Li+ Solvation and Ionic Transport in Lithium Solvate Ionic Liquids

Article pubs.acs.org/JPCC

Li+ Solvation and Ionic Transport in Lithium Solvate Ionic Liquids Diluted by Molecular Solvents Kazuhide Ueno,† Junichi Murai,‡ Kohei Ikeda,‡ Seiji Tsuzuki,§ Mizuho Tsuchiya,‡ Ryoichi Tatara,‡ Toshihiko Mandai,‡ Yasuhiro Umebayashi,∥ Kaoru Dokko,‡ and Masayoshi Watanabe*,‡ †

Graduate School of Medicine, Yamaguchi University, 2-16-1 Tokiwadai, Ube, 755-8611, Japan Department of Chemistry and Biotechnology, Yokohama National University,79-5 Tokiwadai, Hodogaya-ku, Yokohama 240-8501, Japan § Research Center for Computational Design of Advanced Functional Materials (CD-FMat), National Institute of Advanced Industrial Science and Technology (AIST), 1-1-1 Umezono, Tsukuba, Ibaraki 305-8568, Japan ∥ Graduate School of Science and Technology, Niigata University, 8050 Ikarashi, 2-no-cho, Nishi-ku, Niigata City, 950-2181, Japan ‡

S Supporting Information *

ABSTRACT: An equimolar mixture of lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA]) and either triglyme (G3) or tetraglyme (G4) yielded stable molten complexes: [Li(G3)][TFSA] and [Li(G4)][TFSA]. These are known as solvate ionic liquids (SILs). Glyme-based SILs have thermal and electrochemical properties favorable for use as lithium-conducting electrolytes in lithium batteries. However, their intrinsically high viscosities and low ionic conductivities prevent practical application. Therefore, we diluted SILs with molecular solvents in order to enhance their ionic conductivities. To determine the stabilities of the complex cations in diluted SILs, their conductivity and viscosity, the self-diffusion coefficients, and Raman spectra were measured. [Li(G3)]+ and [Li(G4)]+ were stable in nonpolar solvents, that is, toluene, diethyl carbonate, and a hydrofluoroether (HFE); however, ligand exchange took place between glyme and solvent when polar solvents, that is, water and propylene carbonate, were used. In acetonitrile (AN) mixed solvent complex cations [Li(G3)(AN)]+ and [Li(G4)(AN)]+ were formed. [Li(G4)][TFSA] was more conductive than [Li(G3)][TFSA] when diluted with nonpolar solvents due to the greater ionic dissociativity in [Li(G4)][TFSA] mixtures. In view of the stability of the Li−glyme complex cations, the enhanced ionic conductivities, and the intrinsic electrochemical stabilities of the diluting solvents, [Li(G4)][TFSA] diluted by toluene or HFE, can be a candidate for an alternative battery electrolyte.

1. INTRODUCTION Growing demand for large-scale energy storage solutions has stimulated the development of novel combinations of anodic and cathodic materials for advanced lithium-ion batteries and next-generation batteries.1−3 In addition, a number of studies have been conducted on the design of novel electrolytes that support the electrode reactions of these new battery materials.4 The current standard electrolyte used in commercialized lithium-ion batteries is LiPF6 dissolved in a mixed solvent of ethylene carbonate (EC) and linear carbonates such as diethyl carbonate (DEC) at a concentration of approximately 1 mol dm−3.5 These electrolytes are the first choice for testing new electrode materials. However, the standard electrolytes cannot withstand operation at high voltages with 5-V-class cathode materials such as LiNi0.5Mn1.5O4 spinel.6 Furthermore, the electrolyte undergoes side reactions with the reaction intermediates of the high energy density next-generation batteries (Li−sulfur and Li−air batteries).7,8 To meet the © 2015 American Chemical Society

requirements of future batteries, various alternative electrolytes other than those based on carbonates have been proposed. For example, fluorinated solvents,9,10 phosphorus-based solvents,11,12 sulfones,13,14 ionic liquids,15−18 and highly concentrated (or superconcentrated) electrolytes.19−23 Among the new classes of electrolytes, there is particular interest in highly concentrated electrolytes, where the salt concentration typically exceeds 3 mol dm−3 or approaches the salt saturation limit.24 Interest in these materials has arisen because recent findings have shown anomalous behavior concerning their bulk and interfacial properties; and these properties are favorable for electrolytes in advanced lithium-ion and next-generation batteries; for example, (i) higher thermal stability,25 (ii) low volatility, (iii) higher lithium transference Special Issue: Kohei Uosaki Festschrift Received: November 29, 2015 Published: December 15, 2015 15792

