CoâIntercalation of Mg2+ Ions into Graphite for MagnesiumâIon

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Co–Intercalation of Mg Ions into Graphite for Magnesium–Ion Batteries Dong-Min Kim, Sung Chul Jung, Seongmin Ha, Youngjin Kim, Yuwon Park, Ji Heon Ryu, Young-Kyu Han, and Kyu Tae Lee Chem. Mater., Just Accepted Manuscript • DOI: 10.1021/acs.chemmater.8b00288 • Publication Date (Web): 01 May 2018 Downloaded from http://pubs.acs.org on May 1, 2018

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Chemistry of Materials

Co–Intercalation of Mg2+ Ions into Graphite for Magnesium–Ion Batteries Dong–Min Kim†, Sung Chul Jung§, Seongmin Ha†, Youngjin Kim†, Yuwon Park†, Ji Heon Ryuǂ, Young–Kyu Han*,‡, Kyu Tae Lee*,† †

School of Chemical and Biological Engineering, Institute of Chemical Processes, Seoul National University, 1, Gwanak–ro, Gwanak–gu, Seoul 08826, Republic of Korea § Department of Physics, Pukyong National University, Busan 48513, Republic of Korea ǂ Graduate School of Knowledge–based Technology and Energy, Korea Polytechnic University, Gyeonggi, 15073, Republic of Korea ‡ Department of Energy and Materials Engineering, Dongguk University–Seoul, Seoul 04620, Republic of Korea ABSTRACT: Magnesium batteries have been studied as an alternative to lithium–ion batteries because of their potentially lower cost and better safety. However, magnesium batteries suffer from drawbacks such as a few electrolytes suitable for the reversible magnesium metal anode. In this connection, alloy anode materials have been examined to use various electrolytes and have shown good reversibility through an alloying mechanism. However, alloy materials also have issues such as poor capacity retention due to the pulverization of alloy materials during cycling. Therefore, graphite is introduced as an intercalation host for Mg2+ ions through the solvated-Mg-ion intercalation mechanism. Mg2+ ions are reversibly co–intercalated into graphite with linear ether solvents such as dimethoxyethane and diethylene glycol dimethyl ether through the solvated–Mg–ion intercalation mechanism. The co– intercalation of Mg2+ and ether solvents into graphite is supported by ex situ XRD and FT–IR analyses. The first principles calculations suggest the Mg2+–DEGDME double–layer structure in graphite as the structural model of co–intercalation. These findings may provide a new avenue for developing promising anodes in magnesium–ion batteries.

With the development of electric vehicles and smart grid systems, safety and cost of power sources are becoming increasingly important, and in this regard, magnesium batteries have been considered as a promising power source because they are safer and less expensive than commercialized lithium–ion batteries.1-5 In particular, the absence of dendrite formation on the Mg metal anode is one of the most attractive aspects of magnesium batteries, leading to their excellent safety.6-7 However, there are still challenges for realizing practical magnesium batteries. One challenge is the difficulty of utilizing Mg metal as an anode with various electrolytes. Mg metal cannot be used with the conventional carbonate– based solvents used for commercialized lithium–ion batteries, because when Mg metal is in contact with those solvents, a thick passivation layer forms on the surface of Mg metal, leading to failures in the stripping and plating of magnesium.89 The stripping and plating of magnesium metal is only reversible in a few electrolytes such as organohaloaluminate salt in tetrahydrofuran (THF),1 magnesium aluminum chloro complex in THF,10 and ether–based electrolytes containing magnesium salt and chloride additives.11-12 Moreover, organohaloaluminate salts are acidic, leading to corrosive reaction.13-14 In order to overcome these limitations, Mg2+ insertion materials have been considered as alternatives to Mg metal, because they exhibit reversible insertion and de–

