Article pubs.acs.org/IC
A Coordination Chemistry Approach for Lithium-Ion Batteries: The Coexistence of Metal and Ligand Redox Activities in a OneDimensional Metal−Organic Material Gaihua Li,† Hao Yang,† Fengcai Li,† Fangyi Cheng,† Wei Shi,*,†,‡ Jun Chen,† and Peng Cheng*,† †
Department of Chemistry, Key Laboratory of Advanced Energy Materials Chemistry (MOE), Nankai University, Tianjin 300071, China ‡ Collaborative Innovation Center of Chemical Science and Engineering (Tianjin), Nankai University, Tianjin 300071, China S Supporting Information *
ABSTRACT: We demonstrate herein the use of a one-dimensional metal−organic material as a new type of electrode material for lithiumion batteries (LIBs) in place of the classic porous three-dimensional materials, which are subject to the size of the channel for lithium-ion diffusion and blocking of the windows of the framework by organic solvents during the charging and discharging processes. Introducing a one-dimensional coordination compound can keep organic active substances insoluble in the electrolyte during the charging and discharging processes, providing a facile and general new system for further studies. The results show that both the aromatic ligand and the metal center can participate in lithium storage simultaneously, illustrating a new energy storage mechanism that has been wellcharacterized by X-ray photoelectron spectroscopy, electron paramagnetic resonance spectroscopy, and cyclic voltammetry. In addition, the fact that the one-dimensional chains are linked by weak hydrogen bonds rather than strong π−π stacking interactions or covalent bonds is beneficial for the release of capacity entirely without the negative effect of burying the active sites.
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INTRODUCTION Electrical energy storage materials are promising substitutes for traditional fossil fuels because of their renewability, low environmental impact, and low cost.1−3 Lithium-ion batteries (LIBs) are among the most important energy storage materials considering their cost-effective, highly efficient, and environmentally friendly energy storage characteristics. However, poor rate capabilities, relatively low practical capacities, and Liplating issues still limit their applications.4−6 Moreover, high energy consumption for the synthesis of electrode materials prevents their practical applications despite their excellent electrochemical performance. Hence, the development of new types of electrode materials for LIBs remains a great challenge for both chemists and materials scientists. Coordination compounds are a family of compounds containing coordination bonds, which are generally formed between a metal center and a donor atom of an inorganic/ organic ligand.7 A variety of practical applications for coordination compounds, such as in catalytic chemistry,8,9 biochemistry,10−12 and porous materials,13 have been developed since Alfred Werner won the Nobel Prize in Chemistry in 1913. In principle, coordination compounds are promising candidates as electrode materials with the following advantages: (i) the redox-active metal center can accommodate multiple electrons;14,15 (ii) the structural diversity and flexibility of the © XXXX American Chemical Society
ligands can be controlled at the molecular level, allowing the ligands to be redox-active, to be processed efficiently, and to be prepared from renewable resources;16−21 (iii) enhanced electrochemical performance can be achieved by the combination of redox-active metal centers and organic ligands; and (iv) the synthesis of coordination compounds requires relatively low energy consumption. However, regardless of their solubility in a given electrolyte, the charge and discharge processes may cause the decomposition of coordination compounds used as electrodes, which limits the practical applications of these materials. To date, the coordination compounds studied in the context of LIBs are simply based on redox-active metal ions or functional organic groups.22−25 To solve this problem, one promising strategy is to link the metal center with a redoxactive bridging/chelating ligand to form a one-dimensional (1D) chain or a high-dimensional framework, benefiting both the electrochemical and solid-state stability and reversibility. As a continuation of previous research on coordination chemistry using pyridine-2,6-dicarboxylic acid and its derivatives,26−29 we selected [Co1.5L(H2O)4]n (H3L = 4-hydroxypyridine-2,6-dicarboxylic acid) as an advanced anode active material for LIBs. The results of the present study reveal that Received: February 26, 2016
A
DOI: 10.1021/acs.inorgchem.6b00450 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry both the aromatic chelating ligand and the metal centers participate in lithium storage, suggesting a new energy storage mechanism for coordination compounds. More importantly, this work represents the first example of a 1D coordination compound as an active energy storage substance instead of using traditional porous three-dimensional materials, which depend on the size of the channel for lithium-ion diffusion and obstruction of the pore window via organic solvents during the charging and discharging processes. In addition, the 1D chains are linked by weak hydrogen bonds rather than by strong π−π stacking interactions or covalent bonds, allowing the materials to approach their maximum potential capacity without the negative effect of burying active redox sites.