DOI: 10.1021/acs.jpcc.5b11642 J. Phys. Chem. C 2016, 120, 15792−15802

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The Journal of Physical Chemistry C number (>0.5),25 (iv) enhanced stability against electrochemical oxidation and reduction,25−27 (v) fast electrode reaction,28,29 (vi) suppression of aluminum current collector corrosion,30,31 and (vii) suppression of lithium polysulfide dissolution.19,20,32 In highly concentrated electrolytes, the number of solvent molecules and ions are comparable. Therefore, the solvation of lithium ions in the highly concentrated electrolyte differs from that in a corresponding dilute solution of the same components. This unique Li+ solvation is believed to be related to the observed unusual bulk and interfacial phenomena. Equimolar (1:1) mixtures of lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA]) and either triglyme (G3, CH3-O-(CH2-CH2-O)3-CH3) or tetraglyme (G4, CH3-O-(CH2-CH2-O)4-CH3) yield low-melting complexes; these are abbreviated as [Li(G3)][TFSA] or [Li(G4)][TFSA], respectively, and the bulk materials show ionic-liquid-like behavior25 and anomalous interfacial electrochemical properties.33−35 Therefore, the molten complexes are similar to the highly concentrated electrolytes that are currently being investigated for use in advanced lithium-ion and nextgeneration batteries.36,37 Moreover, recent experiments on structural and transport properties38,39 and computational investigations40,41 revealed that the complex cations, [Li(G3)]+ and [Li(G4)]+, are long-lived in the molten state, so they behave as independent cations in a similar way to the organic cation in a typical ionic liquids. Consequently, [Li(G3)][TFSA] and [Li(G4)][TFSA] are also solvate ionic liquids (SILs).42,43 Further investigations have shown that formation of stable complexes is responsible for their higher thermal and enhanced electrochemical stabilities25 and anomalous electrochemical reactions.33,34 High viscosities and low ionic conductivities are common drawbacks for highly concentrated electrolytes including [Li(G3)][TFSA] and [Li(G4)][TFSA], and these drawbacks must be addressed for applications in batteries. Addition of lowviscosity (but often nonpolar) solvents to the electrolyte in polar solvents is a compromise, improving the conductivity, as the standard electrolyte has been optimized in that way. Indeed, [Li(G3)][TFSA] and [Li(G4)][TFSA] are both soluble in low polar solvents such as hydrofluoroethers (HFE),37 as well as in polar solvents. In other words, these molten complexes can serve as the liquid lithium salt soluble in low dielectric solvents in which standard lithium salts are insoluble. Here, we study lithium ion solvation and the ionic transport properties for [Li(G3)][TFSA] and [Li(G4)][TFSA] diluted by various molecular solvents. The primary purpose of this work is to understand the stabilities of the complex cations in the diluted electrolytes because robust complex cations, such as [Li(G3)]+ and [Li(G4)]+, play a key role in the electrochemical properties of the electrolyte. The stability of the complex cations was evaluated by Raman spectra and self-diffusion coefficients of the glyme molecules and lithium ions. The concentration dependence of the conductivity and viscosity are also reported.

(PC) and diethyl carbonate (DEC; from Kishida Chemical) were used as received. A hydrofluoroether, 1,1,2,2-tetrafluoroethyl,2,2,3,3-tetrafluoropropyl ether, which has a chemical structure of HCF2-CF2-O-CH2-CF2-CF2H (hereafter, abbreviated as HFE), was purchased from Kanto Kagaku and dried with molecular sieves. The electrolyte solutions were prepared by mixing appropriate amounts of the lithium salt, glyme, and the additional solvent in an Ar-filled glovebox (VAC, [H2O] < 1 ppm) except for the aqueous solutions. Ultrapure water prepared by a Milli-Q integral water purification system (Millipore) was used to prepare the aqueous solution of the solvate ILs. Measurements. The densities and viscosities of the mixtures were measured using an SVM3000 viscometer (Anton Paar), and the molar concentration of the Li salt was calculated from the density of the solution. Raman spectra of the electrolyte solutions were recorded using a 532 nm laser spectrometer (NRS-4100, JASCO) at ambient temperature (22 ± 2 °C). The resolution of the apparatus was about 2 cm−1. Raman bands were analyzed using the JASCO spectra manager program. When necessary, the intensity was normalized with respect to the Raman bands between 1400 and 1550 cm−1, ascribable to the glyme vibrations. These bands are practically not affected in the intensity and the position upon coordination to the Li+ ion.44 Ionic conductivity was measured by the complex impedance method, in the frequency range of 500 kHz to 1 Hz at an amplitude of 10 mV (VMP3, Bio-Logic). Two platinized platinum electrodes (CG-511B, TOA Electronics, cell constant ≈ 1 cm−1) were dipped in the diluted SILs, and the sample cell was thermally equilibrated at 30 °C for a minimum of 30 min using a thermostat chamber (Espec, SU-222). Pulsed-field-gradient spin echo NMR measurements were carried out to determine the self-diffusion coefficients of the components of the electrolyte solutions. A JEOL ECX-400 NMR spectrometer with a 9.4 T narrow-bore superconducting magnet and a pulsed-field gradient probe was used for the measurements. 1H, 7Li, and 19F NMR spectra were recorded for the solvents, Li+, and [TFSA]−, respectively. The self-diffusion coefficients were measured via the use of a modified Hahn spin echo-based or stimulated echo-based sequence incorporating a pulsed field gradient (PFG) in each τ period. The free diffusion echo signal attenuation, E, is related to the experimental parameters by the Stejskal equation with a sinusoidal PFG: ln(E) = ln(S /Sδ= 0) =

