equilibrium with an isomeric bis-μ-oxo dicopper(III) [Cu^III₂(O)₂]²⁺ species (Scheme 2) [1d,e,10] and it has been shown that these groups have also shown the latter to be capable of arene and o-phenol hydroxylation[1d,e,9b,10b]. As a relevant aside for the studies described above for the oxygenation of [Cu^I₂(R-XYL)]²⁺ complexes, the side-on per-oxo complex intermediate was readily detected and characterized, while no hint of a related bis-μ-oxo species [Cu^III₂(NO₂-XYL)(O)₂]²⁺ was found experimentally; this and theoretical studies support only the side-on peroxo-dicopper(II) species as the reacting electrophile in the Cu₂(XYL) hydroxylation system [6b]. Computational studies on Tyr itself lead to a variety of conclusions. Thus, while a bis-μ-oxo-dicopper(III) entity has never been detected in Tyr, the chemical model studies leave the possibility of an unobservable enzyme [Cu^II₂(μ-η²:η²-O₂²⁻)]²⁺ to [Cu^III₂(O)₂]²⁺ conversion induced by phenol/phenolate substrate copper coordinations[10b,11].

Scheme 2.

Structures of well established dicopper O₂-adducts.

To further elucidate fundamental aspects of the XYL type dicopper(I)/O₂ chemistry involving both O₂-complex formation and substrate hydroxylation, we have here extended our kineticthermodynamic studies on the oxygenation reactions of a related but new series of complexes focusing on the influence of ligand electronic and structural elements. Investigations of variations in ligand electron-donating ability in Cu^I/O₂ chemistry have in the past been extremely revealing [7,12], also see the discussion below. Here, this is achieved by incorporating various 4-pyridyl substituents into the XYL ligands. An unsymmetrical ligand analogue UN (Fig. 1) is utilized to investigate the effect of ligand constraints on the O₂-interaction of the resulting dicopper(I) complex.

Fig. 1.

Ligands used in this study of copper–dioxygen kinetics.

2. Experimental

2.1. Material and methods

Reagents and solvents used were of commercially available reagent quality unless otherwise stated. Methylene chloride and diethyl ether were purified by passing through an alumina column under Ar. Preparation and handling of air-sensitive materials were carried out under an argon atmosphere by using standard Schlenk techniques. Deoxygenation of solvents and solutions was achieved by bubbling with Ar for ~1/2 h or by three freeze/pump/thaw cycles. Solid samples were stored and transferred in an MBraun drybox filled with prepurified nitrogen. Samples for IR, NMR and UV spectra were also prepared in the drybox.

¹H NMR spectra were recorded in CD₃NO₂, or CDCl₃ on a Bruker AMX-300 NMR spectrometer. Chemical shifts are reported as d values downfield from an internal TMS reference. Elemental analyses were performed by Desert Analytics, Tuczon, AZ. EI/CI and FAB Mass spectra were obtained on a VG Instruments 70S Gas Chromatography/Mass Spectrometer by Dr. J. Kachinzki, Mass Spectrometry Facility, Department of Chemistry, the Johns Hopkins University. Infrared spectra were recorded as Nujol mulls on a Mattson Galaxy 4030 series FT-IR spectrophotometer and calibrated with a polystyrene film. UV–Vis spectra were collected on Hewlett-Packard Model 8453 diode array spectrophotometer with HP Chemstation software.

2.2. Syntheses of ligands and copper(I) complexes

2.2.1. ^MeOXYL

To 4-methoxy-2-vinylpyridine (1.6 g, 0.012 mol) in a 15 mL pressure tube (Aldrich) was added m-xylenediamine (0.35 g, 0.0025 mol) and MeOH (2 mL) along with five drops of glacial acetic acid. The resulting mixture was stirred at 60 °C for 5 days. After cooling to room temperature, the volatile components were removed by rotary evaporation. The residual acetic acid was neutralized with Na₂CO₃ aqueous solution. The resulting solution was extracted with CH₂Cl₂ three times. The combined organic extracts were dried over Na₂SO₄ and the volatile components were removed by rotary evaporation. Purification of the crude product on alumina gel using 5% MeOH-95% EtOAc gave 1.1 g (64%) of yellow oil. R_f = 0.4 (silica, 10% MeOH-90% EtOAc). ¹H NMR (300 MHz, CDCl₃): d 8.27 (4H, d, J = 6.0 Hz), 7.16-7.07 (4H, m), 6.63-6.50 (8H, m), 3.76 (12H, s), 3.67 (4H, s), 2.93 (16H, m). FAB-MS (m/z): 677 (M+H⁺).

2.2.2. ^Me2NXYL

To 4-dimethylamino-2-vinylpyridine (3.2 g, 0.022 mol) in a 15 mL pressure tube (Aldrich) were added m-xylenediamine (0.70 g, 0.0051 mol) and MeOH (10 mL) along with glacial acetic acid (0.63 g, 0.011 mol). The resulting mixture was stirred at 60 °C for 5 days. After cooling to room temperature, the volatile components were removed by rotary evaporation. The residual acetic acid was neutralized with Na₂CO₃ aqueous solution. The resulting solution was extracted with CH₂Cl₂ three times. The combined organic extracts were dried over Na₂SO₄ and the volatile components were removed by rotary evaporation. Purification of the crude product on alumina gel using 3% MeOH-97% EtOAc gave 1.0 g (27%) of yellow oil. R_f = 0.19 (silica, 3% NH₄OH, 97% MeOH). ¹H NMR (300 MHz, CDCl₃): d 8.11 (4H, d, J = 6.0 Hz), 7.23-7.12 (4H, m), 6.31-6.29 (8H, m), 3.71 (4H, s), 2.93-2.82 (40H, m). FAB-MS (m/z): 729.6 (M+H⁺).

2.2.3. [Cu^I₂(^MeOXYL)](PF₆)₂ (1b)

Under an Ar atmosphere, 0.297 g (0.439 mmol) of ^MeOXYL in 20 mL of CH₂Cl₂ was deaerated for 20 min and then added to 0.320 g (0.859 mmol) of [Cu(CH₃CN)₄](PF₆)²⁸ in a 100 mL Schlenk flask. The resulting solution was stirred for ½ h and then cooled to −78 °C. Addition of 100 mL of deaerated diethyl ether precipitated out light yellow powder. After warming to room temperature, the solvent was decanted and the solid was washed with CH₂Cl₂/diethyl ether (v:v, 1:5) for five times. The solid was dried under vacuum to afford 0.25 g (27%) of light yellow powder. ¹H NMR (300 MHz, CD₃NO₂): d 8.30 (4H, d), 7.43 (1H, s), 7.09-7.22 (3H, m), 6.85-6.90 (8H, m), 3.92 (12H, s), 3.72 (4H, s), 3.11(16H, s,br). Anal. Calc. for C₄₀H₄₈N₆O₄F₁₂P₂Cu₂: C, 43.92; H, 4.42; N, 7.68. Found: C, 43.92; H, 4.61; N: 7.97%.

2.2.4. [Cu^I₂(^Me2NXYL)](PF₆)₂·1.1CH₂Cl₂ (1c)

Under an Ar atmosphere, 0.213 g (0.292 mmol) of ^NMe2XYL in 15 mL of MeOH was deaerated for 20 min and then added to 0.210 g (0.563 mmol) of [Cu(CH₃CN)₄](PF₆) in a 100 mL Schlenk flask at 0 °C. The resulting mixture was stirred for 20 min and 70 mL of deaerated diethyl ether was added to precipitate out yellow powder. After decanting the solvent, the solid was redissolved in 20 mL of deaerated CH₂Cl₂ at 0 °C and layered with 80 mL of diethyl ether to afford a yellow precipitate. The solvent was decanted and the solid was dried under vacuum for 2 h to yield 140 mg (44%) of yellow powder. ¹H NMR (300 MHz, CD₃NO₂): d 8.03 (4H, d), 7.48 (1H, s), 7.11 (3H, m), 6.5 (8H, s), 5.32 (2.2H, s), 3.57 (4H, s), 3.11-3.04 (40H, m). Anal. Calc. for C₄₄H₆₀N₁₀F₁₂P₂Cu₂·1.1CH₂Cl₂: C, 43.5; H, 5.04; N, 11.22. Found: C, 43.2; H, 5.14; N, 11.22%.