insertion of Mg2+ with conventional electrolytes. For example, Mg2+ ions are reversibly inserted into Bi and Sn.15-16 However, although those alloy materials exhibited the reversible insertion of Mg2+, they suffered drawbacks such as pulverization due to a large volume change during cycling, resulting in poor cycle performance. Therefore, in this study, we introduce graphite as a host material for the intercalation of Mg2+ ions. Unlike the alloying reaction mechanism of alloy materials such as Bi and Sn, graphite stores Mg2+ ions through an intercalation reaction mechanism, leading to no pulverization. Moreover, the Mg2+ ions are reversibly co– intercalated into graphite with linear ether solvents, leading to promising electrochemical performance. Graphite is a commercialized anode material for lithium–ion batteries, and Li+ ions are reversibly intercalated into graphite through a staging phenomenon.17 In addition to Li+ ions, graphite was recently reported to store K+ ions.18-19 However, graphite is known not to store Na+ and Mg2+ ions.20-21 The Li+ and K+ ions exhibit effective ionic radii of 0.76 and 1.38 Å, respectively, whereas the Mg2+ and Na+ ions exhibit smaller radii of 0.72 and 1.02 Å, respectively. This suggests that ion storage capability of graphite does not depend on ionic size. The negligible capacities of graphite for Na+ and Mg2+ can be ascribed to the weakest chemical binding of Na+ and Mg2+ among alkali metals and alkaline earth metals, respectively, to

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various carbon substrates such as graphene and its derivatives with defects or substitutional foreign atoms.22 This implies that the thermodynamic interaction between the Na+, Mg2+ ions and graphite is too weak to allow the intercalation of these ions into graphite. However, a new mechanism has recently been introduced to store Na+ ions in graphite, in which Na+ ions are co–intercalated into graphite with diethylene glycol dimethyl ether (DEGDME) solvents through the solvated–Na– ion intercalation mechanism.23-24 The solvent co–intercalation behavior into graphite was also observed in lithium–ion batteries.25-27 Inspired by the solvated–M–ion intercalation mechanism (M = Li, Na and K), we investigated the co–intercalation of Mg2+ ions into graphite with linear ether solvents such as DEGDME and DME. We performed density functional theory (DFT) calculations to examine the possibility of the co–intercalation of Mg2+ ions into graphite with linear ether solvents, suggesting that the thermodynamically favorable intercalation of Mg2+ ions into graphite through the co–intercalation mechanism. The binding energy between a Mg2+ ion and solvents was calculated to investigate the possibility of co– intercalation of Mg2+ into graphite with solvents. A carrier ion should strongly bind to solvent molecules for successful co– intercalation without ion desolvation at the electrolyte/electrode interface. The binding energies were calculated to be 10.05, 7.60, 5.87, and 5.69 eV for the Mg2+– DEGDME, Mg2+–DME, Mg2+–EC, and Mg2+–DEC complexes, respectively (Figure S1). DME, EC and DEC denote 1,2-dimethoxyethane, ethylene carbonate and diethylene carbonate, respectively. This indicates that the solvation of Mg2+ by DEGDME is most energetically favorable, suggesting that DEGDME is the most advantageous for successful co–intercalation. However, the binding energies for Mg2+–EC and Mg2+–DEC complexes are significantly smaller by at least 4 eV than that for Mg2+–DEGDME complex, implying that Mg2+ desolvation at the electrolyte/graphite interface can occur easily with the EC/DEC electrolyte. This suggests a possibility of unsuccessful co–intercalation of Mg2+–EC and Mg2+–DEC complexes. We also examined the graphite structure co– intercalated with the Mg2+–DEGDME complex. Figure 1a,b shows the single– and double–layer structures of Mg2+– DEGDME complexes in the interlayer space of graphite. We calculated the co–intercalation energies, defined as Eint = [Etot([Mg–solvent]n graphite) nEtot(Mg–solvent) Etot(graphite)]/n, where Etot([Mg–solvent]n graphite), Etot(Mg– solvent), Etot(graphite), and n are the total energy of Mg– solvent co–intercalated graphite, the total energy of a single Mg–solvent complex, the total energy of graphite, and the number of Mg–solvent complexes per unit cell, respectively. The calculated co–intercalation energies are -0.49 and -0.60 eV per Mg2+–DEGDME complex for the single– and double– layer structures, respectively, revealing that the Mg2+– DEGDME into graphite is thermodynamically favorable and that the formation of double–layer structure is preferred over that of single–layer structure. The double–layer intercalant structure was also observed in the study of Na+–DEGDME co–intercalated graphite.28 Therefore, the calculated Mg2+– solvent binding energies and intercalation energies suggest that the Mg2+–DEGDME double–layer structure is the most reasonable structural model of co–intercalation.