−γ 2g 2Dδ 2(4Δ − δ) π2

where S is the spin echo signal intensity, δ is the duration of the field gradient with magnitude g, γ is the gyromagnetic ratio, and Δis the interval between the two gradient pulses. The value of Δ (50 ms) and δ (1−10 ms) were set at constant values, whereas g (0−13 T/m) was varied for the diffusion measurements. All the results were well-described by the Stejskal equation, and the standard deviations of the diffusion data were less than 5%. The PFG-NMR samples were inserted into a NMR microtube (BMS-005J, Shigemi) to a height of 3− 5 mm to exclude convection effects. All measurements were performed at 30 °C. Absorption spectra of a copper complex ([Cu(acac)(tmen)][BPh4], where acac is acetylacetone and tmen is tetramethylethylenediamine), were measured using a UV−vis spectrom-

2. EXPERIMENTAL SECTION Materials. Purified glymes (water content < 50 ppm), triglyme (G3) and tetraglyme (G4), were kindly supplied by Nippon Nyukazai Co., Ltd., and Li[TFSA] (battery-grade, water content < 50 ppm) was obtained from Solvay Japan. Anhydrous molecular solvents, acetonitrile, acetone, and toluene (from Wako), and battery-grade propylene carbonate 15793


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Gutmann donor numbers (DN).5,37,45,49,50 Both SILs were found to be compatible with water, PC, AN, and acetone. Even less-polar solvents, such as HFE and DEC, were suitable over the whole range of concentrations studied. Although mixing nonpolar toluene and the SILs resulted in a biphasic solution, the salt concentration in the SIL-rich phase was sufficient to prepare the battery electrolyte: [Li(G3)][TFSA] and [Li(G4)][TFSA] can be diluted by toluene down to 1.30 and 1.26 mol dm−3, respectively, at room temperature. However, these SILs were barely soluble in hexane. It is likely that the solubility of [Li(G3)][TFSA] and [Li(G4)][TFSA] correlates well with the dielectric constant of the molecular solvents and they can be diluted to a sufficient salt concentration (around 1 mol dm−3) for conductivity by using a solvent that has a dielectric constant greater than ∼2.4. Stability of Li−Glyme Complex Cation. The ratio of selfdiffusion coefficients for lithium and glyme molecules (Dglyme/ DLi) allows the stabilities of the complex cations to be understood. PFG-NMR using 1H, 19F, and 7Li nuclei enables us to measure the diffusivity of each component, and D is obtained on a time scale of 10−2 to 10−3 s. In a previous study, we found that the Dglyme/DLi ratio is unity for [Li(G3)][TFSA] and [Li(G4)][TFSA], suggesting that Li+ diffuses together with glyme molecules in form of long-lived complex cations, that is, [Li(G3)]+ and [Li(G4)]+.51 In contrast, Li[TFSA]−glyme solutions containing excess glyme51 and the molten complexes of [Li(G3)]X and [Li(G4)]X, having associative anions X, have Dglyme/DLi ratios greater than one, suggesting the presence of uncoordinated glyme molecules.52 Figure 1 shows the concentration dependence of the Dglyme/ DLi ratio for the diluted SILs. First, evaluation of the Dglyme/DLi ratio did not show a discernible difference in the stabilities of [Li(G3)]+ and [Li(G4)]+; that is, the changes in the Dglyme/DLi ratios for the SILs were similar in the same solvent. The Dglyme/ DLi ratio was almost unity within experimental error for the solutions of [Li(G3)][TFSA] and [Li(G4)][TFSA] diluted by low polar solvents such as HFE, DEC, and toluene. In addition, in AN and acetone, which are both polar, the Dglyme/DLi ratios were also almost 1 at concentrations greater than 1 mol dm−3. The diffusivity ratios for the SIL solutions in PC remained close

eter (Shimadzu UV-1800) and used to estimate the Gutmann donor number (DN) of each solvent.45 Ab initio calculations were carried out in Gaussian 09.46 Basis sets were used as implemented in the Gaussian program. The geometries of the solvated Li+ ions were fully optimized at the HF/6-311G** level. The interaction energies (Eint) for the solvated Li+ ions were calculated at the MP2/6-311G** level using the optimized geometries. The basis set superposition error (BSSE)47 was corrected for using the counterpoise method.48 The stabilization energy for formation of a complex from an isolated species (Eform) was calculated as the sum of Eint and the deformation energy (Edef), which is the increase in energy associated with deformation of the solvent molecule geometries accompanied by complex formation.