2.2.5. [Cu^II₂(^MeOXYLO⁻)(OH⁻)](PF₆)₂·0.5CH₂Cl₂ (3b)

In a 100 mL Schlenk flask, [Cu^I₂(^MeOXYL)](PF₆)₂ (100 mg, 0.0914 mmol) was dissolved in 50 mL deaerated DMF under an Ar atmosphere. The resulting yellow solution was cooled to 0 °C and exposed to an atmosphere of dry oxygen overnight. The solvent was removed on a high-vacuum rotary evaporator and the green residue obtained was recrystallized from CH₂Cl₂/ether at 0 °C, giving 95 mg (48%) of green crystalline solid. Anal. Calc. for C₄₀H₄₈N₆Cu₂O₆P₂F₁₂·0.5CH₂Cl₂: C, 41.64; H, 4.23; N, 7.19. Found: C, 41.81; H, 4.15; N, 7.20%. μ_eff (Evans method, CD₂Cl₂): 1.08 ± 0.10 B.M. UV–Vis (nm, CH₂Cl₂): 370 (ε = 3450), 637 (ε = 150). IR (nujol, cm ⁻¹): 3643 (br, OH), 840 (s, PF₆). A ¹H NMR spectrum of 3b exhibits broad paramagnetically-shifted resonances and the presence of CH₂Cl₂ solvate was indicated by a relatively sharp peak at 5.32 ppm.

2.3. Stopped-flow kinetics experiments

Rapid kinetics were followed using a SFL-21 (2-mm light path/2-mL syringes) low-temperature flow unit of a SF-3A stopped-flow system (Hi-Tech Scientific) combined with a TIDAS-16 HQ/UV-Vis 512/16B diode array spectrometer (J & M, 507 diodes, 300-720 nm, 1.3 ms minimum sampling time) using flexible light guides connected to a CLH-111 halogen lamp (ZEISS). The two glass coils, containing Cu(I) and oxygen solutions, respectively, and the mixing chamber, were immersed in an ethanol bath. This bath was placed in a dewar, which was filled with liquid nitrogen for low-temperature measurements. The ethanol bath was cooled by liquid nitrogen evaporation, and its temperature was measured by using a Pt resistance thermometer and maintained to 0.1 °C by using a temperature-controlled thyristor power unit (both Hi-Tech).

Data acquisition (up to 256 complete spectra, up to four different time bases) was done based on the kinspec program (J & M). For numerical analysis, all data were pretreated by factor analysis and concentration profiles were calculated by numerical integration using either Specfit (Spectrum Software Ass.) or Globfit (MATLAB).

The solvent CH₂Cl₂ (Uvasol, Merck) was dried with CaH₂ (Merck) and distilled under normal pressure immediately prior to use. For [Cu ^I₂(^Me2NXYL)](PF₆)₂ seven series of Cu(I) concentrations were used to carry out a total of 365 measurements between 90 and +20 °C. Of these 257 were used for the final analysis. The concentrations of Cu(I) solutions used were: 0.10, 0.17, 0.21, 0.37 mM (with a dioxygen concentration of 1.90 mM); 0.13 mM ([O₂] = 0.50 mM); 0.24 mM ([O₂] = 0.24 mM) and 0.12 mM ([O₂] = 0.13 mM). Reaction times measured ranged from 0.02 to 158 s.

For [Cu ^I₂(UN)](PF₆)₂ nine series of Cu(I) concentrations were used to carry out a total of 222 measurements between −90 and −35 °C. Two hundred and eighteen of these were used for the final analysis. For one series, a cutoff filter (480 nm) was used. The concentrations of Cu(I) solutions used varied between 0.14 and 1.08 mM. The dioxygen concentrations varied between 0.25 and 1.9 mM. Reaction times measured ranged from 172 to 227 s.

For [Cu ^I₂(^MeOXYL)](PF₆)₂, three series of Cu(I) concentrations were used to carry out a total of 91 measurements between 91 and +20 °C. Seventy-two of these were used for the final analysis. The concentrations of Cu(I) solutions used were 0.15, 0.30 and 0.41 mM. The dioxygen concentration used throughout was 1.9 mM. Reaction times measured ranged from 4 to 301 s.

2.4. Electrochemistry

Cyclic voltammetry was carried out by using a Bioanalytical Systems BAS-100B Electrochemistry Analyzer. The cell consisted of a modification of a standard three-chamber design equipped for handling air-sensitive solution by utillizing high-vacuum valve seals. A platinum disk (BAS MF 2013) was used as the working electrode. The reference electrode was Ag/AgNO₃. The measurements were performed at room temperature under an Ar atmosphere in DMF solution containing 0.2 M tetrabutylammonium hexafluorophophate and 10 ⁻⁴-10 ⁻³ M copper complex.

2.5. Magnetic moment measurements

Room temperature solution magnetic moments of the complex 3b were determined using the Evans method [13]. NMR sample tubes with screw caps and coaxial insert were purchased from the Wilmad Glass Company. Solutions of 3b with known concentrations were placed in the NMR sample tube and the deuterated reference solvent along with a small drop of undeuterated solvent to facilitate the identification of the reference solvent peak in NMR spectra were placed in the coaxial insert. The ¹H NMR spectra were taken on a Bruker AMX-300 NMR spectrometer. The chemical shift difference of the solvent was used to calculate the molar magnetic susceptibility of the complex based on the following equation:

$$C_M=(-3∕4p)(Dn∕n)(1000∕C)+C_M^{sol}-C_D,$$

where Dn is the paramagnetic shift of the solvent as the product of the chemical shift difference Dd and n (frequency of the NMR spectrometer), c is the sample concentration, $C_M^{sol}$ is the solvent susceptibility and c_D is the diamagnetic correction. $C_M^{sol}$ can be calculated as the product of c_g (value obtained from the CRC hand book) and the molecular weight of the metal complex. The value of c_D is obtained based on the molecular formulation of the complex and Pascal’s constants. The solution magnetic moment is related to the magnetic susceptibility as follows: μ = 2.84((c_MT/n))^1/2, where n is the nuclearity.

2.6. Low-temperature UV–Vis spectroscopy

Low-temperature UV–Vis spectra were recorded on a Hewlett-Packard HP8453 diode array spectrophotometer using a custom manufactured vacuum dewar equipped with quartz windows. To maintain and control low temperatures, a copper-tubing coil was inserted into the methanol-filled UV–Vis dewar through which cold methanol was circulated from an external source (Neslab Endocal refrigerated circulating bath). The temperature in the dewar was monitored by using a resistance thermocouple probe (Omega Model 651). The cuvette assembly consists of a quartz cell with 2 mm pathlength connected to a 14/20 female joint via a 10-cm long glass tube. The apparatus and procedures have also been previously described [14].

A solution of the Cu(I)-complex 1b-d with a known concentration was prepared in a glovebox and transferred into a cuvette with 2 mm pathlength. Then the cuvette was brought out of the glovebox and placed in the UV–Vis low-temperature (193 K) dewar. The UV–Vis spectrum of the copper(I) complex was recorded after 20 min. Then excess O₂ was introduced to the Cu(I) solution through a syringe needle and spectrum of the O₂-adduct was recorded and monitored over time to determine its stability.

3. Results and discussion

3.1. Employing ligand 4-pyridyl substituents in Cu^I/O₂ reactivity studies

The use of 4-pyridyl substituents has already proven very useful in a number of studies involving copper-dioxygen complex formation and reactivity, demonstrating that they mediate changes in the kinetics and thermodynamics of Cu₂O₂ complexes, the equilibria involving different structural isomers of Cu₂O₂ species, and the subsequent reactivity of such complexes [7,12]. In a notable study [12a], we showed that for [{(R-TMPA)Cu^II}₂(O₂²⁻)]²⁺, (Scheme 3), the observed m_O-O value shifts from 827 to 822 to 812 cm ⁻¹ and m_Cu-O (sym) shifts from 561 to 557 to 551 cm ⁻¹, respectively, as R-varies from H- to MeO- to Me₂N-. Increasing the N-donor strength to the copper decreases peroxide π*σ donation to the copper, weakening the Cu-O and O-O bonds. A decrease in ν_Cu-O of the bis-μ-oxo-dicopper(III) complex was also observed with increasing N-donor strength for the R-MePY2 ligand system (Scheme 3). However, no change was observed for m_O-O of the side-on peroxo moiety which is attributed to a compensating reduced charge donation from the peroxide π*σ orbital (to Cu) with increased ligand N-donor strength; this increases the effective negative charge on the peroxide moiety and thus reduces the backbonding from the Cu to the peroxide p*orbital. However, an increase in the amount of bis-μ-oxo-dicopper(III) isomer relative to side-on peroxo-dicopper(II) species is observed for R-MePY2 with R = H < MeO < Me₂N. The stronger donor ligands increase the stability of the Cu^III(μ-O)₂ isomer with its higher oxidation state metal ion. In a related study [15], copper ion complexes of so-called PYAN ligands (R-PYAN = N-[2-(4-R-pyridin-2-yl)-ethyl]-N,N′,N′-trimethyl-propane-1,3-diamine, R = Cl-, H-, MeO- and Me₂N-) were even more severely influenced with respect to the side-on peroxo-dicopper(II)/bis-μ-oxo-dicopper(III) equilibrium. For R = Cl-, the Cu₂O₂ species generated consisted entirely of the former isomer, while for Me₂N-, nearly entirely the latter.