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EXPERIMENTAL SECTION
General Methods and Materials. All of the reagents were purchased commercially and used as received. The single-crystal X-ray diffraction measurement of [Co1.5L(H2O)4]n was recorded on an Oxford Supernova diffractometer using graphite-monochromatized Mo Kα radiation (λ = 0.71073 Å). The elemental analyses (C, H, and N) were performed on a PerkinElmer 2400-II CHNS/O analyzer. Thermogravimetric analysis (TGA) was carried out under an atmosphere of N2 on a Labsys NETZSCH TG 209 Setaram apparatus at a heating rate of 10 °C min−1 from 25 to 800 °C. Powder X-ray diffraction (PXRD) patterns were collected on a Rigaku Ultima IV instrument equipped with a Cu Kα radiation source (λ = 1.54056 Å); the samples were scanned at 10° min−1 in the 2θ range from 3 to 50°. Direct-current (dc) magnetic susceptibilities were measured on a Quantum Design VSM SQUID magnetometer. Electrochemical characterizations were conducted using a two-electrode coin-cell configuration (CR2032). The charge−discharge curves were collected on a LAND battery testing system at room temperature (298 K). Cyclic voltammetry was carried out on a PARSTAT 4000 electrochemical workstation. Battery performance tests were performed using a CR2032 coin-cell battery setup. The anode containing a 60 wt % loading of the coordination compound was prepared according to the following procedure: 60 mg of as-prepared [Co1.5L(H2O)4]n was ground into powder with a mortar and pestle, and 30 mg of Super P and 10 mg of polyvinylidene fluoride were subsequently added to the mortar. The mixture was ground again with a pestle for 30 min. Next, 0.5 mL of Nmethyl-2-pyrrolidone was added to the mixture to form a paste, which was subsequently coated onto copper foil and dried at 100 °C overnight. The cathode and anode were separated with a polypropylene film, and the electrolyte was 1 M LiPF6 in a 1:1 mixture of ethylene carbonate and diethylene carbonate. The battery was assembled in an argon-filled glovebox. Synthesis of [Co1.5L(H2O)4]n. A mixture of 0.3 mmol of Co(OH)2 (0.028 g), 0.3 mmol of H3L (0.055 g), and 10 mL H2O was placed in a 23 mL Teflon-lined stainless steel autoclave and heated at 160 °C for 3 days. Red block-shaped crystals were obtained after the mixture was cooled to room temperature. A yield of 80% based on Co(OH)2 was determined. Anal. Calcd for C7H10Co1.5NO9: C, 24.69; H, 2.96; N, 4.11%. Found: C, 24.52; H, 2.77; N, 4.32%.
Figure 1. (a) The zigzag chain structure. (b) Chains connected by hydrogen bonds. Green dotted lines indicate O−H···O hydrogen bonds.
The PXRD patterns of the as-synthesized [Co1.5L(H2O)4]n are in good agreement with the patterns simulated from the single-crystal data (Figure S2), indicating high phase purity of the material. The application of [Co1.5L(H2O)4]n as the anode material in LIBs features an outstanding charge−discharge efficiency and excellent capacity retention. It delivers a reversible capacity of 431 mAh g−1 at a current density of 50 mA g−1 (the theoretical capacity of [Co1.5L(H2O)4]n is 472 mAh g−1; see below), which is greater than the theoretical capacity of graphite (372 mAh g−1),30 as shown in Figure 2a,b. The irreversible discharge capacity loss from 1978 mAh g−1 for the first cycle to 869 mAh g−1 for the second cycle is a result of the formation of the solid electrolyte interphase (SEI) and the interfacial lithium storage caused by the decomposition of the electrolyte.31,32 The discharge capacity decreased from 869 to 431 mAh g−1 from the second cycle to the 10th cycle because water within the binder reacted with lithium to form Li2O inside the interlayers or on the surface.33,34 We also attempted to use anhydrous [Co1.5L]n for LIB studies; however, this material is not stable because of the removal of the water molecules, which leads to a much lower specific capacity (Figures S3 and S4). Actually, water molecules in the 3D supramolecular structure form hydrogen bonds, which reduce the activation energy for lithium-ion insertion and prevent organic solvent coinsertion.35,36 With respect to the cycle-life performance, the Coulombic efficiency appears to level off almost immediately and then remains constant at 98.3%, indicating that the SEI layer formed in the first cycle is fairly stable. To further examine the cycling stability and rate capacity, discharge and charge cycles were conducted at various current densities (Figure 2c). The discharge capacities were 392, 270, 155, and 85 mAh g−1 at rates of 50, 100, and 200 and 500 mA g−1, respectively. After enduring various charge− discharge rates, the electrode recovered its capacity of 383 mAh g−1 when the rate was returned to 50 mA g−1 and maintained this value without any apparent decay.