3. RESULTS AND DISCUSSION Solubility of [Li(G3)][TFSA] and [Li(G4)][TFSA] in Molecular Solvents. The solubilities of the SILs, [Li(G3)][TFSA] and [Li(G4)][TFSA], were preliminarily tested in a variety of polar and nonpolar solvents. The studied solvents are listed in Table 1, along with their dielectric constants (ε) and Table 1. Dielectric Constants, Gutmann Donor Numbers, and Solubilities of the SILs in the Listed Solvents solvent

dielectric constant, ε (−)

donor number, DN (kcal mol−1)

solubility of the solvate ILs

G3 G4 water PC AN acetone HFE DEC toluene hexane

7.5a 7.7a 78b 65b 36b 21b 6.2c 2.8e 2.4b 1.9b

14a 17a 18b 15a 14b 17b 1.9d 10d 0.1f 0b

soluble soluble soluble soluble soluble soluble partially soluble insoluble

a Ref 49. bRef 45. cRef 37. dDN was estimated from the absorption spectra of a copper complex [Cu(acac)(tmen)][BPh4] in the solvent or a mixture of the solvent and 1,2-dichloromethane using the relationship described in ref 45. eRef 5. fRef 50.

Figure 1. Concentration dependence of the Dglyme/DLi ratio for (a) [Li(G3)][TFSA] and (b) [Li(G4)][TFSA] diluted by molecular solvents at 30 °C. 15794


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The Journal of Physical Chemistry C to Dglyme/DLi ∼ 1 at higher concentrations (cLi > 1.5 mol dm−3), but gradually increased with further dilution. This suggests that glymes coordinating to Li+ were replaced by PC, and the diffusion of the Li+ solvated by PC was slower than that of the free glyme. Hence, the Li-glyme complex cations were unstable in PC in this region. In the aqueous solutions of the SILs, the addition of water to the SILs led to a noticeable decrease in the Dglyme/DLi ratio; however, the Dglyme/DLi ratio was unity when an equimolar quantity of water was added to the SILs (cLi = 2.80 mol dm−3 for [Li(G3)][TSFA] and 2.62 mol dm−3 for [Li(G4)][TFSA]). In dilute aqueous solutions, the Dglyme/DLi ratio was less than 1, probably due to the smaller size of the hydrated Li+ ion than that of the free glymes. Therefore, marked ligand exchange occurs between water and the glymes, and the Li−glyme complex cations were unstable in the diluted aqueous solutions. The stabilities of the Li−glyme complex cations were further investigated using Raman spectroscopy. Spectra were measured for solvents with negligible spectral overlap between the peaks of the glymes and the solvents. For glymes, the Raman bands between 800 and 900 cm−1 have been assigned to a mixture of modes, CH2 rocking vibrations and C−O−C stretching vibrations.53 Pure G3 and G4 have very similar Raman spectra, featuring three bands at 809, 828, and 851 cm−1.54,55 For G3 and G4, a strong band emerged at around 870 cm−1 (the socalled breathing mode) upon their complexation with Li+.56 Therefore, the Raman spectrum in this region reflects the Li+ coordination structure of the glymes in the diluted solution. Figure 2 shows the Raman spectra of [Li(G3)][TFSA] diluted by water, AN, and toluene, with respect to the mixing ratio of the solvent and the SILs (cS/cLi). As shown in Figure 2a, for pure [Li(G3)][TFSA], an intense peak, arising from a breathing mode, is visible at 873 cm−1. However, on further dilution, the intensity of this band became weaker. Simultaneously, in the spectrum of [Li(G3)][TFSA] on addition of water, three bands at 809, 828, and 851 cm−1 due to uncoordinated G3 became more pronounced. This indicates that the complex cations are destabilized in the presence of water, and ligand exchange with water yields free G3. A reduction in the intensity of the peak at 873 cm−1 was also seen at a certain level for the solutions in AN; however, an increase in the peak intensity between 800 and 850 cm−1 for pure G3 was not observed (Figure 2b). The decrease in the intensity of the peak at 873 cm−1 was attributed to the formation of different complex cations consisting of AN and the glymes, and we will discuss this further later. For the toluene solutions (Figure 2c), the reduction of the peak intensity at 873 cm−1 was small; however, the change in the intensity of the peak between 800 to 850 cm−1 was not clear because of overlap with peaks of toluene. The Raman spectra of [Li(G4)][TFSA] diluted by water, AN, and toluene showed a similar trend, as shown in Figure S1. To investigate in detail the solvation of Li+ in dilute SILs, the Raman spectra were further analyzed. This spectral analysis offers quasi-quantitative estimation of the amount of uncoordinated glymes in the solutions. Herein, the Raman spectra were deconvoluted into four bands at around 810, 835, 850, and 873 cm−1 by a Gaussian−Lorentzian function, as described elsewhere.39 As previously reported,39 the integral intensity of the breathing mode at 873 cm−1 (I873) can be represented by contributions from both free (If873) and bound (Ib873) glymes as I873 = If873 + Ib873 = Jf873c f + Jb873cb

Figure 2. Raman spectra (900−800 cm−1) depending on the mixing ratio of the solvent and the SILs (cS/cLi) for (a) water/[Li(G3)][TFSA], (b) AN/[Li(G3)][TFSA], and (c) toluene/[Li(G3)][TFSA].