Scheme 3.

Previously studied systems for copper(I)/O₂ kinetics employing derivatized pyridyl ligands.

Fig. 2.

Structure of the cationic portion of [Cu^I₂(^MeOXYL)]²⁺ (1b).

We also studied the kinetic-thermodynamic behavior for the formation and stability of [{(R-TMPA)Cu^II}₂(O₂²⁻)]²⁺ complexes [12b]. In EtCN, a strongly coordinating solvent, the ligand electronic effects were dampened due to the competitive binding of solvent and dioxygen to the copper centers, especially with weaker electron-donating groups such as Me and tBu, and there was little affect on kinetic and thermodynamic parameters for the formation of both the superoxocopper(II) species [(R-TMPA)Cu^II(O₂²⁻)]⁺ intermediate first formed (not shown) and the final product [{(R-TMPA)Cu^II}₂(O₂²⁻)]²⁺ species (Scheme 3). But with stronger MeO- and Me₂N-electron-donating groups, the rates of formation and stability of the product complexes were increased, as O₂-binding to copper(I) is a redox process, thus favored when oxidation of the metal ion is also preferred. Significant medium (solvent) effects were also seen, with the non-coordinating solvents THF or acetone. The rates of formation of both superoxo and peroxo complexes, along with their thermodynamic stabilities were much enhanced when stronger donor ligands such as MeO- or Me₂N-4-pyridyl substituents were employed.

As mentioned above, the complexes [{Cu(R-MePY2)}₂(O₂)]²⁺ possess a varying side-on-peroxo to bis-μ-oxo dicopper isomer ratio. In reactivity toward substituted anilines (Scheme 4), oxidative N-dealkylation takes place, but the mechanism of reaction shifts as one goes from R = H- to R = Me₂N- [12d]. The former complex, [{Cu(H-MePY2)}₂(O₂)]²⁺, is a stronger one-electron oxidant and possesses a lesser amount of Cu^III₂(μ-O)₂ species, the reaction appears to take place by a predominantly proton-coupled electron-transfer (PCET) or electron-transfer first/subsequent proton transfer (ETPT) mechanism. Hydrogen-atom transfer (HAT, or ET/PT) chemistry occurs for the [{Cu(Me₂N-MePY2)}₂(O₂)]²⁺ species, a weaker oxidant with respect to one-electron ET, but perhaps better able to accommodate a H-atom transfer due to its possession of a more basic bridging oxo group, because of the stronger ligand donors.

Scheme 4.

Cu₂O₂ mediated oxidative N-dealkylation of dimethylanilines.

3.2. Ligand syntheses

The ligand UN was prepared following a literature-based procedure [16], while ^MeOXYL and ^NMe2XYL were synthesized in a similar manner through reactions of m-xylyldiamine with 4-substituted-2-vinylpyridines (Scheme 5); the latter syntheses are published [12a,c]. Also, see the Section 2 for details. As can be seen, the substituents chosen for new studies are R = MeO- and Me₂N-, strong donors compared to R = H- (Fig. 1). Electron-withdrawing groups were not included in these studies since we previously observed that for mononuclear Cu(I) complexes of in our studies on the Cl-MePY2 (R = Cl-) (MePY2 = N-methyl-N,N-bis[2-(2-pyridyl) ethyl]amine), there was no reaction with dioxygen [12c].

Scheme 5.

Outline of the synthetic procedures to generate ^RXYL binucleating ligands.

3.3. Synthesis of copper(I) complexes

The dicopper(I) complexes [Cu^I₂(L)]²⁺ (1b-d, L = ^MeOXYL, ^Me2NXYL and UN, Fig. 1) were prepared by the addition of two equiv of [Cu^I(MeCN)₄](PF₆) to the CH₂Cl₂ solution of the appropriate ligand ^RXYL under an Ar atmosphere. Precipitation with diethyl ether followed by recrystallization affords light-yellow solids, which were characterized by ¹H NMR spectroscopy and elemental analyses. Repeated recrystallization from CH₂Cl₂/diethyl ether or MeOH/diethyl ether was necessary for complete removal of MeCN (derived from [Cu^I(MeCN)₄]PF₆) as a coordinating fourth ligand to copper(I) ion with its tridentate N₃ chelate. This procedure led to the desired three-coordinate dicopper(I) complexes 1b-d (Fig. 1).

3.4. X-ray structure of [Cu^I₂(^MeOXYL)]²⁺ (1b)

Light yellow crystals suitable for X-ray crystallography were obtained by slow diffusion of diethyl ether into an acetone solution of 1b. The structure includes a N,N-dimethyl formamide (DMF) solvate molecule, possibly introduced as a solvent impurity. The cationic portion of the structure is shown in Fig. 2.

The structural parameters are similar to the previously reported dicopper(I) structure with the parent ligand ^HXYL (Table 1) [6a]. As shown in Fig. 1, each Cu(I) ion is three-coordinate with ligation from two pyridine nitrogens and one tertiary amine nitrogen donor. The chelating tridentate ligand causes considerable distortions from ideal trigonal planar coordination. The geometry of the copper centers can be best described as bent “T-shape” geometry. The N_amino-Cu-N_PY (PY = 2-pyridyl) angles are acute within the range of 92-100°, resulting in large N_PY-Cu-N_PY angles of 163.9° and 166.5°. The Cu-N_PY bond distances (1.902-1.929 Å) in complex 1b are typical for Cu-N_heterocyclic (heterocyclic = derivatives of imidazole, pyrazole and pyridine) bond lengths in three-coordinate copper(I)-compounds [6a,17], and these short distances are approaching those found in purely linear two-coordinate systems with similar N-donor ligands, Cu-N ~ 1.86-1.89 Å [18]. The Cu-N_amine distances are much longer than those of Cu-N_PY distances, possibly due to the p back-donation (Cu(I) d electrons to the pyridine π* orbitals. As shown in Table 1, the Cu-N_amine distances in complex [Cu^I₂(^MeOXYL)](PF₆)₂ (1b) are longer by 0.1-0.2 Å than the corresponding bond lengths found in the parent complex [Cu^I₂(XYL)](PF₆)₂ (1a) and the N_PY-Cu-N_PY angles are −15° larger than those found in the parent complex. This may be a ligand electronic effect, where greater electron density around the copper(I) centers resulting from the electron-donating 4-MeO-groups disfavors the coordination of σ-type ligands such as amine nitrogens (i.e., giving longer Cu-N_amine distances). This results in larger N_PY-Cu-N_PY angles and more distortion from trigonal planar geometry.

Table 1.

Selected bond distances and angles for [Cu^I₂(XYL)](PF₆)₂ (1a) and [Cu^I₂(^MeOXYL)](PF₆)₂ (1b) complexes.

[CuI2(HXYL)](PF6)2 (1a)10				[CuI2(MeOXYL)](PF6)2 (1b)
Interatomic distances (Å)
Cu1–N1	2.121(8)	Cu1–N2	1.937(9)	Cu1–N1	2.289(7)	Cu1–N2	1.923(6)
Cu1–N3	1.924(8)	Cu2–N4	2.196(7)	Cu1–N31.917(6)	Cu2–N4	2.287(6)
Cu2–N5	1.904(8)	Cu2–N6	1.918(8)	Cu2–N51.929(8)	Cu2–N6	1.902(8)
Interatomic angles (°)
N1–Cu1–N2	102.5(3)	N1–Cu1–N3	100.4(3)	N1–Cu1–N292.2(3)	N1–Cu1–N3	100.1(3)
N2-Cu1–N3	151.4(4)	N4-Cu2–N5	99.7(3)	N2–Cu1–N3166.5(3)	N4-Cu2–N5	95.5(3)
N4-Cu2–N6	104.4(3)	N5-Cu2–N6	150.8(3)	N4–Cu2–N699.2(3)	N5-Cu2–N6	163.9(3)

Besides T-shaped copper(I) complexes, three-coordinate Y-shaped Cu(I)-complexes have also been described [19]. The bond angle h₃ and the “middle” bond distance d₃ increase as the geometry around the copper(I) distorts from the ideal trigonal planar geometry to T-shape, whereas both h₃ and d₃ decreases when the geometry moves from a trigonal planar to Y-shape geometry (see diagram).

