Solid-State Electrochemistry of Copper (I) Coordination Polymers

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Solid-State Electrochemistry of Copper(I) Coordination Polymers Containing Tetrafluoroborate Anions Hitoshi Kumagai,*,† Satoshi Kawata,‡ and Hideyuki Nakano† †

Toyota Central R&D Laboratories, Inc., 41-1 Yokomichi, Nagakute, Aichi 480-1192, Japan Department of Chemistry, Fukuoka University, Jonan-ku, Fukuoka 810-0180, Japan

‡

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S Supporting Information *

ABSTRACT: Host−guest materials based on coordination polymers (CPs) are currently emerging as potential candidates for battery applications. In this context, we describe the preparation of threedimensional network structures containing BF4 anions and water molecules in the one-dimensional (1D) channels via hydrothermal reactions between Cu(BF4)2 and 4,4′-bipyridine or 1,2-di-4-pyridylethylene. A systematic characterization of the obtained CPs using single-crystal X-ray diffraction, X-ray absorption fine structure, and an electrochemical test was performed. The results showed that the BF4 anions were electrochemically reduced to BF3 in the cavities of the CPs, with concomitant elimination of a leaving fluoride at room temperature. Using this electrochemical property, a prototype battery, in which the CPs act as the anode and graphite as the cathode, was demonstrated. The cell exhibited a practical discharge potential of ∼1.5 V. This constitutes the first demonstration of CPs showing electrochemical B−F bond activation in the 1D channels and rocking-chair-type fluoride insertion and extraction by changes in the electric potential.

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INTRODUCTION Network materials, such as coordination polymers (CPs), also known as metal−organic frameworks (MOFs), are attracting much interest for their applications in the design of molecular magnets because of the strength and rigidity of the extended lattices allowing for gas sorption, intercalation, and connectivity between magnetic ions.1−3 Furthermore, exploitation of the chemical properties, such as selective sorption− desorption and catalysis, which stem from the porous nature of CPs, is receiving increased attention. In this context, the use of CPs as battery materials has recently been added to the list of their potential applications.4−8 CPs usually employ a central metal ion and a multitopic organic linker or a coordination complex having ambidentate ligands, such as cyanide and oxalate. The choice of the metal ions and bridging ligands depends on the desired properties. Thus, we previously reported the use of polycarboxylates or dihydroxybenzoquinone derivatives as bridging ligands for the preparation of host−guest materials, which was accompanied by transformation of the magnetic ground state or a color change of the materials.9−14 Following our studies on host− guest materials, we describe our investigations on the reversible anion extraction and insertion in CPs. These materials could find use as the anode in batteries, provided that the anions are reversibly extracted from and inserted into the CPs. This study focuses on the synthesis and electrochemical properties of CPs synthesized using neutral polypyridyl ligands, i.e., 4,4′bipyridine (4,4′-bpy) or 1,2-di-4-pyridylethylene (dpe), as © XXXX American Chemical Society

bridging ligands for the formation of frameworks that serve as host materials for anions (Scheme 1).15−18 Our concept is Scheme 1. Molecular Structures of the Ligands (a) bpy and (b) dpe

based on the formation of a cationic framework via the reaction between metal cations and neutral bridging ligands that defines the distances between metal centers, in which anion entrapment occurs to balance the residual positive charge resulting from formation of the metal complex.

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EXPERIMENTAL SECTION

Caution! Copper tetraf luoroborate causes severe skin and eye damage, so precautions with suitable care and protection for handling have been followed. Copper tetrafluoroborate hexahydrate, 1,2-di-4-pyridylethylene (dpe),and 4,4′-bipyridine (4,4′-bpy) were purchased from Wako and Aldrich, respectively, and used without further purification. Preparation of [Cu(4,4′-bpy)2(BF4)(0.5H2O)] (1). An aqueous solution (5 mL) of copper tetrafluoroborate hexahydrate (0.47 g, 0.20 mmol) and 4,4′-bpy (0.62 g, 0.40 mmol) was mixed in distilled water Received: October 4, 2018