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RESULTS AND DISCUSSION Red block-shaped crystals of [Co1.5L(H2O)4]n (Figure 1) were synthesized from the reaction of Co(OH)2 and H3L.29 As shown in Figure 1a, two Co2+ ions are chelated and bridged by two L3− ligands to form a binuclear [Co2L2] grid, and these grids are further connected by one additional Co2+ ion through the oxygen at the 4-position of L3− to form a zigzag chain. The 3D structure is sustained by extensive hydrogen-bonding interactions through the chains (Figure 1b). The hydrogenbonding parameters are shown in Figure S1 and Table S1 in the Supporting Information. B
DOI: 10.1021/acs.inorgchem.6b00450 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry
Figure 2. Electrochemical performance of the electrode fabricated using [Co1.5L(H2O)4]n. (a) Charge−discharge curves at 50 mA g−1. (b) Cycling performance and Coulombic efficiency over 50 cycles. (c) Rate performance of [Co1.5L(H2O)4]n at different current densities. (d) Cyclic voltammograms between 0.01 and 3 V at a scan rate of 0.1 mV s−1.
Scheme 1. Electrochemical Redox Reaction Mechanism
of the Li + ions from Li 1 2 Co 3 L 2 (H 2 O) 8 to form Li6Co3L2(H2O)8. As shown in Scheme 1, the redox reaction of Li6Co3L2(H2O)8 at an average potential of 0.8 V proceeds between carboxylate carbonyl and enolate groups, accompanied by insertion/extraction of Li+ ions into/from Li6Co3L2(H2O)8/ Li12Co3L2(H2O)8.39 The aforementioned mechanism suggests that the oxidation state of cobalt could be changed during the charge−discharge process, which can be determined by Co 2p X-ray photoelectron spectroscopy (XPS). As shown in Figure 3a, the peaks at 776.6 and 791.7 eV are assigned to the characteristic peaks of Co2+.40 After the discharge process (Figure 3b), the Co 2p peaks shift to 778.3 and 793.7 eV, indicating the formation of Co0.41 The Co 2p spectrum after the charge process is consistent with the Co 2p spectrum of the original [Co1.5L(H2O)4]n sample. The XPS spectra of C 1s and O 1s were
To study the electrochemical redox reaction mechanism, cyclic voltammetry was conducted first. Figure 2d shows the cyclic voltammograms between 3.0 and 0.01 V vs Li+/Li. The reduction peak at 1.16 V is assigned to the insertion reaction of the Li+ ions into Co3L2(H2O)8 to form Li6Co3L2(H2O)8, and the oxidation peak emerging at 1.29 V corresponds to the extraction reaction of the Li+ ions from Li6Co3L2(H2O)8 to form Co3L2(H2O)8. The redox reaction of Co3L2(H2O)8 at an average potential of 1.22 V proceeds smoothly between Co2+ and Co0 (Co3L2(H2O)8/Li6Co3L2(H2O)8) (Scheme 1).37 The decrease in the electrode potential of Co2+/Co0 caused by the ligand field suggests that [Co1.5L(H2O)4]n is more favorable as an anode material.38 The second reduction peak at 0.45 V is assigned to the insertion reaction of Li + ions into Li6Co3L2(H2O)8 to form Li12Co3L2(H2O)8. The oxidation peak emerging at 1.0 V corresponds to the extraction reaction C
DOI: 10.1021/acs.inorgchem.6b00450 Inorg. Chem. XXXX, XXX, XXX−XXX
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conjugation effect from the aromatic ring and the enol structure, the carbon free radical is fairly stable. As shown in Figure 4a, an electron paramagnetic resonance (EPR) measurement showed the existence of a carbon free radical on the pyridine ring of the electrode material after the first discharge process, with a g value of 2.0064; this result further confirms the aforementioned mechanism. Stability is an important feature of coordination compounds as electrode materials during the charge and discharge