Synthesis and Characterization of Copper ... - ACS Publications

Oct 26, 2016 - Douglas R. Powell,. ‡. Felio Perez,. § ... Integrated Microscopy Center, University of Memphis, Memphis, Tennessee 38152, United Sta...
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Synthesis and Characterization of Copper Complexes with CuICuI, Cu1.5Cu1.5m and CuIICuII Core Structures Supported by a Flexible Dipyridylamide Ligand Ethan P. McMoran,† Douglas R. Powell,‡ Felio Perez,§ Gerard T. Rowe,*,∥ and Lei Yang*,† †

Department of Chemistry, University of Central Arkansas, Conway, Arkansas 72035, United States Department of Chemistry and Biochemistry, University of Oklahoma, Norman, Oklahoma 73019, United States § Integrated Microscopy Center, University of Memphis, Memphis, Tennessee 38152, United States ∥ Department of Chemistry and Physics, University of South CarolinaAiken, Aiken, South Carolina 29801, United States ‡

S Supporting Information *

ABSTRACT: A series of copper complexes supported by a simple dipyridylamide ligand (H2pcp) were isolated and characterized. Treatment of H2pcp with NaH and copper(I) salts led to the formation of [Cu2(2pcp)2] (1a) and {Na[(Cu2(2pcp)2)2]PF6}n (1b). The X-ray crystal structures of both complexes feature CuICuI cores with close Cu···Cu interactions. Electrochemical studies of 1a showed a reversible one-electron oxidation wave in CH2Cl2. On the basis of the work on 1a, we began studying the mixed-valence copper species supported by this ligand. The reaction of H2pcp with Cu(OAc)2 and CuCl in different stoichiometries yielded [Cu2(2pcp)2Cl] (2) and [Cu3(2pcp)2Cl2] (3). X-ray crystallography and spectroscopic characterization suggested delocalized Cu1.5Cu1.5 core structures of both compounds. These results further inspired us to explore the coordination properties of H2pcp toward CuII ions. The complexes [HNEt3][Cu2(2pcp)3(ClO4)](ClO4) (4a), [Cu2(2pcp)3(NO3)] (4b), and [Cu2(2pcp)3(H2O)]BF4 (4c) featuring dinuclear CuIICuII cores were prepared and characterized by X-ray crystallography and spectroscopic methods. Structural analysis of these complexes implied that the accommodation of CuICuI, Cu1.5Cu1.5, and CuIICuII is attributed to the structural flexibility of the ligand H2pcp. Complexes 1a, 2, 3, and 4a were examined by X-ray photoelectron spectroscopy, which confirmed the oxidation state assignments. Computational studies were also performed to provide insight into the electronic structures of these complexes.

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INTRODUCTION The presence of mixed-valence copper clusters in biological systems is of particular interest because of their significant function in electron-transfer processes and catalysis. For instance, the two copper centers at the CuA site switch their oxidation states between CuICuI and delocalized mixed-valence Cu1.5Cu1.5 to transfer electrons in enzymes such as cytochrome c oxidase and nitrous oxide reductase.1 In the active sites of particulate methane monooxygenase (pMMO) and laccase, localized mixed-valence copper clusters with 1+ and 2+ oxidation states are proposed as intermediates in the catalytic process based on recent mechanistic and experimental studies.2 In various efforts to understand the mechanisms of these copper-containing enzymes, small structural and functional models have been developed and characterized. The reported examples based on nitrogen-donor ligands such as amine,3 pyrazolate,4 octaazacrytand,5 amidate,6 β-diketiminate,7 and amidinate8 systems have offered great insight. One major focus in these modeling studies is to manipulate the ligands to support copper centers continually when oxidation states change. Because copper(I) and copper(II) prefer different geometries and valence intermediates need certain coordination © XXXX American Chemical Society

environments to be stabilized, it still remains challenging to have a ligand platform that is flexible enough to accommodate copper ions during the oxidation/reduction process. As part of our ongoing research on biologically related copper complexes, we selected pyridylamide ligands as the supporting scaffolds for copper ions. Our previous work demonstrated the excellent coordination ability of this system and the interesting noncovalent interaction induced into the outer-sphere coordination environment.9 Herein, we report a simple dipyridylamide ligand that can support CuICuI, Cu1.5Cu1.5, and CuIICuII dimeric units. Experimental evidence demonstrated that the tridentate coordination mode and the coordinative flexibility of this ligand are the major contributing factors in the stabilization of these clusters. To our best knowledge, the only ligand system with such a function is phosphinophenylamido/phosphido.10 The complexes prepared in this work were fully characterized by X-ray crystallography, UV−vis, electron paramagnetic resonance (EPR), cyclic voltammetry, and X-ray photoelectron microscopy (XPS). In Received: August 17, 2016