A

DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry (30 mL). The mixture was placed in the Teflon liner of an autoclave, sealed, and heated to 170 °C for 3 days. The bomb was allowed to cool to room temperature. Orange crystals were obtained. The crystals were washed with water and acetone and dried in air (yield 50%). One of these crystals was used for X-ray crystallography. Calcd for 1 (C20H17BCuF4N4O0.5): C, 50.92; H, 3.63; N, 11.88. Found: C, 51.08; H, 3.56; N, 11.95. Selected IR data (ν/cm−1): 3450(br), 2970(br), 1616(m), 1478(m), 1418(s), 1389(s), 1322(s), 1298(s), 1145(w), 1130(w), 1083(w), 1047(m), 1038(m), 804(s), 736(m). Preparation of [Cu(dpe)2(BF4)(1.5H2O)] (2). An aqueous solution (5 mL) of copper tetrafluoroborate hexahydrate (0.47 g, 0.20 mmol) and dpe (0.72 g, 0.40 mmol) was mixed in distilled water (30 mL). The mixture was placed in the Teflon liner of an autoclave, sealed, and heated to 170 °C for 3 days. The bomb was allowed to cool to room temperature. Orange crystals were obtained. The crystals were washed with water and acetone and dried in air (yield 50%). One of these crystals was used for X-ray crystallography. Calcd for 2 (C24H22BCuF4N4O): C, 53.20; H, 4.28; N, 10.34. Found: C, 53.54; H, 4.27; N, 10.56. Selected IR data (ν/cm−1): 3450(br), 2970(br), 1616(m), 1478(m), 1418(s), 1389(s), 1322(s), 1298(s), 1145(w), 1130(w), 1083(w), 1047(m), 1038(m), 804(s), 736(m). General Characterization. Thermogravimetric analysis (TGA; 30−500 °C) was performed on a Shimazu TGA-50 TGA−differential thermal analysis (DTA) system. IR spectra were recorded on a Thermo Nicolet Avatar 360 Fourier transform infrared spectrophotometer. The absorption spectra were measured using a JASCO V-560 spectrophotometer. The nitrogen adsorption isotherm was measured using a Nova 3000e (Quantachrome) at −196 °C after evacuation at 180 °C. The pore diameter was calculated by using the Barrett− Joyner−Halenda method. The specific surface area was calculated by using the Brunauer−Emmett−Teller plot for the adsorption branch. X-ray diffraction (XRD) measurement was carried out with a Rigaku Rint-TTR X-ray diffractometer using Cu Kα radiation. The electrochemical properties of the cells were examined using a Hokuto Denko HJ10/10 mSM8A charge−discharge unit. Cyclic voltammetry (CV) measurements were performed on a Hokuto Denko HZ-5000 polarization system. Electrochemical impedance spectroscopy was measured on a Solartron Analytical 1400 cell test system. X-ray Crystallography and Structure Solution. A suitable single crystal was chosen and glued on the tip of glass fibers. Diffraction data for the complexes were collected on a Rigaku R-AXIS RAPID area detector employing the ω-scan mode at room temperature. The diffractometer was equipped with graphitemonochromated Mo Kα (0.7107 Å) radiation. The data were corrected for Lorentz and polarization effects. The structure was solved by direct methods and expanded using Fourier techniques. The non-hydrogen atoms were refined anisotropically. Hydrogen atoms were refined using a riding model. The final cycle of a full-matrix leastsquares refinement was based on the number of observed reflections and n variable parameters. They converged (a large parameter shift was σ times its estimated standard deviation) with agreement factors of R1 = ∑||Fo| − |Fc||/∑|Fo| and wR2 = [∑(w(Fo2 − Fc2)2)/ ∑w(Fo2)2]1/2. No extinction corrections were applied. Details of the crystallographic data are given in Table 1. The crystal data have been deposited as CCDC 1857066 (1) and 1857065 (2).

Table 1. Detail of Crystallographic Data and Refinements for Compounds 1 and 2 empirical formula fw cryst syst space group a (Å) b (Å) c (Å) α (deg) β (deg) γ (deg) V (Å3) Dc (g cm−3) R1 (wR2) (all data) R1 [I > 2.0σ(I)] GOF

1

2

C40H34B2Cu2F8N8O 943.46 monoclinic P21/c (No. 14) 7.1790(2) 32.2959(7) 17.3595(4)

C24H13BCuF4N4O1.5 541.82 triclinic P1̅ (No. 2) 8.025(2) 12.556(3) 12.669(3) 81.121(6) 75.053(6) 77.378(7) 1197.0(5) 1.503 0.0649(0.1556) 0.0555 1.058