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We further investigated the co–diffusion behavior of the Mg2+–DEGDME complex in the double–layer structure shown in Figure 1c. In the co–diffusion process, a DEGDME molecule surrounding one Mg2+ ion moves parallel to the graphene surface. The diffusion barrier was calculated by using the nudged elastic band (NEB) method, and the calculated diffusion barrier (0.21 eV) was used to evaluate the diffusivity at T = 300 K. It is worth noting that the diffusivity of the Mg2+–DEGDME complex in graphite (1.5 × 10–8 cm2 s–1) is comparable to that of Li+ in graphite (1.8 × 10–9 cm2 s–1) (see detailed procedures in the Supporting Information). This indicates that Mg2+–solvent co–intercalated graphite can perform as well as Li+ intercalated graphite. To demonstrate the co–intercalation of Mg2+ ions into graphite with linear ether solvents, galvanostatic cycling was performed with 0.3 M Mg(TFSI)2 in EC/DEC (1:1 v:v) and 0.3 M Mg(TFSI)2 in DME/DEGDME (1:1 v:v). Mg(TFSI)2 in glyme is known to exhibit the reversible stripping and plating of Mg, although it exhibits high polarization during cycling.28 Therefore, natural graphite and Mg metal were used as the working and counter electrodes for 0.3 M Mg(TFSI)2 in DME/DEGDME. However, carbonate solvents such as EC and DEC are decomposed on the surface of Mg metal to form a thick passivation layer that prevents Mg dissolution. This indicates that carbonate–based electrolytes cannot be used with a Mg metal electrode. As a result, an activated carbon (AC) electrode was used for 0.3 M Mg(TFSI)2 in EC/DEC as the counter electrode instead of Mg metal.29-30 As predicted in the DFT calculation, the graphite electrode with 0.3 M Mg(TFSI)2 in EC/DEC delivered a negligible amount of

Figure 1. Mg2+–DEGDME co–intercalated graphite: (a) single– and (b) double–layer structures. Orange, white, gray, and red balls represent Mg, H, C, and O atoms, respectively. di represents the intercalant gallery height. (c) Energies and structures of the intermediate states obtained from the DFT–NEB calculations of the Mg2+–DEGDME co–diffusion in graphite.