A

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

was stirred for 2 h, and the solvent was removed under vacuum. The resulting brown powder was washed with Et2O (5 mL × 3). The powder was extracted with 2 mL of acetone, and the turbid solution was filtered. Vapor diffusion of Et2O into the brown filtrate at room temperature led to the formation of dark-brown crystals suitable for Xray crystallographic characterization (0.031 g, 41% yield). Anal. Calcd for C44H32N12O4PF6NaCu4: C, 43.50; H, 2.65; N, 13.83. Found: C, 43.22; H, 2.47; N, 13.72. FT-IR (cm−1): 1619, 1586, 1558, 1471, 1432, 1375, 1295, 1161, 1106, 1049, 1008, 830, 776, 751, 739, 698, 634, 617, 558, 514, 449, 426, 407. 1H NMR (DMF-d7, 298 K): 8.53 (m, 4H), 8.38 (m, 2H), 8.13 (m, 2H), 8.01 (m, 2H), 7.90 (m, 2H), 7.70 (m, 2H), 7.11 (m, 2H). UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 457 nm (shoulder, 1370). [Cu2(2pcp)2Cl] (2). Cu(OAc)2 (0.012 g, 0.065 mmol) was added to 1 mL of a H2pcp solution (0.026 g, 0.13 mol) in 2 mL of DMF, and the green mixture was stirred for 2 h. CuCl (0.006 g, 0.065 mmol) was then added to the mixture, and the dark-brownish-green mixture was stirred for 1 h. A total of 15 mL of Et2O was added to the mixture. The resulting deep-brown powder was collected and washed with Et2O (5 mL × 3). The powder was extracted with 2 mL of DMF, and the solution was filtered. Vapor diffusion of Et2O into the dark-browngreen filtrate at room temperature led to the formation of dark-brown crystals suitable for X-ray crystallographic characterization (0.020 g, 55% yield). Anal. Calcd for C22H16ClN6O2Cu2: C, 47.27; H, 2.89; N, 15.04. Found: C, 47.11; H, 2.80; N, 14.95. FT-IR (cm−1): 2360, 1678, 1623, 1586, 1469, 1431, 1385, 1355, 1293, 1270, 1153, 1088, 1047, 1016, 938, 881, 821, 781, 765, 711, 692, 675, 651, 618, 555, 542, 504, 464, 455. UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 566 (570), 715 (320). Solution magnetic moment (Evans’ method, 22.0 °C, 7.16 × 10−3 M, CD2Cl2): 1.86 μB. EPR (9.434 GHz, solution in CH2Cl2, 298 K): g = 2.178 and A = 50 G. EPR (9.434 GHz, frozen solution in CH2Cl2/tolunene, 77 K): g1 = 2.112, g2 = 2.251, g3 = 2.266, A1 = 15 G, A2 = 55 G, and A3 = 133 G. [Cu3(2pcp)2Cl2] (3). Cu(OAc)2 (0.012 g, 0.065 mmol) was added to 1 mL of a H2pcp solution (0.026 g, 0.13 mol) in DMF, and the green mixture was stirred for 1.5 h. CuCl (0.012 g, 0.13 mmol) was added to the mixture, and the dark-brownish-green mixture was stirred for 1 h. A total 15 mL of Et2O was added to the mixture. The resulting deepbrown powder was collected and washed with Et2O (5 mL × 3). The powder was extracted with 2 mL of DMF, and the solution was filtered. Vapor diffusion of Et2O into the dark-brown-green filtrate at room temperature led to the formation of dark-brown crystals suitable for X-ray crystallographic characterization (0.024 g, 55% yield). Anal. Calcd for C22H16Cl2N6O2Cu3: C, 40.16; H, 2.45; N, 12.77. Found: C, 40.28; H, 2.39; N, 12.84. FT-IR (cm−1): 2359, 1623, 1596, 1559, 1471, 1433, 1344, 1294, 1250, 1154, 1090, 1049, 1014, 939, 871, 819, 773, 760, 743, 711, 692, 642, 620, 542, 503, 466, 450, 433. UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 551 (600), 719 (460). Solution magnetic moment (Evans’ method, 22.0 °C, 4.72 × 10−3 M, CD2Cl2): 1.69 μB. EPR (9.434 GHz, solution in CH2Cl2, 298 K): g = 2.167 and A = 47 G. EPR (9.434 GHz, frozen solution in CH2Cl2/toluene, 77 K): g1 = 2.132, g2 = 2.165, g3 = 2.274, A1 = 5 G, A2 = 45 G, and A3 = 103 G. [HNEt3][Cu2(2pcp)3(ClO4)](ClO4) (4a). A solution of Cu(ClO4)2· 6H2O (0.062 g, 0.17 mmol) in 1 mL of CH3CN was added to a stirring mixture of H2pcp (0.050 g, 0.25 mmol) and Et3N (0.025g, 0.25 mmol) in 2 mL of CH3CN. The deep-green solution was stirred for 2 h, and 20 mL of Et2O was added to the mixture. The resulting green powder was collected and washed with Et2O (5 mL × 3). The powder was dissolved in 2 mL of CH3CN and filtered. Vapor diffusion of Et2O into the deep-green solution at room temperature led to the formation of dark-green crystals suitable for X-ray crystallographic characterization (0.062 g, 79% yield). Anal. Calcd for C39H40N10O11Cl2Cu2: C, 45.80; H, 3.94; N, 13.69; Cl, 6.93. Found: C, 45.40; H, 3.96; N, 13.53; Cl, 6.82. FT-IR (cm−1): 1700, 1685, 1622, 1590, 1559, 1525, 1508, 1474, 1436, 1367, 1298, 1253, 1144, 1113, 1087, 1027, 941, 776, 708, 692, 627. UV−vis [CH3CN; λmax, nm (ε, M−1 cm−1)]: 661 (230), 751 (240). UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 632 (260), 761 (250). Solution magnetic moment (Evans’ method, 22.0 °C, 5.52 × 10−3 M, CD2Cl2): 1.79 μB per

addition, theoretical work was performed to investigate the electronic structures of these compounds.

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EXPERIMENTAL SECTION

General Procedures. All reagents were obtained from commercial sources and used as received unless otherwise noted. The ligand 2-[N(2-pyridyl)carbamoyl]pyridine (H2pcp) was synthesized according to the literature.11 For the synthesis of complexes 1a, 1b, 2, and 3 in a glovebox, organic solvents were dried over molecular sieves, CaH2 or sodium/benzophenone, and distilled under vacuum prior to use. 1H NMR spectra were recorded on a Bruker AVANCE 300 MHz spectrometer at room temperature. Chemical shifts (δ) were referenced to the residual solvent signal. UV−vis spectra were recorded on a Cary 50 spectrometer. X-band EPR spectra were recorded on a Freiberg Instruments MiniScope MS5000 spectrometer. The g factors and A coupling constants of EPR spectra were obtained from simulation using Bruker SimFonia software (version 1.25). Elemental analyses were carried out by Atlantic Microlabs, Norcross, GA. Fourier transform infrared (FT-IR) spectra were collected on a Nicolet Magna 560 FT-IR spectrometer with an attenuated-totalreflectance attachment. Solution magnetic susceptibility was measured by using Evans’ method on a Bruker AVANCE 300 MHz spectrometer at room temperature. Cyclic voltammograms were measured using an EG&G Princeton Applied Research scanning potentiostat with a threeelectrode cell containing a platinum disk working electrode, a platinum wire auxiliary electrode, and a Ag/AgCl glass reference electrode. All measurements were performed in organic solutions containing 1.0 mM analyte and 0.1 M NBu4BF4 at room temperature with dinitrogen protection. Recrystallized ferrocene was used as the internal standard. SQUID data were collected by using a Quantum Design MPSM 3 instrument. Caution! The perchlorate salt Cu(ClO4)2·6H2O is potentially explosive. Although we did not experience any problems while conducting the work, it should be handled with great caution. For XPS measurements, analysis of the samples was performed using a Thermo Scientific K-Alpha XPS system equipped with a monochromatic X-ray source at 1486.6 eV, corresponding to the Al Kα line. The X-ray power of 75 W at 12 kV was used for all experiments with a spot size of 400 μm2. The base pressure of the KAlpha instrument was at 1.0 × 10−9 mbar. The instrument was calibrated to give a binding energy of 84.0 eV for Au 4f7/2 and 284.8 eV for the C 1s line of adventitious (aliphatic) carbon present on the nonsputtered samples. Photoelectrons were collected from a takeoff angle of 90° relative to the sample surface. Measurements were done in the constant analyzer energy mode. The survey spectra were taken at a pass energy of 200 eV, while the high-resolution (HR) core-level spectra were taken at a 40 eV pass energy with an energy step size of 0.1 eV and an average of 30 scans. The XPS data acquisition was performed using the “Avantage v5.932” software provided with the instrument. [Cu2(2pcp)2] (1a). CuCl (0.025 g, 0.25 mmol) was added to a stirring mixture of H2pcp (0.050 g, 0.25 mmol) and NaH (0.006 g, 0.25 mmol) in 2 mL of N,N-dimethylformamide (DMF). The brown slurry was stirred for 1 h, and 20 mL of Et2O was added to the mixture. The resulting brown powder was collected and washed with Et2O (5 mL × 3). The powder was extracted with 2 mL of DMF, and vapor diffusion of diethyl ether (Et2O) into the brown filtrate at room temperature led to the formation of dark-brown crystals suitable for Xray crystallographic characterization (0.033 g, 50% yield). Anal. Calcd for C22H16N6O2Cu2: C, 50.48; H, 3.08; N, 16.05. Found: C, 50.26; H, 2.95; N, 16.31. FT-IR (cm−1): 1614, 1581, 1553, 1540, 1464, 1425, 1392, 1343, 1287, 1147, 1094, 1046, 1004, 929, 896, 879, 817, 778, 756, 738, 709, 691, 618, 601, 545, 503, 432, 409. 1H NMR (DMF-d7, 298 K): 8.54 (m, 4H), 8.39 (m, 2H), 8.17 (m, 2H), 8.01 (m, 2H), 7.90 (m, 2H), 7.71 (m, 2H), 7.12 (m, 2H). UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 457 nm (shoulder, 1400). {Na[(Cu2(2pcp)2)2]PF6}n (1b). [Cu(NCCH3)4]PF6 (0.025 g, 0.25 mmol) was added to a stirring mixture of H2pcp (0.050 g, 0.25 mmol) and NaH (0.006 g, 0.25 mmol) in 2 mL of CH2Cl2. The brown slurry B