where cf, cb, Jf873, and Jb873 represent the concentration of free and bound G3 to Li+ ions, and the molar Raman scattering coefficients of the free and bound G3 for the band at ∼873 cm−1, respectively. From the total concentration of the glyme cG = cf + cb, where cG is the concentration of total glymes, we obtained the relationship between the fraction of free glyme (cf/cG) and I873 as eq 2. Jb873 I873 cf = − cG Jb873 − Jf873 cG(Jb873 − Jf873) (2) Because Jb873 and Jf873 are known for the solutions of Li salts of G3 and G4,39 cf/cG can be calculated from I873 for the diluted SILs if Jb873 and Jf873 are assumed to be constant in the diluted SILs. Figure 3 shows the percentage of free glyme (cf/cG), depending on the mixed ratio of the solvents and the SILs (cS/ cLi), calculated by eq 2. The SILs diluted by AN were excluded from this analysis because the change in I873 cannot be described as the simple two-state change between free and bound glymes (vide infra). From the spectral analyses, no significant differences between the diluted [Li(G3)][TFSA] and [Li(G4)][TFSA] SILs were

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complex cations obtained from the Dglyme/DLi ratio (Figure 1) for SILs diluted by water and toluene. In the aqueous solutions, the Li−glyme complex cations were unstable, even at cS/cLi = 1; in contrast, the cations were stable in toluene. Indeed, Raman bands due to toluene did not shift in the SIL mixtures (data not shown). Therefore, for the SILs, toluene is an “innocent” solvent, acting only as a thinner to increase their fluidity. In the SILs diluted by AN, I873 was lowered, as shown in Figure 2b, despite the Dglyme/DLi ratio being almost unity, as shown in Figure 1. At first glance, these results may seem contradictory concerning the stability of the complex cations. However, these findings raise the possibility that the coordinating structures of the Li−glyme complex cations vary in the presence of AN. The Raman band at 920 cm−1 in pure AN was assigned to C−C stretching vibrations and shifted to 930 cm−1 upon coordination with Li+.59 To study the effect of AN addition on the Li−glyme complex cations, the Raman band in this range was deconvoluted into two bands, located at 920 and 930 cm−1. Figure 4a,b shows Raman spectra of [Li(G3)][TFSA] or [Li(G4)][TFSA] diluted by AN between 900 and 960 cm−1. In both spectra, a weak shoulder peak at 930 cm−1 is discernible, and this is indicative of the coordination of AN with Li+ of the SILs. Here, the integral intensity of the band at 920 cm−1, I920, is related to the concentration of uncoordinated, free AN (cf_AN) in the solutions: I920 = J920cf_AN, where J920 is the molar Raman scattering coefficient of free AN molecules. Because cf_AN = cAN − ncLi, where n denotes the solvation number of AN to Li+ and cAN is the total concentration of AN, we obtain the following relationship:60

Figure 3. Estimated percentages of free glyme (cf/cG) in the SILs diluted by water and toluene, depending on the mixed ratio of the solvent and the SILs (cS/cLi).

seen (Figure 3). In aqueous solutions of both SILs, the fraction of uncoordinated, free glymes increased linearly in the range of 1 < cS/cLi < 4, indicating competitive solvation of Li+ by glyme and water in this region. However, water preferentially coordinates to Li+ at cS/cLi > 4. The isosolvation point57 at which the inner solvation shell of Li+ is equally populated by the glyme and water was found at cS/cLi = 2.81 for [Li(G3)][TFSA] and cS/cLi = 2.35 for [Li(G4)][TFSA], implying that [Li(G3)]+ was more stable than [Li(G4)]+ against the ligand exchange reaction with water. This may be attributed to stronger Li+−O interaction for [Li(G3)]+, as suggested by the results of ab initio calculations.58 In toluene solutions of both ILs, the fraction of free glyme was kept negligible and did not change, within experimental error, over the whole range of concentrations. The spectral changes that occur on changing the mixing ratio, cS/cLi (Figure 2), and the results of spectral analyses (Figure 3), generally agree with the stability evaluation of the

I920/c Li = J920 (cAN/c Li − n)

(3)

As shown in the insets of Figure 4a and b, the plots of I920/cLi versus cAN/cLi for the SILs diluted by AN yield a straight line, and, for [Li(G3)][TFSA] and [Li(G4)][TFSA], the average values of n were found to be 1.3 and 1.1, respectively. Considering the above results for SILs diluted with AN, Li+ was coordinated by both glyme and AN. In addition, one oxygen atom of the glyme or [TFSA] anion was unbound, that is, not coordinated to Li+ coordination, and one AN molecule

Figure 4. Raman spectra (960−910 cm−1) of the SILs diluted by AN depending on the mixing ratio of solvent and SIL (cS/cLi) for (a) [Li(G3)][TFSA] and (b) [Li(G4)][TFSA]. Insets show the plots of I920/cLi vs cAN/cLi to estimate the solvation number of AN in the diluted SILs. 15796