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reversible capacity, indicating that Mg2+ ions cannot be intercalated into graphite (Figure 2a). However, remarkably, the graphite electrode with 0.3 M Mg(TFSI)2 in DME/DEGDME exhibited reversible voltage profiles, when the graphite/Mg cell with 0.3 M Mg(TFSI)2 in DME/DEGDME was galvanostatically cycled with a constant capacity of 180 mA h g-1, corresponding to [Mg2+DEGDME]0.24C6 (Figure 2b). The equilibrium redox potential of graphite for the intercalation of Mg2+ ions was approximated by measuring the quasi–open–circuit–voltage (QOCV) during charge and discharge through a galvanostatic intermittent titration technique (GITT). The redox potential of graphite was about 1 V vs. Mg/Mg2+, as shown in Figure 2c, indicating that the plateau at about 0 V during discharge was not due to the electroplating of Mg metal. Figure 2b,c shows that the graphite/Mg cell exhibited a large polarization of about 2 V at the specific current of 2 mA g-1, which is attributed to the sluggish kinetics of the Mg metal electrode with ether–based electrolytes. To measure each contribution of the graphite and Mg electrodes to the large polarization of the graphite/Mg cell, two symmetric cells, one comprising two graphite/graphite electrodes and the other with two Mg/Mg electrodes, were examined by galvanostatic cycling (Figure 2d and Figure S2). To fabricate the graphite/graphite symmetric cell, two graphite electrodes were first discharged with Mg counter electrodes. The discharge capacity was controlled to deliver only 90 mA h g-1, leading to the formation of magnesium graphite intercalation compounds (Mg–GIC) partially intercalated with Mg2+ ions. Then, we disassembled the cells to recover the Mg–GIC electrodes, followed by reassembling symmetric cell using the recovered electrodes. The Mg–GIC/Mg–GIC symmetric cell was glavanostatically cycled with a constant capacity of 40 mA h g-1 and exhibited a small polarization of about 0.2 V at the specific current of 2 mA g-1, as shown in Figure 2d. This value is much smaller than the polarization of the graphite/Mg cell. However, a large polarization was observed for the Mg/Mg symmetric cell with 0.3 M Mg(TFSI)2 in DME/DEGDME (Figure S2), similar to the value of the graphite/Mg cell. Therefore, the large polarization of the graphite/Mg cell can be attributed mostly to the Mg metal electrode. This suggests that the kinetics of intercalation and de–intercalation of Mg2+ ions in graphite is fast enough for magnesium–ion battery electrodes. This is consistent with the high diffusivity of the Mg2+–DEGDME complex in graphite obtained from the DFT–NEB calculations. Moreover, the Mg–GIC/Mg–GIC symmetric cell showed good reversibility over 15 cycles (Figure 2d). To clarify the co–intercalation mechanism of graphite with Mg2+ ions, structural changes in the natural graphite were investigated through ex situ XRD analysis at various states during charge and discharge (Figure 3). The (002) graphite peak located at 26.55° gradually disappeared and split into new two peaks at 25.11° and 28.89° as the Mg2+ ions were intercalated into graphite during discharge. Then, as Mg2+ ions are de–intercalated from the graphite during charge, the two split peaks reversibly merged into one peak located at the (002) peak position of the pristine graphite. The intercalation of Mg2+ ions causes the XRD peak splitting as a result of the periodic sequence of graphite and intercalant Mg2+ planes, which is known as the staging behavior of graphite.19, 24

Figure 2. (a) Voltage profiles of graphite electrodes with 0.3 M Mg(TFSI)2 in EC/DEC (1:1 v:v). (b) Voltage profiles and (c) galvanostatic intermittent titration technique (GITT) profiles of graphite electrodes with 0.3 M Mg(TFSI)2 in DME/DEGDME (1:1 v:v). (d) Voltage profiles of the Mg–GIC/Mg–GIC symmetric cell with 0.3 M Mg(TFSI)2 in DME/DEGDEME.

When the two newly appeared peaks are indexed as (00l) and (00l+1), the value of l can be calculated using equations (1) and (2) which originate from Bragg’s law, resulting in an l value of 6.803.24, 31 Therefore, two peaks located at 25.11° and 28.89° are considered to correspond to the (007) and (008) peaks, respectively, and the c lattice parameter of Mg–GIC is about 24.85 Å after full discharge.

d and d

·················· (1) ·················· (2)

IC is the c lattice parameter of Mg–GIC at each stage, and d(00l) and d(00l+1) are the d–spacing values of the (00l) and (00l+1) planes, respectively. The XRD peak was also observed to broaden and shift during charge and discharge, which is ascribed to the staging disorder of Mg–GIC.32 Although the graphite exhibited reversible intercalation and de–intercalation of Mg2+ ions, Figure 3b reveals that the (002) peak of graphite did not fully recover its original position ((vii) in Figure 3b) after charge, despite the fact that the graphite delivered more charge capacity than discharge capacity. This is attributed to oxidative electrolyte decomposition during charge. Moreover, we calculated the intercalant gallery heights of the Mg2+–solvent complexes after full discharge for various stage numbers (Table S1), and then, compared with the theoretical values obtained from the DFT calculations (Table S2). We found that the intercalant gallery height (di = 11.10 Å) of the double–layer structure where Mg2+–DEGDME complexes are double stacked in Figure 1b is close to the experimentally estimated value (11.45 Å) of the fully discharged graphite obtained from the ex situ XRD analysis, assuming that the stage number is five. d c lattice parameter !24.85 Å( ) pristine lnterlayer distance !3.35 Å( - n ) 1


······ (3)


Figure 3. (a) Voltage profiles of the graphite electrode with 0.3 M Mg(TFSI)2 in DME/DEGDME (1:1 v:v) and (b) the ex situ XRD patterns of graphite electrodes obtained at various points indicated in (a).