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

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09, revision C.01,17 at the M06L18/def2-TZVP19 level of theory. Dunlap’s density-fitting approximation was employed for all geometry optimization calculations, with the auxiliary basis sets generated automatically by the program.20 The identification of each compound’s structural minimum was confirmed through the absence of imaginary vibrational modes in an analytical frequency calculation. For the optimization of 4a structural variants, 4−CH3CN, 4−CH2Cl2, and 4− CH3CN, implicit solvation was included through the use of COSMO.21 The XPS spectra of 1a, 2, 3, and 4a were interpreted by analysis of their Cu 2p orbitals using Orca 3.0.3 for Linux.22 These calculations were carried out at the M06L/def2-TZVP level of theory with the exception of copper atoms, for which the polarized CoreProp basis set [CP(PPP)] was utilized.23,24 A higher accuracy integral grid (Lebedev 770 point grid) was also utilized for the copper atoms. For the openshell molecules 2, 3, and 4a, the restricted open-wave function was calculated in order to allow for easier interpretation of orbital energies. Using Orca 3.0.3, a broken-symmetry calculation was performed on 4a in order to predict the J-coupling value for the dicupric system. Using the def2-TZVP basis set, it was found that the Coulombattenuated functional cam-B3LYP25 gave the best performance. To calculate the numerical value of J, the generalized equation26

copper. EPR (9.434 GHz, solution in CH3CN, 298 K): g = 2.175 and A = 76 G. [Cu2(2pcp)3(NO3)] (4b). Complex 4b was synthesized in methanol (MeOH) by following a procedure similar to that of 4a. Cu(NO3)2· 3H2O was used as the copper(II) salt. Vapor diffusion of Et2O into the deep-green solution at room temperature led to the formation of darkgreen crystals suitable for X-ray crystallographic characterization (0.049 g, 48% yield). Anal. Calcd for C33H24N10O6Cu2: C, 50.57; H, 3.09; N, 17.87. Found: C, 50.68; H, 3.06; N, 17.62. FT-IR (cm−1): 1627, 1593, 1562, 1470, 1433, 1348, 1298, 1253, 1156, 1093, 1047, 1025, 939, 854, 815, 762, 708, 691, 652, 620, 534, 506, 465, 435, 409. UV−vis [CH3CN; λmax, nm (ε, M−1 cm−1)]: 655 (240), 739 (230). UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 637 (250), 749 (280). Solution magnetic moment (Evans’ method, 22.0 °C, 5.24 × 10−3 M, CD2Cl2): 1.78 μB per copper. [Cu2(2pcp)3(H2O)]BF4 (4c). Complex 4c was synthesized in MeOH by following a procedure similar to that of 4a. Cu(BF4)2·6H2O was used as the copper(II) source. Vapor diffusion of Et2O into the deepgreen solution at room temperature led to the formation of dark-green crystals suitable for X-ray crystallographic characterization (0.069 g, 67% yield). Anal. Calcd for C33H26N9O4BF4Cu2: C, 47.95; H, 3.17; N, 15.25. Found: C, 47.89; H, 3.31; N, 15.42. FT-IR (cm−1): 1621, 1591, 1561, 1469, 1434, 1364, 1300, 1253, 1155, 1055, 1018, 939, 854, 816, 761, 709, 690, 653, 620, 510, 463, 434, 409. UV−vis [CH3CN; λmax, nm (ε, M−1 cm−1)]: 658 (220), 747 (230). UV−vis [CH2Cl2; λmax, nm (ε, M−1 cm−1)]: 634 (250), 757 (240). Solution magnetic moment (Evans’ method, 22.0 °C, 6.26 × 10−3 M, CD2Cl2): 1.75 μB per copper. Magnetic Susceptibility of 4a and Magnetic Data Fitting. The magnetic susceptibility of 4a was measured using a Quantum Design MPMS SQUID VSM Evercool magnetometer. The sample was loaded into a polypropylene VSM powder holder. The magnetic susceptibility (zero-field-cooled) was measured with an applied field of 1000 Oe as a function of the temperature between 2 and 300 K. The methods of Morrison and zur Loye12 were used to correct for shape and radial offset effects. The magnetic susceptibility temperature dependence in 4a was modeled using the Bleaney−Bowers equation13 for an isolated spin− spin dimer whose magnetic centers have a coupling constant, J: χmol of Cu =

2NAg 2μB2 k bT (3 + e

−J / k bT

)

+

LS

Jab =

E(DFT ) − HSE(DFT) ⟨S ⟩(DFT) − LS⟨S2⟩(DFT)