3.5. Electrochemistry

The redox potentials for the binuclear copper(I) complexes were measured by cyclic voltammetry (CV) under an Ar atmosphere in N,N-dimethylformamide (DMF) solution containing 0.2 M [Bu₄N][PF₆]. The potentials, listed in Table 2, are reported versus the ferrocene/ferrocenium couple. All the Cu(I)-complexes exhibit single quasi-reversible redox behavior with an i_pa/i_pc ratio close to unity. Fig. 2 shows a typical CV scan here for [Cu^I₂(^MeOX-YL)](PF₆)₂ (1b) in DMF.

Table 2.

Cyclic voltammetry data for Cu(I)-complexes in DMF.^a

Complex	E½ (V)	ΔEp (mV)	ipa/ipc
[CuI2(HXYL)]2+ (1a)	−0.21	210	0.95
[CuI2(MeOXYL)]2+ (1b)	−0.25	170	0.85
[CuI2(NMe2XYL)]2+ (1c)	−0.28	300	1.15
[CuI2(UN)]2+ (1d)	−0.24	90	0.95

^aPotentials are vs. Fc/Fc⁺, measured under the same electrochemical cell conditions.

As shown in Table 2 and Fig. 3, only a single redox process is observed for these dicopper(I) complexes; no distinctive redox processes occur for the two Cu centers. A similar behavior was observed for analogous copper(I) complexes with either xylyl [20] or alkyl [20] linkers connecting two PY2 (=bis(2-(2-pyridyl)ethyl)amine) moieties. By contrast, differential pulse voltammetric measurements on bis(cyclam)dinickel and -dicopper complexes possessing variable methylene or xylyl linkers indicate that electrostatic effects and metal-metal distance as modified by linker length can influence the occurrence or extent of separation in CV experiments of the two metal-centered redox processes [21].

Fig. 3.

Cyclic voltammogram of [Cu^I₂(^MeOXYL)](PF₆)₂ (1b) in DMF.

As shown in Table 2, small redox potential shifts are observed as the ligands become more electron-donating; the E_1/2 values become more negative as the R group goes from H- to MeO- to Me₂N-. The observed ligand effect upon the redox potentials of the corresponding Cu(I)-complexes correlates to the differences in ligand basicity. The pKa values for pyridine, 4-MeO-pyridine and 4-Me₂N-pyridine are 5.21, 6.58 and 9.70, respectively. The more basic ligand provides more electron density to the copper center, making it easier to oxidize Cu(I) to Cu(II). The same correlation between ligand basicity and the E_1/2 values was observed for the mononuclear copper(I) analogues with 4-pyridyl substituted MePY2 ligands. However, more negative E_1/2 values (−0.31 V for parent MePY2, −0.36 for MeO-MePY2 (with 4-MeO-pyridyl groups) and −0.44 for Me₂N-MePY2 (with 4-Me₂N-pyridyl groups)) are observed for the mononuclear copper(I) complexes compared to those for the corresponding dicopper(I) complexes. This may be due to the relative electron-donating capacities of the substituents on the central amine nitrogens, i.e. for the binuclear Cu(I)-complexes with xylyl linker, benzyl (PhCH₂) groups are less electron-donating than Megroups in the mononuclear copper(I) complexes [Cu^I(R-MePY2)]⁺, resulting in more positive redox potential for the binuclear complexes. However, a smaller shift (40 mV) was observed for the analogous mononucleating PhCH₂PY2 complex [22] where a Me- group is replaced with a PhCH₂- group. Therefore, the observed more significant difference in redox potentials for the binuclear Cu(I) complexes 1a-d and the mononuclear Cu(I)-complexes [Cu^I(R-MePY2)]⁺ may arise from some interaction between the two copper ions in the dicopper(I) complexes, where a nearby positively charged center makes the oxidation of one of the Cu(I) ions more difficult. In addition, relatively bigger shifts resulting from modifications in ligand electron-donating ability are observed for the mononuclear copper(I) complexes (130 mV shift from MePY2 to Me₂N-MePY2 complexes) compared to those for the binuclear copper(I) complexes [Cu^I₂(L)]²⁺ (1a-c) (70 mV shift from XYL to ^Me2NXYL complexes), described here, again suggesting possible interaction between the two copper centers in the binuclear complexes. In summary, it seems that the effect of ligand electron-donating ability on the redox potential of the metal ions in these binuclear complexes is reduced due to some (electrostatic) interaction between the two copper(I) centers.

A small negative shift in the E_1/2 value is observed for [Cu^I₂(UN)]²⁺ (1d) with the unsymmetrical ligand environment, compared to that for [Cu^I₂(^HXYL)]²⁺ (1a). Following the arguments made above, ligand constraints in UN may cause the copper ions to be held further apart, making it easier to oxidize Cu(I) to Cu(II) in complex.

Both ligand and structural properties are thus shown to affect the redox potentials of the corresponding copper(I) complexes. We previously observed that the difference in electrochemical behavior of copper(I) complexes with pyridyl and/or quinolyl ligands tended to correlate with variations in their O₂ reactivity [23]. The kinetic studies described below were carried out to investigate both ligand electronic and structural effects on the O₂-interactions with the dicopper(I) complexes described here.

3.6. Isolation of the hydroxylation products

As previously described [6a,c,16], the oxygenation of [Cu^I₂(X-YL)](PF₆)₂ and [Cu^I₂(UN)](PF₆)₂ in DMF at room temperature gives the corresponding phenoxo and hydroxo bridged dicopper(II) complexes [Cu^II₂(XYLO⁻)(⁻OH)](PF₆)₂ (3a) and [Cu^II₂(UN-O⁻)(⁻OH)](PF₆)₂ (3d). The stoichiometry of these reactions has been established by manometric measurements for O₂ uptake (Cu/O₂ = 2:1).¹⁰ Mass spectrometric analysis of the oxygenated product of [Cu^I₂(XYL)](PF₆)₂ (1a) prepared by using isotopically pure ¹⁸O₂ shows that both oxygen atoms in the [Cu^II₂(XY-LO⁻)(⁻OH)](PF₆)₂ complex were incorporated into the final hydroxylation product.

Similarly, when the DMF solution of [Cu^I₂(^MeOXYL)](PF₆)₂ (1b) was oxygenated at room temperature, a solution color change from yellow to dark green was observed. Isolation of the green complex yielded the hydroxylation product [Cu^II₂(^MeOXYLO⁻)(⁻OH)](PF₆)₂ (3b) characterized by elemental analysis, X-ray crystallography, UV–Vis and IR spectroscopies, as well as room-temperature magnetic moment measurements. This complex exhibits UV–Vis features (k_max = 370 nm (ε = 3450), 637 (ε = 150)), which are similar to those for complexes 3a and 3d [7,16]. The intense 370 nm absorption is assigned as an OH⁻ → Cu(II) ligand-to-metal charge-transfer band and the weak absorption band at 637 nm is consistent with a d-d transition. The IR spectrum of complex 3b exhibits a m_O-H at 3643 cm ⁻¹. The room temperature magnetic moment of complex 3b is 1.1 l_B/Cu, comparable with 0.8 l_B/Cu for 3a [24], which indicates efficient antiferromagnetic coupling between the two Cu(II) centers through the bridging phenoxo and hydroxo ligands. Less efficient coupling is observed for the unsymmetrical complex [Cu^II₂(UN-O⁻)(OH⁻)]²⁺ (3d) with a room temperature magnetic moment of 1.5 l_B/Cu [16], manifesting the effect of ligand constraints in a distorted structure, where less efficient overlap of orbitals on the two separate Cu ions occurs. However, under synthetic conditions in CH₂Cl₂ solvent, the hydroxylation reaction seems to be least efficient for [Cu^I₂(^Me2NXYL)]²⁺ (1c), where the unreacted ligand (approximately 40%) as well as the hydroxylated ligand were recovered after the oxygenation reaction. The lower hydroxylation efficiency in 1c may arise from competition for an interrather than intra-molecular O₂-binding process; this has previously been suggested in the oxygenation reactions of a series of analogous binuclear Cu(I)-complexes [7].