This suggests that Mg2+ is co–intercalated in graphite in the form of the Mg2+–DEGDME double–layer structure. The intercalant gallery height can be estimated using eqation (3) if a stage number, n, is specified. Moreover, ex situ Fourier transform infrared spectroscopy (FT–IR) analysis was performed to support the formation of Mg–GICs co–intercalated with linear ether solvents such as DEGDME. To carefully clarify the co–intercalation of the Mg2+–solvent complex with the absence of the effect of residual electrolyte solvents in the pores of the electrodes, we assembled two identical coin cells comprising the graphite and Mg metal electrodes with the 0.3 M Mg(TFSI)2 in DME/DEGDEME electrolyte. While one cell was discharged to intercalate the Mg2+–solvent complex, the other cell remained at rest without applying a current. Then, both cells were dissembled, and the graphite electrodes were washed with DME solvents. Figure 4a shows the FT–IR spectra of the electrolyte (0.3 M Mg(TFSI)2 in DME/DEGDEME) and a bare solvent mixture (DME and DEGDME), in which the bands near 2900 cm-1 and 1100 cm-1 are assigned to the characteristic bands of DME and DEGDME. These bands were detected in the discharged graphite electrode, but not in the rested graphite electrode, as shown in Figure 4b. This supports that the linear ether solvents were co–intercalated into graphite with Mg2+ ions, which is consistent with the theoretical calculations. In addition, we performed ex situ high-resolution transmission electron microscopy (HR-TEM) analysis with energy dispersive X-ray spectroscopy (EDS). The TEM and EDS mapping images of graphite electrodes were compared before and after full discharge (Figure S3). The well-aligned layered structure of graphite was observed in the pristine

Figure 4. FT–IR spectra of (a) 0.3 M Mg(TFSI)2 in DME/DEGDEME electrolyte and a bare solvent mixture (DME and DEGDME), and (b) graphite electrodes before and after the intercalation of Mg2+ ions.

intercalation, partially disordered graphite was observed (Figure S3b, c), and Mg atoms were detected in the disordered region (Figure S3d-f). This suggests that graphite was partially amorphized after Mg2+ intercalation. The discharge capacity of graphite (180 mA h g-1) is considered to be too high for the stage number 5.24 Therefore, we propose that the high capacity of graphite is attributed to the additional intercalation of Mg2+ ions to Mg-GIC after the stage 5, resulting in the formation of the disordered Mg-GIC. In summary, natural graphite was examined as a host material for the intercalation of Mg2+ ions in magnesium–ion batteries. Mg2+ intercalation into graphite is known to be thermodynamically unfavorable. However, we report for the first time that Mg2+ ions are reversibly co–intercalated into graphite with linear ether solvents such as DME and DEGDME. The solvated–Mg–ion intercalation mechanism of graphite was supported by ex situ XRD and FT–IR analyses. Moreover, our DFT calculations suggested the Mg2+– DEGDME double–layer structure in graphite as the structural model of co–intercalation. Finally, we believe that this finding provides a new avenue for developing promising anode materials for magnesium–ion batteries.

ASSOCIATED CONTENT Supporting Information. Experimental methods, calculation details for diffusivity, descriptions for Mg2+–solvent solvating


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structure, Mg2+–DEGDME intercalated graphite structure, electrochemical data, and ex situ TEM analysis are available on supporting information. This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION Corresponding Author *Kyu Tae Lee: [email protected] *Young–Kyu Han: [email protected]

ACKNOWLEDGMENT This work was supported by Samsung Research Funding Center of Samsung Electronics under Project Number SRFC–TA1403– 05 and the National Research Foundation of Korea (NRF) Grant (No. NRF– 2016R1A2B3015956).

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