HS

2

recommended by Soda et al. was employed.27 Using Orca 3.0.3, gas-phase electronic spectra of 1a, 2, and 3 were calculated with the double-hybrid RI-B2PLYP functional.28 A mixture of basis sets (def2-SVP19,29 for carbon, hydrogen, and oxygen; def2TZVP for nitrogen, chlorine, and copper) was employed, and the appropriate correlation auxiliary basis sets (i.e., def2-SVP/c and def2TZVP/c) were used in the RI-MP2 portion of the calculations. Because of the large size of 4a making the B2PLYP calculation intractable, its electronic spectrum was calculated at the B3LYP/def2TZVP level of theory. Selected excited states for these species were visualized as hole−particle pairs using natural transition orbitals (NTOs).30 NTO analysis allows for the condensation of a transition that has many contributions into a small number of hole−particle pairs that describe the excitation process. All molecular orbital images were generated using IboView for Windows.31 Results and Discussion. Synthesis and Characterization of 1a and 1b. The synthesis of 1a is shown in Scheme 1. Treatment of the

Ceff + χTIP T−θ

Scheme 1. Synthesis of 1a

where the second term represents the effective Curie−Weiss law for an impurity of monomeric material and χTIP is temperature-independent paramagnetic susceptibility. Minimization of the reliability factor R [where R = ∑(χmTcalc − χmTobs)2/(χmTobs)2] between temperatures of 2 and 300 K led to the best fit for J = −29 cm−1, and a magnetic field splitting factor g = 1.93, with R = 2.70 × 10−3. X-ray Crystallography. X-ray crystallographic data were collected on crystals with dimensions of 0.56 × 0.39 × 0.37 mm for 1a, 0.39 × 0.04 × 0.02 mm for 1b, 0.16 × 0.12 × 0.10 mm for 2, 0.08 × 0.06 × 0.06 mm for 3, 0.50 × 0.15 × 0.04 mm for 4a, 0.52 × 0.40 × 0.18 mm for 4b, and 0.32 × 0.14 × 0.12 mm for 4c. Data were collected at 100 K using a diffractometer with a Bruker APEX CCD area detector14 and graphite-monochromated Mo Kα radiation (λ = 0.71073 Å). All seven structures were solved by direct methods and refined by full-matrix least-squares methods on F2.15 The crystal parameters, data collection, and refinement of 1a, 2, 3, and 4a are summarized in Table S1. The crystal parameters, data collection, and refinement of 1b, 4b, and 4c are summarized in Table S2. Non-hydrogen atoms were refined with anisotropic displacement parameters. Hirshfeld Surface Analysis. In order to investigate the important intra- and intermolecular interactions occurring within and among the coordination units, analysis of the Hirshfeld surface was performed by using CrystalExplorer (version 3.0) software.16 The calculation was based on the CIF files from the X-ray crystal structures of the metal complexes. Computational Chemistry. Starting with X-ray structures, the optimized geometries of 1a, 2, 3, and 4a were calculated with Gaussian

ligand H2pcp with NaH in DMF yielded a white suspension. The addition of CuCl turned the suspension to a dark-red solution, which yielded a deep-brown powder when Et2O was added. Purification of the powder through recrystallization afforded brown crystals suitable for X-ray diffraction. Solutions of 1a in organic solvents were airsensitive and turned to light-green solutions in a few minutes when exposed to air. The X-ray crystal structure of 1a is shown in Figure 1. Also, see Table 1 for selected bond lengths and angles. Two copper(I) centers are sandwiched by two deprotonated ligands, which display a μ2(κNpy,Namidate;κNpy) coordination mode. For each copper(I) ion, the coordination environment is fulfilled by two nitrogen donors from one ligand and the third nitrogen from the second ligand. The two copper(I) centers have a distance of 2.578 Å that is shorter than the sum of the copper covalent radii (2.64 Å),32 indicating a weak interaction. The dimeric unit has a 2-fold rotation axis that crosses the C

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

Figure 1. X-ray structures of one dimeric unit (left) and two packing units (right) of 1a. Thermal ellipsoids are shown at the 50% level. Most hydrogen atoms are omitted for clarity. Cu−Cu = 2.578 Å. Synthesis and Characterization of 2 and 3. Inspired by the interesting electrochemical behavior of 1a, we attempted to isolate the mixed-valence compounds supported by the ligand H2pcp. In the absence of base, the treatment of 1 equiv of H2pcp with 0.5 equiv of Cu(OAc)2, followed by the addition of either 0.5 or 1 equiv of CuCl, led to the isolation of mixed-valence complexes 2 and 3 in DMF (Scheme 2). Metalation of the amidate group in the ligand H2pcp

Table 1. Selected Bond Lengths (Å) and Angles (deg) of Complex 1a Cu1−N1A Cu1−N3 Cu1−N2 Cu1−Cu1A

1.9180(13) 2.2202(15) 1.9309(14) 2.5780(4)

N1A−Cu1−N2 N1A−Cu1−N3 N2−Cu1−N3 N1A−Cu1−Cu1A N2−Cu1−Cu1A N3−Cu1−Cu1A

169.06(6) 109.40(5) 81.48(5) 85.62(4) 84.89(4) 152.23(4)

Scheme 2. Synthesis of 2 and 3

midpoint of the Cu−Cu bond. In the packing structure, two dimeric units overlap in a face-to-face fashion with a centroid-to-centroid distance of 3.790 Å (Figure 1, right). In fact, the Hirshfeld surface analysis based on one dimeric unit does show the distinct blue and red triangle pairs on the map plotted by the surface index, indicating the presence of intermolecular π···π interactions between the aromatic pyridine rings (Figure S1). One interesting structural feature of 1a is the saddle-shaped conformation in which each ligand twists its two pyridyl rings with a torsion angle of 27.98°. Initially, we thought the intermolecular π···π interactions might be a factor for this unique molecular shape, but the geometry optimization calculations revealed that the bent shape was also stable in the gas phase. Therefore, the intermolecular π···π interactions in the crystal structure of 1a are not necessary to maintain the geometry. When [Cu(NCCH3)4]·PF6 was used in the synthesis, onedimensional polymer 1b with the same dimeric unit as 1a was isolated through a similar procedure. The X-ray crystal structure shows that two pairs of dimeric units are connected by one Na+ cation with a tetrahedral geometry (τ4 = 0.76) that is fulfilled by four oxygen donors from the CO groups. According to X-ray structure data, the average CO bond in 1b is 1.246 Å, which is slightly longer than the CO bond (1.234 Å) in 1a. This elongation of the CO bond might be attributed to the delocalization of electrons through the κ2-amidate group to stabilize Na+ cations. Each Na+ cation is balanced by one PF6− anion, which interacts with one C−H group on the onedimensional chain through an intermolecular hydrogen bond (H···F = ∼2.405 Å; Figure S2 and Table S3). The orange-red solution of 1a in CH2Cl2 shows a shoulder around 460 nm with a coefficient of 1400 M−1 cm−1, suggesting a chargetransfer property between the deprotonated ligand and the copper(I) center (Figure S3). The calculated UV−vis spectra (Figure S4) and NTO analysis (Figure S5) further confirmed the primary ligand-tometal charge-transfer character of this band. Cyclic voltammetry measurement was performed to investigate the oxidation/reduction of the copper centers (Figure S6). A reversible wave was observed at E1/2 = 0.095 V vs Ag/AgCl with a scan rate of 100 mV/s. ΔEp (65 mV) is independent of the scan rates, indicating an electrochemical reversible property of this wave. The feature suggests the formation of a CuIICuI mixed-valence species upon oxidation of the CuICuI unit.