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The Journal of Physical Chemistry C replaced this. The decrease in the Raman band intensity I873, as seen in Figure 2b, may be caused by formation of a mixed solvent complex, for example, [Li(G3)(AN)]+ and [Li(G4)(AN)]+ (and the associated decrease in Jb873). The self-diffusion coefficients of AN were larger than those of Li+ and the glymes over the whole concentration range studied (Figure S2). This suggests that the mixed solvent complexes were not long-lived, and ligand exchange occurred frequently between the bound and free AN, the bound AN and unbound oxygen in a glyme molecule. Ligand exchange becomes more frequent if the diluting solvents form solvated Li+ complexes that are energetically more stable than the Li-glyme complex cation. Stabilization energies (Eform) for the formation of some solvated Li+ ([Li(solvent)n]+) cations from isolated Li+ and the solvent molecules was estimated using ab initio molecular orbital calculations to study Li+−solvent interactions depending on the different solvent species. In particular, this was used for carbonate solvents for which Raman spectral analysis failed. Figure 5 shows Eform of the solvated Li+ ions as a function of

the molecular solvents agrees with the experimental results of Dglyme/DLi that show that the Li−glyme cationic complexes become unstable in the presence of PC but remain stable in the presence of acetone, DEC, and toluene. In the case of AN, glyme and AN were bound to Li+ as mixed solvent complexes (as confirmed by the Raman study shown in Figures 2 and 4). Given the correlation between the stability of the complex cations and the solvent parameter of the diluents is considered, solvents with donor numbers much lower than those of the glymes (HFE, DEC, and toluene) behave as an innocent solvents to dilute the SILs; that is, no ligand exchange reactions occurred in these solvents. The DN values of water, PC, AN, and acetone were comparable to those of the glymes. The glyme molecules remained bound to Li+ at cLi values greater than 1 mol dm−3 in AN and acetone, which have intermediate dielectric constants, even with comparable DNs. This may be due to the high complex formation constant based on the chelate effect of the multidentate glymes.49 However, the highly polar solvents having high dielectric constant such as water and PC can easily replace glyme molecules coordinated to Li+ ions. Ionic Conductivity and Viscosity. The ionic conductivities of [Li(G3)][TFSA] and [Li(G4)][TFSA] are 1.1 and 1.6 S cm−1, respectively, at 30 °C,51 and these values are 1 order of magnitude smaller than that of the standard electrolyte used in commercial lithium ion batteries.5 These solutions are highly viscous (169 and 81 mPas for [Li(G3)][TFSA] and [Li(G4)][TFSA], respectively, at 30 °C), whereas the ionic dissociativity (ionicity) was quite high (0.68 and 0.63 for [Li(G3)][TFSA] and [Li(G4)][TFSA], respectively). Therefore, high viscosity is the reason for their low ionic conductivities.51 Here we report that ionic conductivity can be enhanced by lowering the viscosity of the electrolyte by adding molecular solvents. Figures 6 and 7 show concentration dependence of ionic conductivity and viscosity, respectively, for [Li(G3)][TFSA] and [Li(G4)][TFSA] diluted by the molecular solvents. Maximum conductivity was observed at around 1 mol dm−3, regardless of the identity of the diluting solvent for both [Li(G3)][TFSA] and [Li(G4)][TFSA]. This behavior is similar to the concentration dependence on ionic conductivity for typical electrolyte solutions, where ionic conductivity is governed by both ionic mobility and concentration. The lower ionic mobility, due to higher viscosity, is responsible for lower ionic conductivity at higher salt concentrations, while the lower ionic concentration results in lower ionic conductivities at lower salt concentrations. As seen in Figure 6, the ionic conductivities were enhanced by the addition of the molecular solvent in the following order: for [Li(G3)][TFSA], AN > water > PC > DEC > toluene > HFE; and, for [Li(G4)][TFSA], AN > acetone > water > PC ∼ toluene > DEC > HFE. The viscosity was lowered in the following order, as shown in Figure 7: for [Li(G3)][TFSA], AN < toluene ≤ DEC ∼ water < PC ∼ HFE; and, for [Li(G4)][TFSA], AN < acetone < toluene ≤ DEC ∼ water < PC ∼ HFE. The increase in ionic conductivity upon dilution was more pronounced in the SILs diluted by AN and acetone than by the other solvents. This is primarily due to lower viscosities of the SILs diluted by AN or acetone. The viscosities of the SILs diluted by water were comparable for those of the SILs diluted by DEC and toluene; nevertheless, the ionic conductivities were much higher for SILs diluted by water, probably due to the higher degree of dissociation of the ions in the aqueous electrolytes. This was further investigated and is discussed later.