3.7. X-ray crystal structure of [Cu^II₂(^MeOXYLO⁻)(⁻OH)](PF₆)₂ (3b)

Green crystals suitable for X-ray crystallography were obtained by slow diffusion of diethyl ether into a methylene chloride solution of 3b. Selected bond distances and angles are tabulated in Table 3, along with the corresponding structural parameters for the parent complex [Cu^II₂(^HXYLO⁻)(⁻OH)](PF₆)₂ (3a) previously reported [6a] for comparison. An ORTEP view with atom labeling scheme is shown in Fig. 4.

Table 3.

Selected bond distances and angles for 3c and 1c.

[CuI2(MeOXYLO−)(−OH))](PF6)2 (3c)				[CuI2(XYLO−)(−OH))](PF6)2 (1c)
Interatomic distances (Å)
Cu1–N1	2.054(8)	Cu1–N2	1.980(8)	Cu1–N1	2.034(14)	Cu1–N2	2.006(16)
Cu1–N3	2.180(8)	Cu2–N4	2.056(8)	Cu1–N3	2.258(13)	Cu2–N4	2.028(13)
Cu2–N5	2.190(8)	Cu2–N6	2.003(8)	Cu2–N5	2.149(15)	Cu2–N6	2.027(14)
Cu1–O1	1.969(6)	Cu1–O2	1.959(6)	Cu1–O1	1.979(10)	Cu1–O2	1.938(10)
Cu2–O1	1.987(6)	Cu2–O2	1.968(6)	Cu2–O1	1.972(11)	Cu2–O2	1.962(10)
Interatomic angles (°)
O1–Cu1–O2	75.6(3)	O1–Cu1–N1	93.0(3)	O1–Cu1–O2	76.7(4)	O1–Cu1–N1	92.4(5)
O1–Cu1–N2	162.8(3)	O1–Cu1–N3	98.4(3)	O1–Cu1–N2	157.1(6)	O1–Cu1–N3	97.6(5)
O2–Cu1–N1	159.7(3)	O2–Cu1–N2	92.7(3)	O2–Cu1–N1	161.5(5)	O2–Cu1–N2	88.8(6)
O2–Cu1–N3	101.5(3)	N1–Cu1–N2	94.4(3)	O2–Cu1–N3	100.0(5)	N1–Cu1–N2	96.5(6)
N1–Cu1–N3	96.6(3)	N2–Cu1–N3	96.2(3)	N1–Cu1–N3	96.1(5)	N2–Cu1–N3	102.4(6)
O1–Cu2–O2	74.9(3)	O1–Cu2–N4	92.5(3)	O1–Cu2–O2	76.3(4)	O1–Cu2–N4	93.9(5)
O1–Cu2–N5	98.3(3)	O1–Cu2–N6	161.1(3)	O1–Cu2–N5	98.3(5)	O1–Cu2–N6	161.3(5)
O2–Cu2–N4	162.0(3)	O2–Cu2–N5	97.9(3)	O2–Cu2–N4	163.5(5)	O2–Cu2–N5	98.9(5)
O2–Cu2–N6	92.7(3)	N4–Cu2–N5	96.7(3)	O2–Cu2–N6	91.4(5)	N4–Cu2–N5	95.6(6)
N4–Cu2–N6	95.8(3)	N5–Cu2–N6	97.6(3)	N4–Cu2–N6	94.4(6)	N5–Cu2–N6	97.6(6)
Cu1–O1–Cu2	104.2(3)	Cu1–O2–Cu2	105.3(3)	Cu1–O1–Cu2	102.5(5)	Cu1–O2–Cu2	104.4(5)
Cu1–O1-C17	128.8(6)	Cu2–O1-C7	127.0(6)	Cu1–O1-C17	129.6(10)	Cu2–O1-C17	127.1(10)

Fig. 4.

ORTEP diagram of the cationic portion of [Cu^II₂(^MeOXYLO⁻)(OH)]²⁺ (3b).

Each copper ion is coordinated with three nitrogen atoms from the tridentate chelate along with the bridging-phenolate (O1) and hydroxide oxygen (O2) atoms in a slightly distorted square pyramidal geometry. The structural parameters are very similar to those of the parent complex, as shown in Table 3. The N3 and N5 pyridine nitrogen atoms occupy axial positions in an anti conformation with respect to the Cu₂O₂ plane. These axially coordinated Cu-N_PY distances are 0.20 and 0.19 Å longer than those in the basal plane. The basal Cu-N_PY (N2 and N6) bond lengths average 1.99 Å, which are ca. 0.2 Å shorter than that in the parent complex [Cu^II₂(XYLO⁻)(⁻OH)](PF₆)₂ (3a) [6a], possibly due to the greater electron-donating abilities of the 4-MeO- compared to the unsubstituted pyridine nitrogen atoms.

Similar to that in the parent complex, the Cu₂O₂ unit is essentially planar and the phenolate ring is twisted with respect to the planar Cu₂O₂ unit. The ligand orientation in these two complexes show substantial differences from the phenolate- and hydroxybridged binuclear Cu(II) complex with UN ligand, where the unsymmetrical ligand environment enforces the Cu₂O₂ unit into a butterfly shape [16] and the two axial nitrogens engage in a cis conformation, in contrast to the anti conformation observed in the two symmetrical complexes (3a and 3b). These structural differences have been shown to decrease the antiferromagnetic coupling between the two Cu(II) ions and to affect the O₂-adduct formation and subsequent hydroxylation reactions, based on bench-top observations and kinetic measurements discussed below.

3.8. Bench-top oxygenation reactions in CH₂Cl₂ at low temperature

When a bench-top low-temperature (193 K) oxygenation of [Cu^I₂(^HXYL)]²⁺ (1a) is carried out, no intermediate species [Cu^II₂H₂(XYL)(O₂²⁻)]²⁺ could be observed spectroscopically. However, under the same conditions, the O₂-adducts [Cu^II₂(L)(O₂²⁻)]²⁺ (2b-d, L = ^MeOXYL, ^Me2NXYL, UN) with characteristic UV–Vis absorption features could be detected. The stability of [Cu^II₂MeO₂(XYL)(O₂²⁻)]²⁺ (2b) and [Cu^II(^Me2NXYL)(O₂²⁻)]²⁺ (2c) was still somewhat limited as indicated by the decrease in intensity of their characteristic electronic absorptions over time (5 min). By contrast, the UN complex [Cu^II₂(UN)(O₂²⁻)]²⁺ (2d) was much more stable at 193 K (no intensity decrease over 1 h) [16]. These bench-top observations manifest the ligand electronic or structural effects upon copper(I)-dioxygen adducts. However, rapid kinetic studies were needed to further understand, in a quantitative manner, how the ligand environment affects the O₂-interactions of these dicopper(I) complexes (1b-d), followed by arene hydroxylation chemistry.

3.9. Stopped-flow kinetic studies

3.9.1. Overall findings

As mentioned in the Section 1, detailed kinetics studies for the oxygenation reactions of [Cu^I(R-XYL)]²⁺ (R on the central arene para to the position which is hydroxylated; R = -NO₂, -H, -C(CH₃)₃, -F) conform to Scheme 1, with the results providing corroborative evidence for a hydroxylation reaction mechanism involving an electrophilic peroxo species. Here, in our full investigation and kinetic analysis, we find that the reactions of the dicopper(I) complexes [Cu^I₂(L)]²⁺ (1b-d, L = ^MeOXYL, ^Me2NXYL and UN, Fig. 1) with O₂ in dichloromethane follow the same mechanism found for [Cu^I₂(XYL)]²⁺ (1a) (Schemes 1 and 6). At low temperature, [Cu^I₂(L)]²⁺ (1a-d, L = XYL, ^MeOXYL, ^Me2NXYL and UN) binds dioxygen reversibly (k₁, k₋₁) to form a μ-g²:g²-(side-on) peroxo species [Cu^II₂(L)(O₂²⁻)]²⁺ (2a-d), which subsequently hydroxylates the arene (k₂) to give [Cu^II₂(LO⁻)(⁻OH)]²⁺ (3a-d) products. Fig. 5 shows the UV–Vis behavior, when, for example, [Cu^I₂(^MeOXYL)]²⁺ (1b) reacts with excess O₂ at −50 °C. Upon rapid mixing of solution of O₂ and 1b, there is very rapid formation of [Cu^II (^MeOXYL)(O₂²²⁻)]²⁺ (2b) with the predominant charge-transfer absorption at λ_max = 363 nm (ε = 14000) and the accompanying band 437 nm (ε = 2700), past proven to be a species with the μ-η²:η²-side-on peroxodicopper(II) core. In a unimolecular process 2b subsequently decays and is transformed into [Cu^II₂(^MeOXYLO⁻)(⁻OH)]²⁺ with 370 and 650 nm absorptions. The inset of Fig. 5 shows the absorbance versus time trace at 363 nm, demonstrating a perfect correlation between experimental data and the suggested mechanism (Scheme 6).

