Hydrate-melt electrolytes for high-energy-density aqueous batteries

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1 ARTICLE NUMBER: Hydrate-melt electrolytes for high-energy-density aqueous batteries Yuki Yamada, 1,2 Kenji Usui, 1 Keitaro Sodeyama, 2,3,4 Seongjae Ko, 1 Yoshitaka Tateyama, 2,4 and Atsuo Yamada 1,2, * 1 Department of Chemical System Engineering, The University of Tokyo, 7-3-1, Hongo, Bunkyoku, Tokyo , Japan 2 Elements Strategy Initiative for Catalysts & Batteries (ESICB), Kyoto University, 1-30, Goryo- hara, Nishikyo-ku, Kyoto , Japan 3 PREST, Japan Science and Technology Agency (JST), 4-1-8, Honcho, Kawaguchi, Saitama , Japan 4 Center for Green Research on Energy and Environmental Materials and Center for Materials Research by Information Integration, National Institute for Materials Science (NIMS), 1-1, Namiki, Tsukuba, Ibaraki , Japan Corresponding Author * yamada@chemsys.t.u-tokyo.ac.jp NATURE ENERGY 1

2 log ( / S cm 2 mol 1 ) 1, superionic (decoupled) hydrate melt LiTFSI 22 mol kg 1 (sat.) good-ionic (full dissociation) KCl line 1.2 mol kg 1 poor-ionic (associated) non-ionic log ( 1 / Poise 1 ) Supplementary Figure 1 Walden plot for the Li(TFSI)0.7(BETI)0.3 2H2 hydrate melt (red closed circle) and LiTFSI/H2 solutions (blue closed circles) at various concentrations at 30 C. The Walden plot was generated from the molar conductivities ( ) and viscosities ( ) of the electrolyte solutions. In a Walden plot, electrolyte solutions can be classified in terms of their performance as ionic conductors: superionic (upper left region above the ideal KCl line), good-ionic (on the ideal line), poor-ionic (bottom right region under the ideal line), or non-ionic (far below the ideal line) liquids and solutions. 1,2 For the LiTFSI/H2 solutions, the plot approaches the ideal line with increasing concentration, and finally joins with the ideal line at saturation. By contrast, the hydrate melt lies above the ideal line and is thus classified as a superionic solution, in which ionic conductivity and viscosity are decoupled. 3 2 NATURE ENERGY

3 SUPPLEMENTARY INFRMATIN 0.8 solvating at all times N Li, average LiTFSI K 3-LiBETI 100,00teps 20-H ID number of H 2 molecules Supplementary Figure 2 Average numbers (NLi,average) of Li + ions coordinated with each water molecule during 100,000 DFT-MD steps. Here, a water molecule with its atom located within 0.25 nm of a Li + ion is defined as a solvating molecule. All water molecules in the supercell of 7-LiTFSI/3- LiBETI/20-H2 were tracked over 100,00imulation steps and then labelled with identification numbers from 1 to 20 in descending order of their NLi,average values. An NLi,average value of indicates that the water molecule was coordinated with Li + during all simulation steps. For 7-LiTFSI/3- LiBETI/20-H2, all water molecules had NLi,average values that were equal or quite close to, suggesting that free water molecules (outside Li + hydration shells) are rare in this liquid and thus that it can be classified as a hydrate melt. NATURE ENERGY 3

4 Current / ma Current / ma mol kg 1 LiTFSI/H 2 /Zn 0.1 mv s 1 Li(TFSI) 0.7 (BETI) 0.3 2H 2 /Zn 0.1 mv s Potential / V vs. Li + /Li Supplementary Figure 3 Cyclic voltammograms of Li4Ti512 on a Zn current collector in a typical aqueous solution of 1.2 mol kg 1 LiTFSI/H2 and in a hydrate melt electrolyte of Li(TFSI)0.7(BETI)0.3 2H2 at 0.1 mv s 1. 4 NATURE ENERGY