Figure 5. Stabilization energies (Eform) for the formation of [Li(solvent)n]+ complexes from isolated Li+ and the chosen solvents were estimated using ab initio calculations at the MP2/6-311G**// HF/6-311G** level. The broken lines represent the reported Eform values for [Li(G3)]+ and [Li(G4)]+.58

number of coordinated solvent molecules. The optimized structures for each solvated Li+ ion are also shown in Figure S3. For toluene, [Li(toluene)4]+ did not form because of steric hindrance due to toluene molecules. The reported values of Eform for [Li(G3)]+ (−95.6 kcal mol−1) and [Li(G4)]+ (−107.7 kcal mol−1)58 were higher than those of [Li(solvent)4]+ (i.e., cS/ cLi = 4). This was also true, even when using the relatively nonpolar DEC. In DEC, the DG/DLi values, in the range of cLi > 1 mol dm−3 corresponding to cS/cLi < 4, do not suggest ligand exchange. The inconsistency between the experiments and the calculations (for AN, acetone, and DEC) may arise from the fact that the calculations did not take into account the chelate effect (less pronounced entropic penalty) for G3 and G4; thus, the Li+−glyme interaction free energy may be underestimated in the calculation. Nonetheless, we can compare the Li+− solvent interactions in [Li(solvent)n]+ complexes consisting of monodentate solvents. We found that Eform of solvated Li+ decreased in the following order: PC < AN < acetone < DEC < toluene. Furthermore, the ordering of formation energies for 15797


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Figure 6. Concentration dependence of ionic conductivity for (a) [Li(G3)][TFSA] and (b) [Li(G4)][TFSA] diluted by the molecular solvents at 30 °C. The solid lines are guides for the eye.

Figure 7. Concentration dependence of the viscosity in logarithmic scale for (a) [Li(G3)][TFSA] and (b) [Li(G4)][TFSA] diluted by the molecular solvents at 30 °C. The solid lines are guides for the eye.

Figure 8. Walden plots for (a) [Li(G3)][TFSA] and (b) [Li(G4)][TFSA] diluted by the molecular solvents at 30 °C.

When the ionic conductivity of the diluted [Li(G3)][TFSA] and [Li(G4)][TFSA] are compared, a noteworthy difference is

that the increase in the conductivity upon dilution by nonpolar toluene or HFE was more pronounced for [Li(G4)][TFSA]. 15798


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Figure 9. Concentration dependence of the DTFSA/DLi ratio for (a) [Li(G3)][TFSA] and (b) [Li(G4)][TFSA] diluted by molecular solvents at 30 °C.

solvents such as toluene and HFE, and the deviation was clearly smaller for [Li(G4)][TFSA], indicating that the ionic association in this system was less pronounced. Previous studies have shown that, in [Li(G3)][TFSA], an oxygen atom from the [TFSA] anion coordinates to Li+ in addition to four oxygen atoms from G3, forming a contact-ion pair (CIP).39,40,56 In contrast, in [Li(G4)][TFSA], free [TFSA] was dominant, and the ions were present as solvent separated ion pairs (SSIP), because the five oxygen atoms of G4 can coordinate to Li+, satisfying the most probable coordination number of Li+ in electrolyte solutions, that is, 4 or 5.39,40,56 Hence, distinct Li+ solvation modes between [Li(G3)][TFSA] and [Li(G4)][TFSA] are relevant concerning differences in the ionization states of the diluted SILs. The Coulombic attraction between the cations and anions are enhanced with decreasing dielectric constants of the solutions, leading to more significant deviation from the ideal line on dilution, as shown in Figure 8. It is plausible that the dissociation of the CIPs in [Li(G3)][TFSA] was affected by the reduction in solution permittivity due to the close proximity of the cation and the anion. Therefore, [Li(G4)][TFSA] present as SSIPs showed a greater degree of dissociation in comparison with [Li(G3)][TFSA] when diluted by nonpolar solvents. Figure 9 shows the ratio of self-diffusion coefficients of [TFSA] anions and Li+, DTFSA/DLi, for [Li(G3)][TFSA] and [Li(G4)][TFSA] diluted by the molecular solvents. The ratio DTFSA/DLi can also indicate the different ionization states that occur on dilution. In Figure 9a, DTFSA/DLi values for [Li(G3)][TFSA] diluted by HFE, DEC, and toluene are shown to approach unity on dilution, indicating that Li+ and [TFSA] diffuse together in form of an ion pair in nonpolar solvents. In contrast, the DTFSA/DLi values for [Li(G3)][TFSA] diluted by PC and AN increased, while DTFSA/DLi of [Li(G3)][TFSA] diluted by water decreased on addition of molecular solvents. These results suggest that Li+ and [TFSA] can diffuse separately in the polar solvents, reflecting their ionic dissociation. Li+ and [TFSA] have similar self-diffusion coefficients in neat [Li(G4)][TFSA] (DTFSA/DLi ∼ 0.97). However, this does not suggest significant ion pairing, as indicated by the high ionic dissociativity (ionicity: 0.63) in neat [Li(G4)][TFSA], but, instead, suggests comparable hydrodynamic radii of these ions.51 For this reason, the change in