without a base has been observed in our previous work.9b The resulting dark-brownish-green crystals of both complexes are quite stable in air for months, but the solutions in organic solvents turned to green within ∼10 min in air. X-ray crystallography analysis revealed a dimeric structure in 2 (Figure 2, top) and a trinuclear cluster in 3 (Figure 2, bottom), with selected bond lengths and angles listed in Table 2. In the X-ray crystal structure of 2, a dinuclear core structure similar to that of complex 1 was observed. A chloride ligand sits above the dimeric unit and interacts with both copper centers through a distance of 2.5446(19) Å. Because the two negatively charged amidate ligands and one chloride have to be balanced by two copper centers, the total charge of the two copper ions should be 3+. Interestingly, the Cu1− Cu1A distance in 2 [2.3577(15) Å] is significantly shorter than the Cu−Cu distance in 1 [2.5780(4) Å], indicating a stronger interaction between the copper centers. These features suggest a delocalized mixed-valence (Cu1.5Cu1.5) property of 2, which is confirmed by EPR, UV−vis, and XPS evidence (see the following discussion). The X-ray crystal structure of 3 is featured by a trinuclear copper core structure with a trianglar arrangement of the three copper centers. D

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

[Cl1−Cu2−Cl1A = 177.59(5)°] is fulfilled by two chlorides with a Cu−Cl bond of 2.1176(9) Å. On the basis of the charge balance (two negatively charged amidate ligands and two chlorides), geometries of the copper centers, and spectroscopic studies, we assign the Cu1 and Cu1A dimeric structure as a delocalized mixed-valence (Cu1.5Cu1.5) unit, while Cu2 as a copper(I) center. As demonstrated by the spectroscopic evidence and theoretical calculations (see the following discussion), the [CuCl2]− motif in 3 does not have any significant impact on the electronic structure of the delocalized mixed-valence dimeric unit. Therefore, the formation of the trimeric structure is attributed to the electrostatic attraction between the negatively charged [CuCl2]− unit and positively charged [Cu2(2pcp)2]+ unit. A close examination of the Hirshfeld surface map plotted with dnorm for 3 revealed an interesting feature shown in Figure S7. Two bright-red dots are observed on the surface map of the [Cu2(2pcp)2]+ unit, indicating weak Cu1···Cl1 and Cu1A···Cl1A interactions. This feature is confirmed by the large Cu1···Cl1 and Cu1A···Cl1A distances (2.810 Å) in the X-ray crystal structure. Apparently, as a result of the weak interactions, the [CuCl2]− unit is able to anchor on top of the [Cu2(2pcp)2]+ unit and nearly align with the Cu1−Cu1A bond (torsion angle of Cl1−Cl1A−Cu1A−Cu1 = 12.84°). UV−vis spectra of complexes 2 and 3 in a CH2Cl2 solution were collected to understand their electronic structures (Figure S8). Both compounds show similar bands around 560 and 710 nm, indicating similar solution structures. Much like the experimental data, the calculated spectra of the mixed-valence complexes 2 and 3 are similar to each other (Figure S9). The spectrum of 3 contains all of the major features seen in that of 2 but is blue-shifted by about 70 nm. This similarity in their electronic structures is more evidence for our description of the [CuCl2]− group being electrostatically bound rather than participating in a discrete metal−metal bond. NTO analysis of the transitions in 2 and 3 support this assignment, as well. The similarity between the two compounds is illustrated particularly well in the representation of the corresponding absorption features of 2 at 583 nm and 3 at 507 nm (Figure 3). The nature of these two transitions is nearly identical, with a majority of the transitions having d → d character and a smaller contribution that have Cu−Cu σ → σ* character. The only participation that the [CuCl2]− group has in the 507 nm transition of 3 is as part of the ligand orbitals that make up the hole of the minor σ → σ* component. Despite this participation, the overall nature of the hole orbital is not fundamentally changed compared to that of 2. A similar situation is observed in the lowestenergy transitions of 2 and 3, in which the features have strong [Cu1.5Cu1.5] → Cl− and [Cu1.5Cu1.5] → [CuCl2]− character, respectively (Figures S10 and S11). X-band EPR spectra of 2 and 3 were collected in CH2Cl2 solutions (room temperature; Figure 4) and frozen CH2Cl2/toluene solutions [1:1 (v/v), 77 K; Figure 5]. The seven-line hyperfine signals observed at both temperatures suggest that an unpaired electron delocalizes over two copper nuclei with I = 3/2, confirming the assigned delocalized mixed-valence Cu1.5Cu1.5 states based on the X-ray crystal structures. Simulation of experimental spectra gave close g values and A coupling constants, again suggesting the similar solution structures of the two complexes. The solid-state EPR spectra of the two complexes at room temperature were also collected, but unfortunately no hyperfine features were observed (Figure S12). From the above discussion, we demonstrated the similar structural and spectroscopic characteristics of 2 and 3. Theoretical work was also performed to better understand the electron structure of the two compounds. The identical Mulliken atomic charges and spin densities of the two copper atoms in 2 (Cu1, charge = 0.1474, spin = 0.3477; Cu1A, charge = 0.1474, spin = 0.3477) and the analogous copper centers in 3 (Cu1, charge = 0.0339, spin = 0.3218; Cu1A, charge = 0.0339, spin = 0.3218) strongly suggest a delocalized state. The singly occupied molecular orbitals (SOMOs) of 2 and 3 derived from UNO analysis allow for an intuitive visualization of their unpaired electron densities, as seen in Figure 6. The UNO visualization highlights the similarity of the frontier orbitals of these two compounds, and the high symmetry of these orbitals emphasizes the extent of unpaired electron

Figure 2. X-ray structures of complexes 2 (top) and 3 (bottom). Thermal ellipsoids are shown at the 50% level. Hydrogen atoms are omitted for clarity. Complex 2: Cu1···Cu1A = 2.358 Å. Complex 3: Cu1−Cu1A = 2.436 Å; Cu1···Cu2 = 2.867 Å; Cu1A···Cu2 = 2.867 Å.