101.0070(9) 3950.8(2) 1.586 0.0667 (0.1908) 0.0631 1.328

demonstrating that the hydrothermal conditions are essential for the synthesis of 1 and 2. These results are similar to those of Cu(4,4′-bpy)1.5(NO3)(H2O)1.25 and Cu(dpe)2(NO3).15,18 The X-ray crystal structure analyses of compounds 1 and 2 revealed the extended cationic framework constructed from the tetrahedral copper(I) ions and neutral bridging ligands. Because of the similarity between the structures of 1 and 2, we will restrict our discussion to 1, and the pertinent points regarding 2 will be mentioned when appropriate. Figure 1 shows the framework structures of 1 and 2, which contain metal ions, bridging ligands, water molecules, and BF4 anions. Details of the crystallographic data are listed in Table 1, and selected bond distances and angles around the metal centers along with the atomic numbering scheme are shown in Tables S1 and S2 and Figures S1 and S2, respectively. The structure is composed of a 4-fold interpenetrated three-dimensional (3D) network of copper(I) ions bearing four bridging ligands in tetrahedral coordination. The key feature of the structure is the presence of sizable channels running along the a crystallographic direction of the 3D network. The distance between copper(I) ions bridged by the bpy ligand is 10.18 Å, and the nearest distance between copper(I) ions is 7.18 Å. The structure of 1 comprises two types of channels (A and B). Type A is filled with BF4 anions and water molecules, whereas type B is only filled with BF4 anions. The distorted tetrahedral geometry of Cu(1) is characterized by four Cu−N distances in the range of 1.996(3)−2.060(5) Å and N−Cu−N angles ranging from 101.93(17)° to 118.25(18)°. Compound 1 exhibits two types of bpy ligands, planar and distorted from the planar structure, in which the torsion angles between two pyridyl groups range from 10.1° to 41.2°. Despite having similar 3D structures, there are some structural differences between 1 and 2. The structure of 2 is related to the previously reported isomorphous structures of Cu(dpe)2(BF4)·0.5CH2Cl2 and Cu(dpe)2(BF4)·MeCN, which comprise similar 5-fold interpenetrated network structures.16,17 The degree of interpenetration has been increased from four independent lattices for 1 to five independent arrays for 2. This comes from the ligand length, π−π, and steric interactions.16 The structure of 2 contains only one type of channel, which contains water molecules and BF4 anions. The CuI···CuI distance bridged by the dpe ligand is 13.46 Å, and the nearest CuI···CuI distance is 8.30 Å. The longer CuI···CuI distances

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RESULTS AND DISCUSSION Compounds 1 and 2 were obtained as orange or brown crystals, respectively, from conventional hydrothermal reactions between copper(II) ions and the corresponding ligand, 4,4′-bpy or dpe, in a 1:2 molar ratio. The reduction reactions and the crystal structures of the compounds were confirmed via an X-ray structure analysis performed on single crystals isolated from the reaction mixtures. Similar reduction reactions under hydrothermal conditions were previously reported to afford Cu(4,4′-bpy)1.5(NO3)(H2O)1.25 and Cu(dpe)2(NO3).15,18 Attempts to synthesize 1 and 2 under ambient conditions or refluxing in water failed, thus B

DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX

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compared to those of 1 are due to the −CC− bond between pyridyl groups. Moreover, all of the dpe ligands of 2 exhibit a planar structure, which contrasts with the presence of planar and nonplanar ligands in 1. To assess the thermal properties of each compound, TGA was performed in air. Figure S3 depicts the results of the TGA of 1 and 2, which are characterized by three weight losses up to 500 °C. The weight was found to gradually decrease with increasing temperature in the first step from room temperature to 200 °C, followed by an abrupt decrease. The first weight loss amounts to 1.1% and 3.4% for compounds 1 and 2, respectively, from room temperature to 170 °C, and the second weight loss corresponds to 36% at 320 °C for both compounds. The first weight loss step corresponds to the loss of water molecules in the one-dimensional (1D) channel to give the dehydrated phases of 1 and 2. The second weight loss is due to thermal decomposition and combustion of the organic ligand to form the oxide phase. Compounds 1 and 2 were heated to 170 °C in vacuo to yield the dehydrated samples, and the corresponding IR spectra were measured. Figure S4 shows the IR spectra of the as-synthesized and dehydrated samples. The IR spectra of the complexes in the 400−4000 cm−1 range, which are very similar to one another, are dominated by the vibrational modes of the ligands, BF4 anions, and water molecules. The stretching modes of the water molecules are broad and locate around 3600 cm−1. Formation of the dehydrated samples of 1 and 2 was evidenced by the disappearance of the absorption bands of the water molecules from the IR spectrum, whereas those of the bridging ligands and BF4 anions remained visible. These results are in good agreement with those of the TGA measurements. The electrochemical properties of 1 and 2 were then evaluated. The corresponding electrochemical cells were assembled using 1 and 2 dehydrated at 180 °C, with the negative electrode containing 70% by weight of 1 or 2, 25% of