5 SUPPLEMENTARY INFRMATIN a Current / ma - Sat. Li 2 S 4 /H mv s Potential / V vs. Li + /Li b 5.0 Sat. LiN 3 /H 2 Current / ma mv s Potential / V vs. Li + /Li Supplementary Figure 4 Cyclic voltammograms of Li4Ti512 (on a Zn current collector) in (a) saturated Li2S4/H2 (2.9 mol kg 1 ) and (b) saturated LiN3/H2 (12.6 mol kg 1 ) electrolytes at 0.1 mv s 1. NATURE ENERGY 5

6 a 5.2 mol kg 1 LiTFSI/H 2 Current / ma mv s Potential / V vs. Li + /Li b Current / ma mol kg 1 LiTFSI/H mv s Potential / V vs. Li + /Li Supplementary Figure 5 Cyclic voltammograms of Li4Ti512 (on a Zn current collector) in (a) 5.2 and (b) 9.5 mol kg 1 LiTFSI/H2 electrolytes at 0.1 mv s 1. 6 NATURE ENERGY

7 SUPPLEMENTARY INFRMATIN a Current / ma mol kg 1 Li(TFSI) 0.7 (BETI) 0.3 /H mv s Potential / V vs. Li + /Li b 11 mol kg 1 Li(TFSI) 0.7 (BETI) 0.3 /H 2 Current / ma mv s Potential / V vs. Li + /Li Supplementary Figure 6 Cyclic voltammograms of Li4Ti512 (on a Zn current collector) in (a) 1.1 and (b) 11 mol kg 1 Li(TFSI)0.7(BETI)0.3/H2 electrolytes at 0.1 mv s 1. NATURE ENERGY 7

8 a LiCo 2 (10C) 25 o C 1st b LiNi 0.5 Mn (6.8C) 25 o C 1st Supplementary Figure 7 Initial charge-discharge voltage profiles of two Li-ion full cells, (a) LiCo2/Li4Ti512 and (b) LiNi0.5Mn1.54/ Li4Ti512, with Li(TFSI)0.7(BETI)0.3 2H2 hydrate melt electrolytes (the same cells shown in Fig. 4a, b). The cell capacity was calculated based on the total weight of the positive and negative electrodes. 8 NATURE ENERGY

9 SUPPLEMENTARY INFRMATIN a LiCo 2 (0.2C) b LiCo 2 (0.5C) 3rd 2nd 1st 3rd 2nd 1st Capacity / Ah kg Cycle number Efficiency / % c LiCo 2 (10C) Capacity / Ah kg Cycle number Efficiency / % 3rd 2nd 1st Capacity / Ah kg Cycle number Efficiency / % Supplementary Figure 8 Charge-discharge voltage profiles at various C-rates for a 2.4 V LiCo2/Li4Ti512 full cell (LiCo2: 5~8 mg cm 2, Li4Ti512: 2.5~4.5 mg cm 2 ) with an alkaline hydrate melt electrolyte, Li1+x(TFSI)0.7(BETI)0.3(H)x (2 x)h2 (x=35), and a quartz crystal separator. Charging and discharging were performed at (a) 0.2C (27.0 ma g 1 for LiCo2), (b) 0.5C (67.5 ma g 1 for LiCo2), and (c) 10C (1.35 A g 1 for LiCo2) at 25 C. Al and Ti were used as the negative and positive current collectors, respectively. The cell capacity was calculated based on the total weight of the positive and negative active materials. The discharge capacity and coulombic efficiency of the cells upon cycling are shown in the insets. NATURE ENERGY 9

10 a LiCo 2 (0.2C) 25 o C b LiCo 2 (0.5C) 25 o C 2nd 10th 20th 50th 100th nd 10th 20th 50th 100th c LiCo 2 (10C) 25 o C 2nd 10th 20th 50th 100th 500th Supplementary Figure 9 Comparison of the charge-discharge voltage profiles of 2.4 V LiCo2/Li4Ti512 full cells (shown in Supplementary Fig. 8) after prolonged cycling at rates of (a) 0.2C, (b) 0.5C, and (c) 10C. 10 NATURE ENERGY