Table 2. Selected Bond Lengths (Å) and Angles (deg) of Complexes 2 and 3 Complex 2 Cu1−N1 Cu1−N2A Cu1−N3A Cu1−Cl1 Cu1−Cu1A

1.938(5) 1.934(5) 2.046(4) 2.5446(19) 2.3577(15)

N1−Cu1−N2A N1−Cu1−Cu1A N1−Cu1−N3A N1−Cu1−Cl1 N2A−Cu1−N3A N2A−Cu1−Cu1A N2A−Cu1−Cl1 N3A−Cu1−Cu1A N3A−Cu1−Cl1 Cu1−Cl1−Cu1A

151.15(19) 87.63(13) 107.46(19) 108.96(14) 83.34(19) 88.26(15) 94.23(13) 161.62(14) 101.852(13) 84.89(4)

1.938(5) 1.934(5) 2.046(4) 2.4356(8) 2.8763(8) 2.8763(8) 2.1176(9) 2.1176(9)

N1−Cu1−N2A N1−Cu1−Cu1A N1−Cu1−N3A N1−Cu1−Cu2 N2A−Cu1−N3A N2A−Cu1−Cu1A N2A−Cu1−Cu2 N3A−Cu1−Cu1A N3A−Cu1−Cu2 Cu1−Cu2−Cu1A Cl1−Cu2−Cl1A

169.23(10) 85.67(8) 102.02(10) 79.59(7) 83.35(10) 85.51(7) 102.16(7) 153.33(6) 141.35(6) 50.27(2) 177.59(5)

Complex 3 Cu1−N1 Cu1−N2A Cu1−N3A Cu1−Cu1A Cu1···Cu2 Cu1A···Cu2 Cu2−Cl1 Cu2−Cl1A

In addition to the Cu1−Cu1A dimeric unit similar to that of 2, complex 3 has a third copper center Cu2 that is located right above the Cu1−Cu1A bond. Cu1 and Cu1A interact with each other through a distance of 2.4356(8) Å, which is slightly longer than the Cu1−Cu1A bond in 2 [2.3577(15) Å] but is still shorter than the Cu1−Cu1A bond in 1 [2.5780(4) Å]. For the Cu2 center, the linear geometry E

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

assignment of these compounds as delocalized mixed-valence Cu1.5Cu1.5 species with either a bridging chloride or a [CuCl2]− group. Cyclic voltammetry characterization of 2 and 3 was performed in CH2Cl2 solutions with dinitrogen protection. The two complexes showed reversible oxidative waves around E1/2 = −0.11 V (complex 2; Figure S13, left) and E1/2 = −0.095 V (complex 3; Figure S14, left), both of which are assigned as the CuICuI/CuIICuI couple of the dimeric unit. ΔEp (264 mV in 2 and 256 mV in 3) of the waves showed no significant change with different scan rates. In comparison to the CuICuI/CuIICuI wave of 1a (E1/2 = 0.09 V), the lower oxidation potentials of 2 and 3 might result from the presence of an electrondonating chloride ligand. In addition to the CuICuI/CuIICuI couple, 2 showed an irreversible oxidation wave at E = 0.79 V (Figure S13, right). 3 also showed a reversible oxidation wave at E1/2 = 0.82 V (Figure S14, right), which gradually became irreversible at low scan rates. Both waves are assigned as the CuIICuII state formed from the oxidation of CuIICuI core structures. Because of coordination of the chloride ligand to the copper center in 2, it is possible that structural reorganization during oxidation of the copper ions leads to the formation of geometries that favor the copper(II) state, therefore causing irreversibility. However, for 3, dissociation of the linear [CuCl2]− unit exposes the mixedvalence dimeric unit to oxidation with less influence from Cl−, thus leading to the slightly higher oxidation potential and its reversibility between the CuIICuI and CuIICuII states at high scan rates. No oxidation behavior was observed for the Cu2 center in 3 because of the fact that a copper center with a linear geometry is usually resistant toward oxidation.33 Synthesis and Characterization of 4a−4c. On the basis of the coordination structures and electrochemical behavior of complexes 2 and 3, we envisioned that the ligand H2pcp would also be an excellent platform to support fully oxidized copper(II) clusters. In fact, in our previous work based on the ligand H2pcp, we have demonstrated the interesting coordination properties of the amidate group through amide proton migration or metalation in the absence of a base.9b In the resulting mononuclear and dinuclear copper(II) complexes, strong coordinating anions, such as Cl− and OAc−, were required as terminal ligands to support the copper(II) center(s). In an effort to minimize the influence from anions, we employed copper(II) salts with weak or noncoordinating anions such as ClO4−, NO3−, and BF4−. Scheme 3 illustrates the synthesis and structure of the dinuclear copper(II)

Figure 3. NTO comparison between the 583 nm transition of 2 (top) and the 507 nm transition of 3 (bottom). The percent contribution to the overall transition is given for each pair of holes (left, blue and purple) and particles (right, orange and red). delocalization across the two copper centers, with virtually no unpaired electron density residing on the formally cuprous Cu2 in 3. Therefore, the calculated electronic structures of 2 and 3 provide support for the

Figure 4. EPR spectra (black, experimental; red, simulation) of complexes 2 (left) and 3 (right) in CH2Cl2 at room temperature. EPR parameters for 2: g = 2.178; A = 50 G. EPR parameters for 3: g = 2.167; A = 47 G. F

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

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Figure 5. EPR spectra (black, experimental; red, simulation) of complexes 2 (left) and 3 (right) in frozen 1:1 CH2Cl2/toluene (v/v) at 77 K. EPR parameters for 2: g1 = 2.112; g2 = 2.251; g3 = 2.266; A1 = 15 G; A2 = 55 G; A3 = 133 G. EPR parameters for 3: g1 = 2.132; g2 = 2.165; g3 = 2.274; A1 = 5 G; A2 = 45 G; A3 = 103 G.

Figure 6. UNO-SOMO images of complexes 2 (A) and 3 (B).

Scheme 3. Synthesis of 4a

Figure 7. X-ray structures of 4a. Thermal ellipsoids are shown at the 50% level. Cu1···Cu2 = 2.9896(14) Å. fashion, leading to different coordination geometries of the copper sites. Cu1 is chelated by one amidate nitrogen and one pyridyl nitrogen from the ligand highlighted in blue, and the coordination environment is fulfilled by two pyridyl nitrogens from two other ligands (red and green) and one weakly coordinated ClO4− anion. The τ5 value of Cu1 is 0.18, suggesting a distorted square-pyramidal geometry with an elongated axis [Cu1−O1E = 2.623(4) Å; dashed line in Figure 7], while for the Cu2 center, two deprotonated ligands (red and green) chelate it through amidate and pyridyl nitrogen donors, and the fifth position is occupied by the pyridyl nitrogen from the third ligand (blue). The τ5 value of 0.56 indicates that the geometry of Cu2 is a structural intermediate between the standard squarepyramidal and trigonal-bipyramidal geometries. The Cu1···Cu2 distance of 4a (2.990 Å) is much larger than those in the CuICuI (∼2.578 Å in 1a) and Cu1.5Cu1.5 (∼2.358 Å in 2) dimers, suggesting no direct interactions between the copper centers. Complexes 4b and 4c have dimeric core structures nearly identical with that of 4a but with different axial ligands on the Cu1 center (Figure S15). The NO3− anion coordinates to the Cu1 atom in 4b