processes. In our case, [Co1.5L(H2O)4]n exhibits specific structural features and excellent stability. First, the chelation of L3− to Co2 results in a relatively stable structure; second, according to crystal-field theory, Co1 (D4h coordination symmetry) and Co2 (D3h coordination symmetry) both have crystal-field stabilization energy before and after the charge and discharge processes, which is beneficial to both stability and reversibility. The magnetic measurement gave a χMT value of 5.12 cm3 mol−1 K at 300 K, which is larger than the spin-only value of 2.82 cm3 mol−1 K for 1.5 free CoII ions as a consequence of the orbital contribution (Figure 4b). The magnetic result is consistent with the energy splitting diagrams in Figure 4c. Recently, metal−organic frameworks (MOFs), a new generation of porous crystalline materials constructed via coordination bonds between inorganic nodes and organic linkers, have attracted great attention as electrode materials.22−25,44,45 For example, FeIII(OH)0.8F0.2(BDC)·H2O (MIL53(Fe)) was studied as a cathode and exhibited a reversible capacity of 75 mAh g−1. Its relatively low energy density is a consequence of its poor electronic conductivity.44 MOF-177 (Zn4O(BTB)2, BTB3− = 1,3,5-benzenetribenzoate), with a very high Brunauer−Emmett−Teller (BET) surface area of 5250 m2 g−1, was studied for lithium storage as an anode material; however, structural destruction after lithium storage was
Figure 3. Co 2p XPS spectra of [Co1.5L(H2O)4]n as-synthesized (a) and following an electrochemical discharge process (b) and a charge process (c) from the assembled electrodes.
further studied (Figure S5). In the C 1s XPS spectra, the energies of 284.8, 288.6, and 290.6 eV are assigned to the characteristic peaks of the carbon−carbon double bond of the pyridine ring, the carboxyl, and the carbon−carbon single bond of the enol structure, respectively.42 According to the spectra, the carboxyl of L3− transformed into the enol structure after the discharge process and appeared again after the charge process, suggesting a reversible interconversion between the carboxyl and enol structures. In the O 1s spectra, the two peaks at 531.6 and 529.3 eV are assigned to the lattice oxygen and the Li−O bond, respectively.43 These results suggest that the Li−O bond formed after gaining electrons and disappeared after losing electrons, further indicating the reversibility of the redox reaction. In the discharge process, the carboxylic oxygen acquires an electron, which leads to the formation of a carbon− carbon double bond. The electron on the pyridine ring results in the formation of a carbon free radical. Because of the
Figure 4. (a) EPR spectrum of the electrode material after the first discharge process at 298 K. (b) Plot of χMT vs T at an applied field of 1000 Oe. (c) Energy splitting diagrams for the d orbitals of Co1 and Co2 in the charge and discharge processes. D
DOI: 10.1021/acs.inorgchem.6b00450 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry observed.45 As an alternative, the 1D metal−organic material [Co1.5L(H2O)4]n exhibited a reversible capacity of 431 mAh g−1 when used as an anode material for LIBs; this capacity is much higher than the theoretical capacity of graphite (372 mAh g−1). Such 1D metal−organic materials benefit from the following two structural features: (i) dual redox-active metal centers and organic components are coordinated together for enhanced insolubility and stability in the electrolyte, and (ii) the coordination chains are weakly connected via hydrogen bonds, which provide channels for lithium-ion diffusion.