11 b Capacity / Ah kg 1 SUPPLEMENTARY INFRMATIN a LiCo 2 (0.2C) 30th 10th 2nd 22 mol kg 1 LiTFSI/H 2 1st mol kg 1 LiTFSI/H 2 hydrate melt Efficiency / % Cycle number Supplementary Figure 10 a, Charge-discharge voltage profiles of a 2.4 V LiCo2/Li4Ti512 full cell with a saturated 22 mol kg 1 LiTFSI aqueous electrolyte (not a hydrate melt). Charging and discharging were performed at 0.2C (27.0 ma g 1 for LiCo2) and 25 C. The cell was composed of 8 mg cm 2 LiCo2, 4 mg cm 2 Li4Ti512, and a quartz crystal separator. Al and Ti were used as the negative and positive current collectors, respectively. The cell capacity was calculated based on the total weight of the positive and negative active materials. b, Discharge capacity and coulombic efficiency upon cycling of the cells with the saturated LiTFSI aqueous electrolyte and the alkaline hydrate melt electrolyte (shown in Supplementary Fig. 8a) at 0.2C (27.0 ma g 1 for LiCo2) and 25 C. NATURE ENERGY 11

12 a LiNi 0.5 Mn (0.5C) st b LiNi 0.5 Mn (C) st Capacity / Ah kg Cycle number Efficiency / % Capacity / Ah kg Cycle number Efficiency / % Supplementary Figure 11 Charge-discharge voltage profiles of 3.1 V LiNi0.5Mn1.54/Li4Ti512 full cells (LiNi0.5Mn1.54: 6~7 mg cm 2, Li4Ti512: approximately 2 mg cm 2 ) with a Li(TFSI)0.7(BETI)0.3 2H2 hydrate melt electrolyte and a quartz crystal separator. Charging and discharging were performed at (a) 0.5C (73.3 ma g 1 for LiNi0.5Mn1.54) and (b) C (146.6 ma g 1 for LiNi0.5Mn1.54) at 25 C. Al and Ti were used as the negative and positive current collectors, respectively. The cell capacity was calculated based on the total weight of the positive and negative active materials. The discharge capacity and coulombic efficiency upon cycling are shown in the insets. 12 NATURE ENERGY

13 SUPPLEMENTARY INFRMATIN a LiCo 2 (10C) 25 o C b 1st 2.35 V 130 Wh kg Ah kg LiNi 25 o 0.5 Mn (6.8C) C 1st 7 V 90.6 Wh kg Ah kg Supplementary Figure 12 Calculation of the actual energy densities from the initial discharge voltage profiles of (a) a LiCo2/Li4Ti512 cell with an alkaline hydrate melt electrolyte (Li1+x(TFSI)0.7(BETI)0.3(H)x (2 x)h2 (x=35)), and (b) a LiNi0.5Mn1.54/Li4Ti512 cell with a hydrate melt electrolyte (Li(TFSI)0.7(BETI)0.3 2H2). Charging and discharging were performed at (a) 10C (1.35 A g 1 for LiCo2) and (b) 6.8C (0 A g 1 for LiNi0.5Mn1.54) at 25 C. The LiCo2/Li4Ti512 cell was composed of 1.7 mg cm 2 LiCo2, mg cm 2 Li4Ti512, and a glass fiber separator. The LiNi0.5Mn1.54/Li4Ti512 cell was composed of mg cm 2 LiNi0.5Mn1.54, mg cm 2 Li4Ti512, and a quartz crystal separator. Al and Ti were used as the negative and positive current collectors, respectively. The cell capacity was calculated based on the total weight of the positive and negative active materials. NATURE ENERGY 13

14 F1s PVdF salt damaged salt CF CF and/or passivation 3 2 LiF pristine F1s TFSI CF 3 damaged TFSI LiF LiTFSI powder F1s BETI damaged BETI CF 3 CF 2 LiF LiBETI powder S2p salt -S 2 - passivation film sulfide S x pristine S2p TFSI -S 2 - damaged TFSI LiTFSI powder S2p BETI -S 2 - damaged BETI LiBETI powder Li1s Ti3s LT (Ti3s) LiF Li salt LT (Li1s) pristine Li1s Ti3s damaged LiTFSI LiF LiTFSI LiTFSI powder Li1s Ti3s damaged LiBETI LiF LiBETI LiBETI powder NATURE ENERGY