Figure 1. (a) Projection of the 3D network structure of 1, viewed along the a crystallographic axis. (b) Projection of the 3D network structure of 2, viewed along the a crystallographic axis. Color code: C, white; B, orange; F, violet; N, green; O, red; Cu, blue. Hydrogen atoms are omitted for clarity.

Figure 2. Potential−capacity profiles for 1 (a) and 2 (b). Insets: Corresponding capacity−retention curves. CV curves for 1 (c) and 2 (d). C

DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry carbon black as a conductive additive, and 5% of poly(tetrafluoroethylene) (PTFE) as a binder. In contrast, 95% by weight of graphite and 5% of PTFE as a binder were mixed and used as a positive electrode, and a 1:1 (v/v) ethylene carbonate/diethyl carbonate (EC/DEC) solution of LiBF4 (1 M) was used as the electrolyte. Figure 2a shows voltage versus capacity plots for the cells of 1 and 2 under cell voltages from 0 to 3.5 V at room temperature. The current density was 100 mA g−1. The charge cycle was induced by supplying electrons to the CPs from an external bias, which corresponds to the extraction of anions, whereas the discharge cycle was induced by the release of electrons as a result of an external bias, i.e., the insertion of anions. Initially, the potential increased gradually until it reached a plateau near 3.5 V. The first discharge capacities for 1 and 2 were found to be 36.5 and 25.5 mAh g−1, respectively. Because the expected specific capacities associated with one-electron redox exchange per unit formula of the CPs are 58 mAh g−1 for 1 and 53 mAh g−1 for 2, the obtained values correspond to 63% and 48% of the theoretical capacities for 1 and 2, respectively. Interestingly, when the same measurements were performed using a similar cell in the absence of 1 or 2, the charge−discharge capacities were not observed. These results indicate that the anions of 1 and 2, and not those present in the electrolyte solutions, are responsible for the observed electrochemical activities. Despite having an irreversible capacity of 57% for 1 and 56% for 2 between the first charge and discharge cycles, the reversible reduction− oxidation process in the CPs is most likely primarily due to the irreversible reaction of the graphite used as the cathode. After 100 cycles, the capacity retention slowly decays to values of about 15% for 1 and 40% for 2, respectively (Figure 2a,b, insets). Structural/textural/morphological material changes upon cycling or material solubility issues pertaining to the nonoptimized electrolyte are often invoked as the reasons for such fading. Therefore, further optimization of the electrode and formation of the cells may lead to an increase of the capacities. The electrochemical properties of the CPs were investigated by CV using the same cells mentioned above. Parts c and d of Figure 2 show the initial first and second CV curves of the 1 and 2 electrodes at a scan rate of 1 mV s−1, respectively. In the first cathodic cycle, the peak at around 2.8 V can be associated with oxidation of the BF4 anion in the cavity of the CPs to form BF3 (i.e., BF4 + xe− → BF3 + F−). In the reverse scans, one broad anodic peak appears at 2.0 V, corresponding to the reduction of BF3 to the BF4 anion. It is worth noting that the fluorine-ion redox mechanisms of the CPs after the first cycle primarily come from the reversible redox reactions of the BF4 anion at 2.8/2.0 V. Moreover, the enveloped redox peak area shows a slight decrease upon the initial cycling after the second cycle, indicating a decrease in the electrochemical stability and reversibility, which is consistent with the cycling performance. To investigate the redox mechanism that occurred during the electrochemical reactions, the structures of both electrode were investigated using XRD to gain insight into the structural changes of the materials. Figure 3a shows the first-cycle charge−discharge curves of 1 and 2, in which we monitored the starting (black), 3.5 V charge (red), and 0.05 V discharge (blue) samples. Figures 3b and S5 show the XRD peaks of the negative electrode (CPs 1 and 2) after the electrochemical reactions were slightly shifted to lower 2θ values compared to those of the starting samples, indicating lattice expansion. Although the diffraction patterns remained virtually unaltered,