(2) (a) Scrosati, B.; Hassoun, J.; Sun, Y. K. Energy Environ. Sci. 2011, 4, 3287−3295. (b) Dunn, B.; Kamath, H.; Tarascon, J. M. Science 2011, 334, 928−935. (3) Armand, M.; Tarascon, J. M. Nature 2008, 451, 652−657. (4) Morozan, A.; Jaouen, F. Energy Environ. Sci. 2012, 5, 9269−9290. (5) (a) Sun, W. W.; Wang, Y. Nanoscale 2014, 6, 11528−11552. (b) Nam, K. T.; Kim, D.-W.; Yoo, P. J.; Chiang, C.-Y.; Meethong, N.; Hammond, P. T.; Chiang, Y.-M.; Belcher, A. M. Science 2006, 312, 885−888. (6) (a) Poizot, P.; Laruelle, S.; Grugeon, S.; Dupont, L.; Tarascon, J.M. Nature 2000, 407, 496−499. (b) Bruce, P. G.; Scrosati, B.; Tarascon, J. M. Angew. Chem., Int. Ed. 2008, 47, 2930−2946. (7) Werner, H. Angew. Chem., Int. Ed. 2013, 52, 6146−6153. (8) Belda, O.; Moberg, C. Coord. Chem. Rev. 2005, 249, 727−740. (9) Dhakshinamoorthy, A.; Garcia, H. Chem. Soc. Rev. 2014, 43, 5750−5765. (10) Tascon, R. E.; Colston, M. J.; Ragno, S.; Stavropoulos, E.; Gregory, D.; Lowrie, D. B. Nat. Med. 1996, 2, 888−892. (11) Gill, M. R.; Thomas, J. A. Chem. Soc. Rev. 2012, 41, 3179−3192. (12) Sigel, H. Chem. Soc. Rev. 2004, 33, 191−200. (13) Zhou, H. C.; Long, J. R.; Yaghi, O. M. Chem. Rev. 2012, 112, 673−674. (14) Kim, H.; Jeong, G.; Kim, Y.-U.; Kim, J.-H.; Park, C.-M.; Sohn, H.-J. Chem. Soc. Rev. 2013, 42, 9011−9034. (15) Cao, R. G.; Xu, W.; Lv, D. P.; Xiao, J.; Zhang, J. G. Adv. Energy Mater. 2015, 5, 1402273. (16) Wang, S. W.; Wang, L. J.; Zhu, Z. Q.; Hu, Z.; Zhao, Q.; Chen, J. Angew. Chem., Int. Ed. 2014, 53, 5892−5896. (17) Nishide, H.; Oyaizu, K. Science 2008, 319, 737−738. (18) Wang, S. W.; Wang, L. J.; Zhang, K.; Zhu, Z. Q.; Tao, Z. L.; Chen, J. Nano Lett. 2013, 13, 4404−4409. (19) Zhu, Z. Q.; Hong, M. L.; Guo, D. S.; Shi, J. F.; Tao, Z. L.; Chen, J. J. Am. Chem. Soc. 2014, 136, 16461−16464. (20) Wu, J. S.; Rui, X. H.; Wang, C. Y.; Pei, W. B.; Lau, R.; Yan, Q. Y.; Zhang, Q. C. Adv. Energy Mater. 2015, 140, 2189−2212. (21) Song, Z. P.; Zhou, H. S. Energy Environ. Sci. 2013, 6, 2280− 2301. (22) Wu, D.; Guo, Z.; Yin, X.; Pang, Q.; Tu, B.; Zhang, L.; Wang, Y. G.; Li, Q. Adv. Mater. 2014, 26, 3258−3262. (23) Zhang, Z. Y.; Yoshikawa, H.; Awaga, K. J. Am. Chem. Soc. 2014, 136, 16112−16115. (24) Aubrey, M. L.; Long, J. R. J. Am. Chem. Soc. 2015, 137, 13594− 13602. (25) Maiti, S.; Pramanik, A.; Manju, U.; Mahanty, S. ACS Appl. Mater. Interfaces 2015, 7, 16357−16363. (26) Zhao, B.; Chen, X. Y.; Cheng, P.; Liao, D. Z.; Yan, S. P.; Jiang, Z. H. J. Am. Chem. Soc. 2004, 126, 15394−15395. (27) Zhao, B.; Gao, H. L.; Chen, X. Y.; Cheng, P.; Shi, W.; Liao, D. Z.; Yan, S. P.; Jiang, Z. H. Chem. - Eur. J. 2006, 12, 149−158. (28) Gao, H. L.; Yi, L.; Zhao, B.; Zhao, X. Q.; Cheng, P.; Liao, D. Z.; Yan, S. P. Inorg. Chem. 2006, 45, 5980−5988. (29) Zhao, X. Q.; Li, Y. C. J. Coord. Chem. 2013, 66, 937−945. (30) Yoshio, M.; Wang, H. Y.; Fukuda, K. J. Angew. Chem., Int. Ed. 2003, 35, 4203−4206. (31) Sahay, R.; Kumar, P. S.; Aravindan, V.; Sundaramurthy, J.; Ling, W. C.; Mhaisalkar, S. G.; Ramakrishna, S.; Madhavi, S. J. Phys. Chem. C 2012, 116, 18087−18092. (32) Wang, J.; Yang, N.; Tang, H.; Dong, Z.; Jin, Q.; Yang, M.; Kisailus, D.; Zhao, H.; Tang, Z.; Wang, D. Angew. Chem., Int. Ed. 2013, 52, 6417−6420. (33) Torardi, C.; Miao, C. R. Chem. Mater. 2002, 14, 4430−4433. (34) Gao, S. K.; Chen, Z. J.; Wei, M. D.; Wei, K. M.; Zhou, H. S. Electrochim. Acta 2009, 54, 1115−1118. (35) Nam, K. W.; Kim, S.; Lee, S.; Salama, M.; Shterenberg, I.; Gofer, Y.; Kim, J. S.; Yang, E.; Park, C. S.; Kim, J. S.; Lee, S. S.; Chang, W. S.; Doo, S. G.; Jo, Y. N.; Jung, Y.; Aurbach, D.; Choi, J. W. Nano Lett. 2015, 15, 4071−4079. (36) Kim, H.; Hong, J.; Park, K. Y.; Kim, H.; Kim, S. W.; Kang, K. Chem. Rev. 2014, 114, 11788−11827.