15 SUPPLEMENTARY INFRMATIN N1s salt S-N-S reduced anion? pristine N1s TFSI S-N-S LiTFSI powder N1s BETI S-N-S LiBETI powder C1s binder salt AB or contamination CF 3 CF 2 sp pristine C1s LiTFSI CF 3 damaged anion or contamination sp 3 LiTFSI powder LiBETI damaged anion CF 3 CF 2 or contamination LiBETI powder C1s sp Supplementary Figure 13 XPS spectra (F1s, S2p, Li1s, Ti3s, N1s, and C1s) of pristine and singlecycled Li4Ti512 electrodes in LiCo2/Li4Ti512 full cells with hydrate melt electrolytes. To clarify the damage to the residual Li salts in the electrode caused by X-ray radiation or Ar + sputtering, the XPS spectra of LiTFSI and LiBETI powders deposited on carbon tape were obtained. The depth profiles of the cycled electrode and the Li salt powders were obtained via Ar + sputtering at 1 kv for 60~300. NATURE ENERGY 15

16 Supplementary Table 1 Electrode capacities and potentials used to calculate the theoretical energy densities of the batteries depicted in Fig. 4c. Electrode (active Li) Reversible capacity [mah g 1 ] Average potential [V vs. Li + /Li (standard) a ] LiCo2 ( 0.5Li) LiNi0.5Mn1.54 ( Li) LiNi0.5Mn1.54 ( 0.5Li) LiMn24 ( 0.7Li) Graphite, C (+1/6Li) Li4Ti512 (+Li) a The term standard denotes conditions such that both the Li + and solvent activities are 1 and the temperature is K. 16 NATURE ENERGY

17 SUPPLEMENTARY INFRMATIN Supplementary Note 1: Relationship between potential shift and Li + activity According to the Nernst equation, the equilibrium potential (E) of the Li + intercalation/deintercalation reaction depends on the activity of Li + (ali) in the electrolyte solution: EE = EE o + RRRR FF ln aa Li (1) where E o, R, T, and F denote the standard potential, gas constant, absolute temperature, and Faraday constant, respectively. With the equilibrium potential in a 1.2 mol kg 1 LiTFSI aqueous solution as the standard (Estd = E o + (RT/F) ln ali,std, where ali,std is the Li + activity in a 1.2 mol kg 1 LiTFSI aqueous solution), the potential shift ( E = E Estd) can be expressed as a function of the Li + activity as follows. EE = RRRR FF ln aa Li RRRR FF ln aa Li,std = RRRR FF ln aa Li aa Li,std (2) In this work, the measurement was performed at 25 C, and thus, the equation can be transformed in the following way. aa Li EE = 59 log (V) aa Li,std (3) According to this equation, the experimentally observed E (>0, as shown in Fig 5b) corresponds to an increased Li + activity compared with the 1.2 mol kg 1 LiTFSI/H2 standard. The Li(TFSI)0.7(BETI)0.3 2H2 hydrate melt exhibits the highest ali, which is times higher than that in the standard electrolyte. Supposing that ali is expressed as a product of the activity coefficient and the molality, the activity coefficient of Li + in the hydrate melt can be evaluated to be times higher than that in the standard electrolyte. Such a strongly increased ionic activity coefficient of far greater than 1 is widely observed in various aqueous solutions at high concentrations, 4 and this phenomenon cannot be explained by the conventional Debye-Hückel theory (based on dilute solutions). Stokes and Robinson rationalized the increased ionic activity coefficient by considering NATURE ENERGY 17