Figure 3. (a) Charge−discharge curves of cells of 1 (green) and 2 (violet). (b) XRD pattern of 2 showing the reversibility of the discharge−charge process. (c) B K-edge XANES spectra of 2 showing the reversibility of the discharge−charge process. Color code: starting, black; discharge, red; charge, blue.

new diffraction peaks were observed. These results are indicative of the stability of the framework during the electrochemical process. The network structure of Cu(dpe)2 exhibited remarkable thermal stability and anion-exchange properties without producing any damage.18 In the case of the positive electrodes (Figure S6), a peak at 2θ = 26.50° was observed at the initial stage of the reaction. After reduction, this peak shifted to lower 2θ values, 26.34° for 1 and 26.29° for 2, which is indicative of the insertion of the anions into the graphite layers. In contrast, when the reduced samples were oxidized, the peaks shifted to higher 2θ values of 26.45° and 26.51° for 1 and 2, respectively, suggesting that extraction of the anions from the graphite layers occurred. These results are indicative of the reversibility of the electrochemical anion extraction and insertion into the graphite layers. The expanding space was, however, smaller than the previously D

DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry reported BF4− intercalated graphite.19 To gain further insight into the insertion and extraction of anions induced by the change in the electric potential, we performed ex situ B and N K-edge X-ray absorption near-edge structure (XANES) measurements on the samples before electrochemical measurements: the reduced samples in which anions were extracted at 0 V and the oxidized samples after reduction in which anions were inserted at 3.5 V. The results of the B K-edge XANES measurements are shown in Figures 3c and S7, in which a peak can be observed at ∼200 eV before electrochemical measurements. This peak corresponds to the σ* transition of the tetrahedral boron atom. The XANES spectrum of the oxidized sample after electrochemical extraction of the anions exhibits two peaks at 194 and 200 eV. The peak of lower energy is attributable to the π* transition of the boron atom, whereas the other peak coincides with that of the virgin sample. The presence of the new peak is indicative of the structural change of the BF4 anions to give planar BF3 molecules during electrochemical oxidation. Notably, the presence in the XANES spectrum of the peak attributed to BF4 anions after electrochemical reduction suggests that some of these anions remained in the 1D cavity, which is in accordance with the observed charge−discharge values corresponding to 42% and 67% of the theoretical values. The XANES spectrum of the sample in which anions are inserted electrochemically after electrochemical extraction of the BF4 anions shows the peak at 200 eV, whereas that at 194 eV disappears. This result is indicative of the reaction of the planar BF3 molecule with a fluorine anion, affording the tetrahedral BF4 anion.20,21 Some batteries based on fluoride conductors attaching or detaching into or from electrode materials have been previously reported.22 However, these systems require working temperatures above 150 °C because they use solid electrodes. In the absence of 1 or 2, electrochemical activities are not observed, which suggests that electrochemical activation of the B−F bonds of the CPs in the cavity is required. It has been reported that the CPs could adsorb and stabilize in their frameworks toxic or explosive molecules such as BF3, CO, and acetylene. The BF3 molecule, which is a strong Lewis acid, would then be trapped in the cavity, and its reaction with other molecules would be prevented. In contrast, the van der Waals radius of the BF4 anion had been previously determined from MM2 and ab initio molecular orbital calculations as 0.227 nm.23,24 Considering this value, the size of the cavity of 1 and 2 seems too small for effective electrochemcal extraction of the BF4 anions, which provides further support to the key role of the B−F bond activation (Figure 4). CPs have been applied to