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CONCLUSIONS We introduced a 1D metal−organic material as a new anode active material for LIBs. The coordination between the metal ion and the bridging ligand forming the chain structure not only enhances the solid-state stability for LIBs but also provides a means of combining two active elements for enhanced electrochemical performance. The independent redox-active sites on both the metal and ligand simultaneously contribute to the capacity, resulting in a stable capacity and a Coulombic efficiency of 98.3%. The newly discovered mechanism, in which both cobalt ions and ligands are involved in the redox reaction, was revealed via the Co 2p, C 1s, and O 1s XPS spectra, EPR spectra, and cyclic voltammograms, which are consistent with the observed electrochemical performance. Moreover, the introduction of a coordination chain compound does not have restrictions for the size of the channels of the electrode materials for lithium-ion diffusion and prevents the solubility of active organic substances in the electrolyte during the charging and discharging processes. This result provides a new approach to use coordination compounds as electrode materials for rechargeable LIBs. Continuations of the research using 1D/2D metal−organic materials with both a redox-active metal ion and a ligand are underway in our lab.
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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.6b00450. Interchain hydrogen bonding in [Co1.5L(H2O)4]n, cycling performance of [Co1.5L]n, and PXRD, TGA, FTIR, XPS, and SEM data (PDF)
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AUTHOR INFORMATION
Corresponding Authors
*E-mail:
[email protected]. *E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work was supported by the NNSFC (Grants 21331003, 21373115, and 91422302) and the MOE (Grants NCET-130305 and IRT-13R30).
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REFERENCES
(1) (a) Vu, A.; Qian, Y. Q.; Stein, A. Adv. Energy Mater. 2012, 2, 1056−1085. (b) Morozan, A.; Jaouen, F. Energy Environ. Sci. 2012, 5, 9269−9290. E
DOI: 10.1021/acs.inorgchem.6b00450 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry (37) Fawcett, W. R. Langmuir 2008, 24, 9868−9875. (38) Nguyen-Thanh, D.; Frenkel, A. I.; Wang, J. Q.; O’Brien, S.; Akins, D. L. Appl. Catal., B 2011, 105, 50−60. (39) Armand, M.; Grugeon, S.; Vezin, H.; Laruelle, S.; Ribière, P.; Poizot, P.; Tarascon, J.-M. Nat. Mater. 2009, 8, 120−125. (40) Sun, C. C.; Yang, J.; Rui, X. H.; Zhang, W. N.; Yan, Q. Y.; Chen, P.; Huo, F. W.; Huang, W.; Dong, X. C. J. Mater. Chem. A 2015, 3, 8483−8488. (41) Wagner, C. D.; Riggs, W. M.; Davis, L. E.; Moulder, J. F.; Muilenberg, G. E. Handbook of X-ray Photoelectron Spectroscopy; PerkinElmer: Waltham, MA, 1978. (42) Beamson, G.; Briggs, D. High Resolution XPS of Organic Polymers: The Scienta ESCA300 Database; Wiley: Chichester, U.K., 1992. (43) Marco, J. F.; Gancedo, J. R.; Gracia, M.; Gautier, J. l.; Ríos, E.; Berry, F. J. J. Solid State Chem. 2000, 153, 74−81. (44) Férey, G.; Millange, F.; Morcrette, M.; Serre, C.; Doublet, M. L.; Grenèche, J. M.; Tarascon, J. M. Angew. Chem., Int. Ed. 2007, 46, 3259−3263. (45) Li, X. X.; Cheng, F. Y.; Zhang, S. N.; Chen, J. J. Power Sources 2006, 160, 542−547.
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