18 the effect of hydration. 4,5 They found a trend of increase in the ionic activity coefficient with an increasing extent of hydration (e.g., a decreasing radius of the solute cations). 4 According to these authors, hydration increases the ionic activity coefficient by eliminating the role of the water as a solvent and instead causing the water to act as part of the solute, and thus, it is necessary to distinguish the bound water molecules in the hydration shells from the free water molecules outside the hydration shells. Based on this discussion, they introduced a solvent term ( (n/ )log aw) into the Debye-Hückel equation: 5 log γγ ± = AA II 1 + BBBB II nn νν log aa ww log[1 18(nn νν)mm] (4) where ± is the mean ionic activity coefficient, A and B are constants, r is the mean distance of closest approach of the ions, I is the ionic strength, n is the hydration number, is the number of ions formed by one solute, aw is the water activity, and m is the molality. Notably, a decrease in the water activity causes the mean ionic activity coefficient to increase. Although this equation might not be valid for the present hydrate melt, with its extremely low water content, this approach can qualitatively explain the high Li + activity coefficient, which deviates from the conventional Debye-Hückel theory. 18 NATURE ENERGY

19 SUPPLEMENTARY INFRMATIN Supplementary Note 2: Verification of high Li + activity based on water activities According to the Gibbs-Duhem equation at constant pressure and temperature, the changes in the chemical potentials of the Li salts ( s) and water ( w) in the electrolyte are related to each other: xx s dμ s + xx w dμ w = 0 (5) where xs and xw denote the mole fractions of Li salts and water in the solution (xs+xw=1). Thus, using the activities of the Li salts (as) and water (aw), the following equation can be derived: xx s d ln aa s + xx w d ln aa w = 0 d ln aa s = xx w xx s d ln aa w xx w d ln aa s = d ln aa 1 xx w w (6) (7) (8) By integrating the above equation from the values corresponding to a standard 1.2 mol kg 1 LiTFSI/H2 solution (as,std and aw,std) to those corresponding to the hydrate melt state (as,hm and aw,hm), the following equation is derived: ln aa s,hm = aa s,std aa w,hm aa w,std xx w daa (1 xx w )aa w w (9) Hence, we can obtain the ratio of the Li salt activities (as,hm/as,std) from the activities and molar fractions of water in the electrolyte solutions. Here, using the available water activity values of aw,std=0.965 (at xw,std= in 1.2 mol kg 1 LiTFSI/H2) and aw,hm=0.118 (at xw,hm= in the hydrate melt), the ratio of the Li salt activities (as,hm/as,std) was roughly estimated to be Note that this is not the Li + activity but rather the Li salt activity, which is the product of the Li + activity and the anion activity. At present, there is no way to estimate the individual Li + and anion activities from the Li salt activity. However, the average combined Li + and anion activity can be used to obtain a rough estimate of the individual activities. The average activity in the hydrate melt was found to be ( ) 1/2 = times higher than that in the standard electrolyte, which NATURE ENERGY 19

20 corresponds to a 0.36 V upward shift in potential. This value, though roughly estimated, is quite close to the experimental observation. 20 NATURE ENERGY

21 SUPPLEMENTARY INFRMATIN Supplementary References 1. Xu, W., Cooper, E. I. & Angell, C. A. Ionic liquids: ion mobilities, glass temperatures, and fragilities. J. Phys. Chem. B 107, (2003). 2. Bressel, R. D. & Angell, C. A. Fluidity and conductance in aqueous electrolyte solutions. An approach from the glassy state and high-concentration limit. I. Ca(N3)2 solutions. J. Phys. Chem. 1086, (1979). 3. McLin, M. & Angell, C. A. Contrasting conductance/viscosity relations in liquid states of vitreous and polymer solid electrolytes. J. Phys. Chem. 92, (1988). 4. Stokes, R. H. A thermodynamic study of bivalent metal halides in aqueous solution. Part XVII-Revision of data for all 2:1 and 1:2 electrolytes at 25, and discussion of results. Trans. Faraday Soc. 44, (1947). 5. Stokes, R. H. & Robinson, R. A. Ionic hydration and activity in electrolyte solutions. J. Am. Chem. Soc. 70, (1948). NATURE ENERGY 21

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