lithium-ion battery or capacitor materials; however, reports on their application to asymmetric anion rocking-chair-type battery materials are still sparse.4−8 This is the first example of secondary battery materials using electrochemical B−F bond activation. The N K-edge XANES spectra of the reduced and oxidized samples after electrochemical reduction are shown in Figure S8, wherein the N K-edge XANES spectrum of 1 exhibits similar shapes after reduction and oxidation. In contrast, the N K-edge XANES spectrum of the reduced sample of 2 shows a shift from the initial value of 399.7 to 399.9 eV, which returns to the original value of 399.7 eV after oxidation. The energy level of the transition of the XANES spectrum of 2 increases when the electron density on the nitrogen atom decreases. This difference between the spectra of 1 and 2 suggests the presence of interactions such as π−π interaction between the bridging ligands and BF3 molecules. Thus, the planar BF3 molecules would interact with the dpe ligand in 2, decreasing the electron density of the ligand due to the strong Lewis acidity of the BF3 molecule, whereas the distorted bpy ligands of 1 would prevent the formation of π−π interactions with BF3. Therefore, the higher value found for the specific capacity of 2 is most likely due to the enhanced stability of BF3 molecules in the lager cavity of 2. Finally, high C-rate tests were performed to evaluate their fast charge−discharge capabilities in practical applications, which is critical for EV. As can be extracted from Figure 5a,b, the 2 anode delivers much improved rate capability upon discharge compared with the 1 anode. Figure 5c clearly demonstrates the promising rate capability of the 2 anode, which is superior to that of the 1 anode in the tested range of C-rates. Meanwhile, the Nyquist plots of both anodes exhibit similar tendencies (Figure 5d,e), showing a depressed semicircle in the high-to-medium frequency range and a sloped line in the low-frequency range, which represents charge-transfer kinetics and mass diffusion processes, respectively. Notably, the charge-transfer resistance (Rct) of the 2 electrode (12.8 Ω) is much smaller than that of the 1 electrode (24.7 Ω) at room temperature, with this difference being more pronounced with decreasing temperature. Graphite as the cathode active material has low fluorine diffusion coefficients and low overall capacity; therefore, the differences found in the rate properties and Rct could be considered to result from the CP cavity size. Considering the cycling performance, we can conclude that CP cavities of larger size are suitable for fast charging; on the other hand, the smaller cavities have longer lifetime, depending on the stability of the frame.

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CONCLUSION In summary, we have demonstrated the preparation of anion secondary batteries using copper(I)-based 3D CPs as cathode materials. Elucidation of the redox mechanism revealed the occurrence of electrochemical B−F bond activation and the reversible reaction BF4− ⇄ BF3 + F− in the cavity of the CPs at room temperature. During the cycling process, the frameworks of the materials remained virtually unaltered, which suggests the potential application of 1 and 2 as anode materials for anion batteries. By designing suitable molecules (e.g., different anion species or framework sizes), we believe that the realization of anion secondary batteries beyond lithium-ion batteries will become possible. Our findings offer insights into the novel mechanism of secondary batteries and prospects for the development of unique electrostorage devices. Further

Figure 4. (a) Structure of 2 with the van der Waals radius. Color code: C, gray; Cu, blue; F, violet; N, green. (b) Schematic representation of the electrochemical B−F bond activation. E

DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX

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Figure 5. Fast-charging rate capacities of 1 (a) and 2 (b). (c) Variations in the normalized charge capacities depending on the C-rates at 25 °C. Nyquist plots of the 1 (d) and 2 (e) anodes at different temperatures.

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ACKNOWLEDGMENTS We thank Dr. T. Nonaka for Cu XANES spectroscopy measurements. We thank Dr. K. Yamanaka and Prof. T. Ohta for B and N K-edge XANES spectroscopy measurements and helpful discussions. We are grateful to M. Nakai for assistance with material synthesis.

studies to develop a rocking-chair-type anion battery with enhanced performance are currently in progress.

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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.8b02768. Results and additional data of XRD patterns, IR spectra of as-synthesized and dehydrated compounds, TGA, and B and N K-edge XANES spectra (PDF)

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Accession Codes

CCDC 1857065−1857066 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing [email protected], or by contacting The Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

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REFERENCES

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Hitoshi Kumagai: 0000-0001-7996-7836 Hideyuki Nakano: 0000-0002-2866-3282 Author Contributions

These authors contributed equally. Notes

The authors declare no competing financial interest. F

DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.inorgchem.8b02768 Inorg. Chem. XXXX, XXX, XXX−XXX