complex 4a. Et3N was applied first to deprotonate the amide group in the ligand H2pcp, followed by the addition of 2/3 equiv of Cu(ClO4)2· 6H2O in CH3CN. Similar reactions with different stoichiometry ratios (1:1 ratio of ligand to metal) or in different solvents (MeOH and acetone) afforded the same compound, suggesting the excellent stability of 4a. X-ray crystallography revealed a novel dimeric structure supported by three deprotonated ligands. Because of the slight disorder of the copper(II) centers in 4a, Cu(NO3)2·3H2O and Cu(BF4)2·6H2O were applied in the synthesis to obtain better structures. As a result, structural analogues of 4a−4c were isolated and purified. All three complexes were very stable in both solid and solution states, and no changes were observed when stored in air for months. Figure 7 depicts the X-ray crystal structure of 4a, and Table 3 shows the selected bond lengths and angles. Two copper(II) centers are supported by three bridging amidate ligands in an unsymmetrical G

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

Cyclic voltammetry measurement of 4a in CH2Cl2 showed a reversible wave at E1/2 = −0.13 V in both scan directions, which is attributed to a CuICuII/CuIICuII couple (Figures S20, left, and S21). Consideration of the copper center geometries in 4a suggests that Cu1 is the oxidation/reduction site because its square-pyramidal geometry with a weakly coordinated ClO4− anion requires less structural rearrangement during reduction and oxidation. Further scans showed the independence of ΔE (200 mV) toward scan rates (Figure S20, right), indicating the electrochemical reversibility of this wave. One irreversible peak was also observed at −0.83 V, which might be the reduction of the Cu2 center with the distorted trigonal-bipyramidal geometry. The structure rearrangement required for the accommodation of copper(I) during the reduction of copper(II) may cause the irreversibility. In an effort to directly compare the oxidation states of the copper centers in complexes 1a, 2, 3, and 4a, XPS measurements were preformed. The binding energy of Cu 2p electrons obtained from this technique can reveal important information regarding copper oxidation states.34 As shown in Table 4 and Figure 8, a solid sample

Table 3. Selected Bond Lengths (Å) and Angles (deg) of Complex 4a Cu1−N1A Cu1−N1B Cu1−N2C Cu1−N3C Cu1−O1E

2.024(4) 1.985(4) 1.964(4) 2.023(4) 2.623(4)

Cu2−N2A Cu2−N1C Cu2−N2B Cu2−N3A Cu2−N3B

1.950(4) 2.059(4) 2.133(5) 2.120(4) 1.987(4)

Cu1···Cu2

2.9896(14)

N1A−Cu1−N1B N1A−Cu1−N2C N1A−Cu1−N3C N1B−Cu1−N2C N1B−Cu1−N3C N2C−Cu1−N3C N1C−Cu2−N2A N1C−Cu2−N2B N1C−Cu2−N3A N1C−Cu2−N3B N2A−Cu2−N2B N2A−Cu2−N3A N2A−Cu2−N3B N2B−Cu2−N3A N2B−Cu2−N3B N3A−Cu2−N3B

97.11(16) 95.19(17) 164.7(2) 154.0(2) 92.20(17) 81.19(17) 94.11(17) 130.69(18) 139.8(2) 92.08(17) 98.2(2) 80.90(17) 173.17(18) 89.39(17) 79.82(18) 92.50(18)

Table 4. Cu 2p3/2 and Cu 2p1/2 Binding Energies from XPS Measurements

through a long Cu1−O1D distance of 2.494(4) Å, while the water (H2O) molecule coordinates to the Cu1 atom in 4c with a Cu1−O1S bond of 2.333(5) Å due to the noncoordinating property of the BF4− anion. A comparison of the three structures revealed some interesting features: (a) In the solid state, the Cu1 center is stabilized by weak coordination from the fifth ligand, either an anion from the copper(II) salt or a H2O molecule from the solvent. (b) The increasing size of the weakly coordinated axial ligands H2O (28.33 Å3) < NO3− (51.62 Å3) < ClO4− (66.76 Å3)16 is reflected by the increasing Cu1−O bond distances Cu1−O1S [H2O, 2.333(5) Å] < Cu1−O1D [NO3−, 2.494(4) Å] < Cu1−O1E [ClO4−, 2.623(4) Å]. (c) No disorder was observed on copper(II) sites when noncoordinating the BF4− anion was used. The last feature provided detailed information about the structure with better accuracy. For instance, the τ5 values of Cu1 and Cu2 in 4c are 0.35 and 0.57, respectively, confirming the geometry assignments based on the X-ray data of 4a. The Cu1···Cu2 distances of 4a [2.9896(14) Å] and 4c [2.9778(11) Å] are also comparable. The selected bond lengths and angles of 4b and 4c are presented in Table S4. Electronic absorption spectra of 4a were collected in CH3CN and CH2Cl2 (Figure S16). Two d−d bands were observed at 640 and 756 nm in both solvents. In computational analysis of the UV−vis properties of 4a, we examined the role of the axial ClO4− ligand on Cu1 (Figure S17) with solvent effects. The calculated UV−vis spectra without ClO4− showed two peaks at 650 and 771 nm in both solvents (Figure S17, right), which are comparable to the experimental results. In contrast, the coordination of ClO4− on Cu1 significantly shifted the bands (Figure S17, left) relative to the experimental spectra. Therefore, we conclude that ClO4− dissociates from the Cu1 center when dissolved in weak or noncoordinative solvents, and the influence of these solvent molecules on the structure is negligible. The X-band EPR spectrum of 4a in a CH3CN solution was collected at room temperature (Figure S18). The four-line hyperfine pattern indicated that the unpaired electron interacts with only one copper(II) center. This observation was evidenced by the magnetic studies of the complex. Effective magnetic moment (μeff) measurement of 4a in CD2Cl2 at room temperature by using Evans’ method gave a value of 1.79 μB for each copper(II), suggesting localization of the single electron on each copper center. SQUID measurement from 2 to 300 K was also performed (Figure S19). The χm data fitting gave a J value of −29 cm−1, which means the two copper centers interact with each other with a weak antiferromagnetic coupling. The μeff value at room temperature obtained from SQUID (1.78 μB per copper) also matches well with the data obtained from Evan’s method (1.79 μB per copper), indicating conservation of the dimeric structure of 4a in a CH2Cl2 solution. These data match well with the results of a brokensymmetry CAM-B3LYP calculation on the dimeric structure of 4a, which predicts a J value of −30 cm−1.