Scheme 6.

Kinetic scheme deduced for the [Cu^I₂(^RXYL)]²⁺ (1a-c) oxygenation reactions employing PY′ donor ligands.

Fig. 5.

Time-dependent UV–Vis spectra for the oxygenation of [Cu^I₂(^MeOXYL)]²⁺ (1b) at −50 °C. Inset: absorbance vs. time at 363 nm; (x) experimental data; (−) calculated curve, based on the model described in Scheme 6. See text for further details.

It is found that at low temperature, the formation of the side-on peroxo species [Cu^II₂(L)(O₂²⁻)]²⁺ (2) is complete, thus k₋₁ being irrelevant. However, at higher temperatures, k₋₁ has to be taken into account since the intermediate 2 is then only partially formed. Similar to the previously studied analogous systems [7,25], the hydroxylation step (k₂) is composed of a thermal and of a photochemical temperature-independent term, and the latter dominates below ~−40 °C. Thus, data acquired only at higher temperatures, where the photochemical terms could be accounted for and subtracted out, were used for determination of the activation parameters for k₂. An exceptional case was that for the oxygenation reaction of [Cu^I₂(UN)](PF₆)₂ (1d), where the photochemistry dominated in the whole temperature range (183-238 K) studied, initially precluding the determination of the thermal parameters of k₂ for 1d. The problem could be solved however, as discussed below.

The variable-temperature stopped-flow data of the oxygenation reactions of [Cu^I₂(L)]²⁺ (1a-d, L = ^HXYL, ^MeOXYL, ^Me2NXYL and UN) were analyzed by global analysis in the eigenvector space [23] and the kinetic and thermodynamic parameters deduced are listed in Tables 4 and 5. For the temperature-dependent reactions of interest (k₁, k₋₁ and k₂), plots of the kinetic and thermodynamic parameters log(k/T) or log(K/T) are all well behaved. Eyring plots of k₁, k₋₁ and k₂ for the ^MeOXYL complex are shown in Fig. 6.

Table 4.

Kinetic parameters for O₂-interactions with dicopper(I) complexes 1a–d.

Parameters	HXYL (1a)a	MeOXYL (1b)	Me2NXYL (1c)	UN (1d)
k1 (M−1 s−1)
ΔH≠ (kJ mol−1)	8.2 ± 0.1	6.9 ± 0.3	2.3 ± 0.2	12.3 ± 0.2
ΔS≠ (J K−1 mol−1)	−146 ± 1	−149 ± 1	−156 ± 1	−132.4 ± 0.8
183 K	385 ± 5	(6.9 ± 0.2) × 102	(6.3 ± 0.2) × 103	(1.38 ± 0.02) × 102
223 K	(1.24 ± 0.01) × 103	(1.91 ± 0.03) × 103	(1.01 ± 0.01) × 104	(7.23 ± 0.01) × 102
298 K	(5.1 ± 0.1) × 103	(6.5 ± 0.3) × 103	(1.84 ± 0.06) × 104	(5.2 ± 0.2) × 103
k−1 (s−1)
ΔH≠ (kJ mol−1)	70 ± 1	78 ± 1	71 ± 2	67 ± 1
ΔS≠ (J K−1 mol−1)	50 ± 6	76 ± 5	60 ± 8	63 ± 6
183 K	(1.5 ± 0.4) × 10−5	(2.5 ± 0.7) × 10−6	(3 ± 1) × 10−5	(5.4 ± 0.9) × 10−4
223 K	(7.20 ± 0.05) × 10−2	(2.9 ± 0.3) × 10−2	(1.3 ± 0.2) × 10−1	1.79 ± 0.1
298 K	(1.3 ± 0.2) × 103	(1.4 ± 0.2) × 103	(2.8 ± 0.5) × 103	(2.2 ± 0.5) × 104
k2 (s−1)
ΔH≠ (kJ mol−1)	50 ± 1	48.9 ± 0.6	54 ± 2	57 ± 1b
ΔS≠ (J K−1 mol−1)	−35 ± 2	−47 ± 2	−39 ± 7	−26 ± 6b
183 K	(3.4 ± 0.2) × 10−4	(1.5 ± 0.2) × 10−4	(2.1 ± 0.6) × 10−5	(1.1 ± 0.2) × 10−5 b
223 K	(1.49 ± 0.03) × 10−1	(6.0 ± 0.2) × 10−2	(1.4 ± 0.1) × 10−2	(1.04 ± 0.06) × 10−2 b
298 K	172 ± 8	61 ± 3	26 ± 4	31 ± 8b

^aPreviously published results [7].

^bData obtained using single wavelength (361 nm) interrogation for the kinetic study.

Table 5.

Equilibrium parameters for O₂-interaction with dicopper(I) complexes 1a–d.

Parameters	HXYL (1a)	MeOXYL (1b)	Me2NXYL (1c)	UN (1d)
K1 (M−1) (=k1/k−1)
ΔH≠ (kJ mol−1)	−62 ± 1	−71 ± 2	−69 ± 2	−55 ± 1
ΔS≠ (J K−1 mol−1)	−196 ± 6	−225 ± 5	−222 ± 9	−196 ± 6
183 K	2.6 × 107	(2.8 ± 0.8) × 108	(3 ± 1) × 108	(2.6 ± 0.4) × 105
223 K	1.7 × 104	(6.6 ± 0.6) × 104	(7.5 ± 0.8) × 104	(4.0 ± 0.1) × 102
298 K	3.9	4.5 ± 0.5	6 ± 1	(2.4 ± 0.5) × 10−1

Fig. 6.

Eyring plots for k₁, k₋₁, and k₂ for [Cu^I₂(^MeOXYL)]²⁺ (1b): k, Boltzmann constant; h, Planck constant; [O₂] = 1.9 × 10 ⁻³ M; [1b] (Δ), 1.5 × 10 ⁻⁴ M; (O), 3.0 × 10 ⁻⁴ M; (1), 4.1 × 10 ⁻⁴ M.

3.9.2. Reversible O₂-binding to the dicopper(I) complexes [Cu^I₂(L)]²⁺ (1)

As seen from Table 4, the binding of O₂ (k₁) to the dicopper(I) complexes 1 to form the side-on peroxo species [Cu^II₂(L)(O₂²⁻)]²⁺ (2) is characterized by very low activation enthalpies of less than 12 kJ mol ⁻¹ and large compensating activation entropies of −136 to −156 J K ⁻¹ mol ⁻¹. Such observations may be rationalized on the basis of having reactive, easily oxidizable and coordinatively unsaturated (three coordinate) cuprous species. An additional reason for the observed low activation enthalpies may arise from the composite nature for the formation of the peroxodicopper(II) species [Cu^II₂(L)]²⁺ (2) since it is unlikely that O₂ would bind to both copper ions simultaneously. Rather, as described in previous studies, we suggest that the dioxygen first binds to one copper moiety to generate the superoxo intermediate (i.e. Cu^ICu^II-O₂⁻) in a rapid, left-lying equilibrium and only in a second step forms the spectroscopically detectable μ-peroxo dicopper(II) moiety. Similar low or even negative activation enthalpies (i.e. consistent with a multi-step process) are observed for the formation of the peroxodicopper(II) complexes with a series of binucleating ligands Nn having the identical ligand set (i.e. the PY2 tridentate) linked instead by alkyl chains [26]. In addition, the formation of the μ-1,2-(end-on)-peroxodicopper(II) complexes with a series of mononucleating tetradentate ligands exhibit negative activation enthalpies[12b,23,27] as a result of a similar pre-equilibrium step (formation of the initial Cu^I/O₂ adduct, i.e. a superoxocopper(II) complex).

By contrast, Tolman, Zuberbühler and co-workers [9a,10a,28], through detailed stopped-flow kinetic studies, demonstrated that the formation of the O₂-adducts of mononuclear or binuclear copper(I) complexes containing substituted triazacyclononane (tacn) tridentate ligands is associated with considerable high activation enthalpies (ΔH^≠ around 40 kJ mol ⁻¹). These relatively large ΔH^≠ values are comparable to those seen for the formation of Cu-O₂ 1:1 (i.e. superoxocopper(II) species) adducts in reactions of well-studied mononuclear copper(I) complexes with TMPA (=tris(2-pyridylmethyl)amine) and its analogs [12b,23]. Thus, it is concluded that the formation of superoxo intermediates, which are not spectroscopically observable in these tacn derivative cases, are the rate-determining steps in the oxygenation reactions of the corresponding copper(I) complexes. However, a stopped-flow kinetic study[12b] on the oxygenation reactions of a series of mononuclear copper(I) complexes with 4-pyridyl substituted TMPA ligands in propionitrile show that significant activation enthalpies (29-32 kJ mol ⁻¹) associated with the formation of the superoxocopper(II) complexes involves the (rapid) dissociation of a coordinated nitrile in a strongly left-lying equilibrium. In addition, when weakly coordinating solvents such as tetrahydrofuran (THF) and acetone are used as solvents, much faster formation of superoxocopper(II) species are observed, demonstrating the importance of the solvent medium, where nitriles, which are excellent strong ligands for copper(I) centers [29], can compete with the O₂-binding process.

More interestingly, for the oxygenation reactions of the dicopper(I) complexes [Cu^I₂(^RXYL)]²⁺ (1a-c), the rate of dioxygen binding (k₁) increases (Table 4) as the ligand becomes more electron-donating (i.e. XYL to ^MeOXYL to ^Me2NXYL), consistent with oxygenation being accompanied by electron transfer from copper(I) to O₂ (giving a superoxo or peroxo anion). The activation enthalpy ΔH^≠ for k₁ also exhibits a trend, where ΔH^≠ decreases as the ligand electron-donating ability increases, again suggesting more favorable dioxygen binding with strong electron-donating ligands. Presumed ligand constraints (as seen in X-ray structure of [Cu^II₂(UN-O⁻)(⁻OH)]²⁺ (3d), discussed above) are also shown to affect Cu₂-O₂ binding, where a slower dioxygen binding (i.e. smaller k₁) and a higher activation enthalpy ΔH^≠ (for k₁) are observed for the copper(I) complex [Cu^I₂(UN)]²⁺ (1d) with the unsymmetrical ligand UN (Fig. 1), compared to that for the parent ligand XYL (Table 4). The observed differences suggest that for these binuclear copper(I)-complexes, the formation of the superoxo intermediates (i.e. initial O₂-adduct) is unlikely to be the rate-determining step since it is hard to conceive that individual copper(I) centers with nearly identical coordination environments (UN versus XYL) would exhibit different intrinsic reactivity toward O₂ (for the formation of superoxo intermediates), again manifesting the composite nature of these oxygenation reactions.

In contrast to the observed very low activation enthalpies (<12 kJ mol ⁻¹) for k₁, large activation enthalpies ΔH^≠ (67-78 kJ mol ⁻¹) along with positive activation entropies are associated with the dissociation (k₋₁) of O₂ from Cu₂O₂ complexes. Although not dramatic effects, a larger rate constant k₋₁, as well as a lower activation enthalpy (associated with k₋₁) are observed for the UN complex compared to that for XYL complexes (Table 4), suggesting that ligand constraints in [Cu^II₂(UN)(O₂²⁻)]²⁺ (2d) decreases its copper-dioxygen bonding strength, k₋₁, representing the breaking of a Cu-O bond and release of O₂. However, no direct correlation between the ligand electron-donating ability and the copper-dioxygen binding strength (ΔH⁻ for k₋₁, Table 4) is observed. In fact, the ^MeOXYL complex exhibits a higher activation enthalpy k₋₁) compared to that for the parent XYL ligand, which is in line with the notion that electron-donating ligand somehow increases the Cu-O bond strength in the peroxo complex 2. Perhaps Cu(II) → O (peroxide) back-bonding is not insignificant in [Cu^II₂(L)(O₂²⁻)]²⁺ (2) complex, as has been discussed with respect to the electronic structure and bonding in μ-η²:η²-(side-on)-peroxodicopper(II) species [30]. Nevertheless, ΔH^≠ and k₋₁ values are almost identical for the ^Me2NXYL (most electron-donating ligand) and ^HXYL complexes, exhibiting no effect of a more electron-donating ligand (^Me2NXYL). As mentioned, it was proposed that there is some interaction between the dicopper-bound electrophilic peroxo group and the xylyl p systems in [Cu^II₂(L)(O₂²⁻)]²⁺ (2) due to their close proximity. Perhaps this can provide an explanation for the absence of a correlation between the copper-dioxygen bond strength in the peroxo complex 2 and the ligand electronic-donating ability.

From k₁ and k₋₁, equilibrium constants k₁ = k₁/k₋₁ and the corresponding thermodynamic parameters ΔH° and ΔS° can be calculated; they are collected in Table 5. Here, again, the thermostability (ΔH°) of the peroxo complex [Cu^II₂(L)(O₂²⁻)]²⁺ (2) appears to correlate with the effect of ligand constraints, where the unsymmetrical ligand UN affords a less stable peroxo complex compared to that with XYL ligand, reflected by their k₁ and ΔH° values. Ligand constraints or steric hindrance have well been demonstrated to influence the stability of copper-dioxygen adducts formed [27,31]. A complex possessing a binucleating analogue (DO, Fig. 7) of TMPA with -CH₂OCH₂-linked pyridyl groups exhibits significant stability in acetone and in fact, the thermodynamic parameters could not be obtained (i.e. since its formation is essentially complete at all temperatures) [31a]. By contrast, an analogous ligand (D¹, Fig. 7) with -CH₂CH₂-linker leads to a strained and much less stable dioxygen adduct. In addition, in the oxygenation reactions of a series of binuclear copper(I)-complexes (Nn, Fig. 7) containing the same tridentate PY2 units [26], the length of the alkyl chain (i.e. -(CH₂)_n-) connecting the two PY2 units appears to have a remarkable effect on the equilibrium constants k₁ for the formation of the μ-η²:η²-peroxodicopper(II) complexes [Cu^II(Nn)(O₂²⁻)]²⁺. The [Cu^I₂(N5)]²⁺ complex with the longest alkyl linker (-(CH₂)₅-) affords the largest k₁ value (i.e. most stable peroxo complex), which is in accord with the view that the N5 complex forms the most favorable Cu₂O₂ structure due to the relatively unconstrained nature of the ligand. This is consistent with a resonance Raman study [32] of Cu₂O₂ complexes with Nn ligands, where the N5 complex exhibits the lowest m(O-O) value (due to enhance back donation form Cu to the peroxo σ* orbital within the planar Cu₂O₂ core), closest to that for the unconstrained mononuclear Cu-MePY2 complex, [{Cu^II(MePY2)}₂(μ-η²:η²-O₂²⁻)]²⁺. Comparing to these Nn ligands, the ^RXYL and UN ligands are more constrained because of the bulky and less flexible xylyl groups connecting the PY2 tridentate moieties. Thus, the overall reaction enthalpies ΔH° for the XYL and UN systems are considerably less negative than those with Nn ligands, −56 kJ mol ⁻¹ for UN, −62 to −71 kJ mol ⁻¹ for ^RXYL, −81 to the −84 kJ mol ⁻¹ for Nn, all forming dicopper complexes with [Cu^II₂(μ-η²:η²-O₂²⁻)]²⁺ core.

Fig. 7.

Dinucleating ligands used in other Cu^I₂/O₂ kinetic studies.

The substituted XYL ligands (^MeOXYL and ^Me2NXYL) afford more stable peroxo complexes compared to that with the parent XYL ligand (i.e. more negative ΔH° and bigger k₁, Table 5). However, the peroxo complexes with ^MeOXYL and ^Me2NXYL ligands exhibit nearly identical thermostabilities (Table 5), again reflecting no direct correlation between the stability of the peroxo species 2 and the ligand electron-donating ability, as said before possibly due to interaction between the peroxo core and π orbitals of xylyl linker group.

In summary, the thermodynamic parameters (ΔH° and ΔS°) clearly indicate that the dioxygen binding to these dicopper(I) complexes [Cu^I₂(L)]²⁺ (1) is entirely driven by enthalpy, while large negative entropies result in their room-temperature instability. Ligand constraints are shown to have a significant effect on the dioxygen binding (k₁) as well as the stability (K₁, ΔH°) of the resulting peroxo species [Cu^II₂(L)(O ^2-₂)]²⁺ (2). Ligand electron-donating ability appears to affect the binding of dioxygen (k₁ and ΔH^≠ for k₁) to the copper(I) complexes, where more electron-donating ligands favor O₂-binding with smaller activation enthalpies. However, no direct correlation between the stability of the resulting peroxo complexes 2 with ligand electronic properties are observed, possibly due to some interaction between the Cu₂O₂ core and xylyl p orbitals (as previously mentioned).