binding energy (eV) Cu 2p3/2 1a 2 3 4a

931.9 932.8 932.6 936.3

± ± ± ±

0.06 0.02 0.03 0.46

Cu 2p1/2 951.7 952.5 952.4 956.5

± ± ± ±

0.06 0.02 0.03 0.29

Figure 8. XPS spectra of complexes 1a, 2, 3, and 4a recorded in the Cu 2p region.

of 1a exhibits a Cu 2p3/2 peak at 931.9 ± 0.06 eV that is comparable with that of CuICl (2p3/2 = 932.2 eV).35 In contrast, 4a shows a Cu 2p3/2 peak at 936.3 ± 0.46 eV and a strong satellite peak, which is a typical feature for copper(II) compounds.34−36 The most interesting feature is the Cu 2p3/2 peaks at 932.8 ± 0.02 and 932.6 ± 0.03 eV exhibited by 2 and 3, which locate between the Cu 2p3/2 peaks of 1a and 4. This observation clearly suggests intermediate oxidation states of 2 and 3 between I and II. This Cu 2p3/2 binding energy trend (1a < 2 ≈ 3 < 4a) was also observed from the Cu 2p1/2 peaks in the 951.7− 956.5 eV range. This is consistent with oxidation state assignment for the four complexes based on X-ray crystallography and spectroscopic evidence. H

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry By analysis of the calculated Cu 2p orbital energies of complexes 1a, 2, 3, and 4a, XPS of these compounds can also be rationalized. To increase the accuracy of these calculations, the integral grid was increased significantly and the Orca core-property basis set [CP(PPP)] was used for the copper centers. To aid in the interpretation of the results, a restricted open-wave function was calculated for the openshell molecules 2, 3, and 4a. The resulting Cu 2p orbital energies for each of these compounds are shown in Scheme 4. The trend observed

complexes are superposed to compare the conformations for the rest of the structures. Complex 4 is not shown because of the presence of three amidate ligands. Foremost of the observed features is the different Cu−Cu distances in complexes 1a (∼2.578 Å), 2 (∼2.436 Å), 3 (∼2.358 Å), and 4a (∼2.990 Å) supported by the ligands with the same μ2-(κNpy,Namidate;κNpy) coordination mode. Additionally, as shown in Figure 9, the top ligands of complexes 1a, 2, and 3 have comparable conformations, but the bottom ligands showed very different spatial orientations. As a result, the average Cu−N bond length (∼2.020 Å) of 1a is longer than the ∼1.973 Å length observed in 2 and 3. Another distinctive conformation feature is the torsion angle of the two pyridyl rings within an amidate ligand. It varies from 27.98° in 1 to 26.48° in 2 to 12.37° in 3. For 4, the three ligands adapt even larger dipyridyl torsion angles (39.23−43.91°) to fulfill the distorted square-pyramidal and trigonal-bipyramidal geometries of the two copper(II) centers. Therefore, our analysis based on these features revealed that the flexible coordination behavior exhibited by the ligand 2pcp− should be attributed to the rotation of the pyridyl rings and bending of the amidate group.

Scheme 4. Calculated Cu 2p Orbital Energies for Compounds 1a, 2, 3, and 4a

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CONCLUSION With X-ray crystallography and spectroscopic characterization, we have demonstrated the ability of the dipyridylamide ligand H2pcp to support copper complexes with different oxidation states. The CuICuI core in complex 1a is stabilized by two deprotonated H2pcp ligands with a μ2-(κNpy,Namidate;κNpy) coordination mode. Electrochemical studies showed a reversible oxidation/reduction wave assigned as a CuICuI/CuIICuI couple, which inspired us to establish the synthetic procedures for valence intermediates. The resulting dinuclear and trinuclear copper complexes 2 and 3 with mixed-valence core structures were characterized by X-ray crystallography and spectroscopic methods. EPR of both complexes at room temperature and 77 K showed characteristic seven-line hyperfine signals, suggesting delocalized Cu1.5Cu1.5 mixed-valence states of the complexes. These results prompted us to further isolate complex 4 featuring a dimeric CuIICuII core structure supported by three amidate ligands. In contrast to the close Cu−Cu interactions in 1a, 2, and 3, spectroscopic and magnetic studies of 4 have demonstrated a weak antiferromagnetic interaction between the two copper(II) centers because of the large Cu···Cu distance. XPS measurements of all four complexes have clearly shown the increasing trend of Cu 2p electron binding energy with the increasing oxidation states (1a < 2 ≈ 3 < 4a). These results served as clear evidence about the intermediate valence of 2 and 3 between copper(I) and copper(II). The unrestricted NTO analysis revealed the similar electronic structures of mixed-valence complexes 2 and 3, which are responsible for the similar spectroscopic properties of the two compounds. In addition, the calculated Cu 2p orbital energies rationalized the XPS features observed for complexes 1a, 2, 3, and 4a.

here is consistent with the experimental XPS data, in both the binding energies and the widths of the peaks. In 3, there are two distinct copper environments and binding energies corresponding to the mixed-valence Cu1.5Cu1.5 and cuprous [CuCl2]− groups. These two Cu 2p energy regimes in 3 match well with the individual environments in 1a and 2. The expected double peak is not observed in XPS of 3, which might be due to convolution of the copper(I) signal with the mixed-valence signal because the latter is rather broad. Another XPS feature that is explained by these calculations is the broadness of the peak from 4a. Unlike the other compounds, 4a is a polar and less symmetric molecule, both of which give rise to greater directional anisotropy for the Cu 2p orbitals. This results in a predicted breadth of energy of 0.21 eV, over 20 times greater than that for the CuI species (0.01 eV). Given the X-ray crystal structure data of all four complexes, it would certainly be interesting to inspect the coordinative flexibility exhibited by the amidate ligand 2pcp− that accommodates the copper centers with different oxidation states. Figure 9 shows the graphical overlay of the dimeric copper core structures of 1a (green), 2 (blue), and 3 (red), in which the pyridyl rings with the N1 donors from the three

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.6b02006. Spectroscopic data and cyclic voltammetry (PDF) X-ray crystallographic data in CIF format (CIF)

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Figure 9. Dimeric copper core structure overlay of complexes 1a (green), 2 (blue), and 3 (red) based on the pyridyl ring with the N1 donor. The Cl1 donor of 2 and the [CuCl2]− unit in 3 are omitted for clarity. Complex 4 is not shown because of the presence of three amidate ligands.

AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Fax: +1-(501)-450-3623. I

DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

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Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the start-up fund provided by the College of Natural Sciences and Mathematics and University Research Council from the University of Central Arkansas, Arkansas Science & Technology Authority, from the State of Arkansas. We thanks INBRE of Arkansas for the purchase of the EPR instrument. We also thank Magnus Pauly for preparing samples of complex 1a.

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DOI: 10.1021/acs.inorgchem.6b02006 Inorg. Chem. XXXX, XXX, XXX−XXX