3.9.3. Arene hydroxylation step k₂

As shown in Fig. 5, [Cu^II₂(^MeOXYL)(O₂²⁻)]²⁺ (2b) as well as the other side-on peroxo species [Cu^II₂(L)(O₂²⁻)]²⁺ (2) are not stable (except the UN complex 2d, vide infra) even at low temperature, as reflected by the disappearance of their UV–Vis features, and their subsequent decay leads to arene hydroxylation (Schemes 1 and 6). However, bench-top observations mentioned above clearly indicate that complex 2d is much more stable (i.e. longer-lived by bench-top UV–Vis monitoring) than the other peroxo complexes with XYL or substituted XYL ligands (^MeOXYL and ^Me2NXYL). This suggests that subsequent hydroxylation reaction is less favorable with the UN ligand, which possesses an unsymmetrical and strained ligand environment. A nearly ideal orientation of the arene substrate and the peroxo core in [Cu^II₂(R-XYL)(O₂²⁻)]²⁺ complexes (Schemes 1 and 6) are proposed for these facile hydroxylation reactions [6b,27]. Thus, possible disruption of this desirable orientation caused by ligand constraints of the unsymmetrical UN ligand inhibits the occurrence of a fast arene hydroxylation reaction which otherwise might be expected for 2d at low temperature (<193 K). However, at higher temperatures, 2d is not stable and it does decay, leading to the formation of the well-characterized hydroxylation product [Cu^II₂(UN-O⁻)(⁻OH)]²⁺ (3d) [16].

As mentioned, the experimental methods used for most of the kinetic studies leads to dominating photochemistry relating to the hydroxylation step (k₂) for [Cu^II₂(UN)(O₂²⁻)]²⁺ (2d); the diode array spectrometer inputs rather intense light over the whole spectral region. As copper-dioxygen complexes in general posssess very strong peroxo-to-Cu(II) charge-transfer absorptions, it is not too surprising that photochemical reactions would exist. In fact, metal-dioxygen complex photochemistry in general has and does attract general interest [33], and Cu_n(O₂)]ⁿ⁺ complexes are candidates for future studies. We found that the way around the photochemistry issue for the UN system was to go to single-wavelength studies. Thus, Fig. 8 shows data obtained via 361 nm monitoring of the [Cu^I₂(UN)]²⁺ (1d) + O₂ reaction. From such data, valid kinetic parameters were obtained for k₂ (Table 4). Fig. 9 shows an Eyring plot for k₂ avoiding photochemistry and a comparison to the data obtained when using the diode array spectrometer.

Fig. 8.

Time-dependent spectra and 361 nm monitoring of the [Cu^I₂(UN)]²⁺ (1d) + O₂ reaction in CH₂Cl₂ at 183 K.

Fig. 9.

Eyring plots on data obtained for the hydroxylation reaction (k₂) of [Cu^II₂(UN)(O₂²⁻)]²⁺ (2d) using either the diode array spectrometer (red) or single wavelength monitoring (blue; lower straight line). (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.)

In line with the benchtop spectroscopic findings for [Cu^I₂(UN)]²⁺ (1d) oxygenation mentioned above, the k₂ kinetic parameters reveal that indeed, the rate of the arene hydroxylation reaction, [Cu^II₂(UN)(O₂²⁻)]²⁺ (2d) → [Cu^II₂(UNO⁻)(⁻OH)]²⁺ (3d), is less than is found for the other ^RXYL dicopper complex analogues. This especially obvious if one compares k₂ values and corresponding activation parameters with the hydroxylation reaction for the parent complex [Cu^II₂(^HXYL)(O₂²⁻)]²⁺ (Table 4). For example, at 223 K, k₂ is 15 times smaller for the UN system. While not a huge difference, it is clearly sufficient to have allowed facile bench-top UV–Vis spectroscopic monitoring of 2d at 183 K (vide supra) [16], while [Cu^II₂(^HXYL)(O₂²⁻)]²⁺ (2a) cannot be detected. As discussed, the slowed arene hydroxlyation rate for [Cu^II₂(UN)(O₂²⁻)]²⁺ (2d) undoubtedly reflects a less than ideal proximity or orientation of the complex’s electrophilic peroxo group toward the arene psystem.

Ligand substituents also affect the arene hydroxylation reactions of the peroxo complexes [Cu^II₂(L)(O₂²⁻)]²⁺ (2). As seen in Table 4, k₂ decreases with increasing ligand electron-donating ability. This is consistent with the proposed reaction mechanism involving an electrophilic peroxo group. Strong electron-donating groups provide more electron density to the peroxo core through bonding to the copper centers, which should disfavor the electrophilic attack of the peroxo species. The enhanced electrophilic nature of a μ-η²:η²-peroxo group relative to an end-on peroxo-dicopper(II) structure (see Scheme 2) is supported by theoretical studies[30,34], where the peroxide has two antibonding interactions with each of the two copper ions, resulting in a much less negative peroxide due to extremely large r-donor interactions between peroxide and copper. The electrophilic behavior of the side-on peroxo moiety has also been demonstrated in a detailed study comparing reactivity patterns of the side-on versus end-on peroxo dicopper(II) species [35]. However, theoretical calculations on the electronic structure of side-on peroxodicopper(II) moieties show that the bonding interaction between the Cu d_x2-y2 orbitals and the O₂²⁻ σ* orbital facilitate the back-donation of electron-density which is responsible for the weak m(O-O) stretch (~740 cm ⁻¹) in side-on peroxodicopper(II) species compared to the end-on analogues (~820 cm ⁻¹).

In summary, the proposed reaction mechanism for the endogenous arene hydroxylation reaction effected by the side-on peroxo complexes with XYL and related ligands are further confirmed by the kinetic/thermodynamic studies described here. Strong electron-donating ligands afford the formation of more electron-rich peroxo moieties, which disfavor their electrophilic arene hydroxylation (i.e. smaller rate constants k₂ Scheme 1). Ligand constraints in the unsymmetrical UN ligand results in a significantly slower hydroxylation due to the disruption of the ideal orientation of the peroxo core and arene substrate present in the symmetrical ligand systems.

4. Conclusions

The effect of ligand electronic variations on the dioxygen binding and the arene hydroxylation reactions of the resulting Cu₂O₂ species has been achieved by incorporating 4-pyridyl substituents in the extensively studied XYL system. The detailed kinetic/thermodynamic studies of the oxygenation reactions of the corresponding copper(I) complexes with substituted XYL ligands reveal that the more electron-donating ligand favors O₂-binding to form the peroxodicopper(II) complexes [Cu^II₂(L)(O₂²⁻)]²⁺ (2). However, no direct correlation between the ligand electron-donating ability and the stability of the peroxo complexes is observed, which is rationalized by some interaction between the Cu₂O₂ core and the xylyl π-orbitals. Less favorable arene hydroxylation of the peroxo complexes with strong electron-donating ligands is consistent with the proposed reaction mechanism involving electrophilic attack of the peroxo moiety on the xylyl ring.

Ligand structural effects are also investigated utilizing the analogous but unsymmetrical UN ligand. The kinetic and thermodynamic data show that the ligand constraints have significant effect on the formation, stability and reactivity of the peroxo complex. A significantly slower thermal arene hydroxylation reaction is observed for the UN complex compared with the XYL complex, suggesting disruption of the ideal positioning of the peroxo core towards the arene substrate.

The present detailed systematic investigations provide a wealth of new chemistry and information concerning how ligand electronic and structural elements affect copper-dioxygen interactions as well as C-H activation, i.e. biomimetic hydroxylation mediated by copper-dioxygen chemistry. Such fundamental studies help to unravel the detailed nature of copper-dioxygen reactivity and provide a basis for a deeper understanding of relevant biological systems as well as potential practical applications for copper mediated oxidative chemistry.

Appendix A. Supplementary material

CCDC 867041 contains the supplementary crystallographic data for complex [Cu^II₂(^MeOXYLO⁻)(⁻OH)](PF₆)₂ (3b). These data can be obtained free of charge from The Cambridge Crystallographic Data Centre via www.ccdc.cam.ac.uk/data_request/cif. Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.ica.2012.01.042.