The conductivity of [Li(G4)][TFSA] diluted by toluene (7.6 mS cm−1 at cLi = 1.28 mol dm−3) is similar to typical, nonaqueous electrolytes despite the difference in viscosities being small between diluted [Li(G3)][TFSA] and [Li(G4)][TFSA]. This may be due to a greater degree of dissociation for [Li(G4)][TFSA] in the presence of toluene (vide infra). State of Ionization. An empirical Walden rule states that the product of the molar conductivity (Λ) and the viscosity (η) is constant in dilute electrolyte solutions (Λη = constant).61 The Walden plot approach, namely, plots of log Λ versus log η−1, is a simple and versatile method that is applicable to many types of electrolytes if molar conductivities and viscosities are available. Indeed, it has been applied to evaluate the apparent degree of ionization (or ionicity) in ionic liquids62,63 and other ionic conductive materials, including nonaqueous electrolytes.64 To characterize the state of ionization, experimental data are compared to a straight reference line, calibrated by measurements from a 1 M KCl aqueous solution;62 in this situation, the ions are assumed to be fully dissociated and to have equal mobility. Thus, the apparent degree of dissociation can be estimated by any departure from the ideal behavior in the Walden plot.65 The Walden plots for the diluted SILs are shown in Figure 8. All the data lie below the ideal reference line. Further, the deviation from the ideal line becomes apparent with dilution when low polar solvents are used, illustrating that the incremental increase in conductivity is not as high as that expected from the reduced viscosity caused by addition of the molecular solvents. These facts imply that ionic association takes place in all diluted SILs. Deviation from the reference line became greater in the following order: for [Li(G3)][TFSA], water < PC < AN < DEC < HFE ∼ toluene; and, for [Li(G4)][TFSA], water < PC < AN ∼ acetone < DEC ∼ HFE ∼ toluene. The order of the solvents correlate well with the dielectric constants for each solvent, and the ionic dissociation of the SILs was lower in solvents with lower dielectric constants; therefore, this accounts for higher conductivity of the SILs diluted by water than by DEC and toluene despite their showing comparable viscosity values. Furthermore, significant differences in the deviations from the reference line can be seen when compared with [Li(G3)][TFSA] and [Li(G4)][TFSA] diluted by nonpolar 15799


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The Journal of Physical Chemistry C DTFSA/DLi for [Li(G4)][TFSA] diluted by nonpolar solvents was not apparent, remaining almost unity over the whole range of measured concentrations (Figure 9b). At lower cLi, the diluted [Li(G4)][TFSA] showed similar behavior concerning DTFSA/DLi to those of diluted [Li(G3)][TFSA]. [Li(G4)][TFSA] also formed an ion-pair in the nonpolar solvents although Li[(G4)][TFSA] was likely more dissociative than [Li(G3)][TFSA].

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AUTHOR INFORMATION

Corresponding Author

*Tel./Fax: +81-45-339-3955. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

4. CONCLUSIONS In this study, we investigated the solvation of Li+ in SILs diluted by molecular solvents using diffusivity measurements and Raman spectroscopy. From the diffusion experiments and Raman spectral analysis, we concluded that competitive solvation between the glyme and the diluents occurred in the SILs diluted by highly polar solvents such as water and PC; thus, the glyme−Li complex cations were unstable in the presence of these solvents. However, in less polar solvents, the glyme molecules preferentially coordinated to Li+, for example, toluene, DEC, and HFE, even at dilutions of less than 1 mol dm−3. In intermediate AN, the glyme remained bound to Li+, but AN also participated in the solvation of Li+ in the form of mixed solvent complex cations such as [Li(G3)(AN)]+ and [Li(G4)(AN)]+. Consequently, the glyme−Li complex cations were still stable in moderately polar solvents, such as AN and acetone, that have a DN comparable to those of the glyme ligands and an intermediate dielectric constant (in the range of cLi > 1 mol dm−3 at least) as well as in less polar solvents having lower DN values. Analysis with Walden plots illustrated that the ionic dissociativity (ionicity) is strongly dominated by the dielectric constant of the solvents. The increase in ionic conductivity upon dilution was greater for [Li(G4)][TFSA] when diluted by nonpolar solvents such as HFE and toluene. This was due to higher dissociativity of [Li(G4)][TFSA] in the nonpolar solvents. Among the diluting solvents studied here, water, AN, acetone, and DEC are not stable in direct contact with Li metal and are not stable at the electrochemical potentials of typical negative electrodes used in lithium-ion batteries (e.g., graphite). Furthermore, PC caused an unfavorable ligand exchange reaction with the Li−glyme complex cation. However, SILs diluted by toluene or HFE are applicable for use as battery electrolytes. In particular, diluted [Li(G4)][TFSA] is the most suitable in terms of its higher ionic conductivity. The excellent electrolyte properties of the SILs diluted by nonpolar HFE have been reported for Li−sulfur batteries: the redox shuttle mechanism was greatly suppressed due to the low solubility of polysulfide into the SILs diluted by HFE, leading to high Coulombic efficiencies and a longer cycle life of the Li−sulfur cell.34,37 To demonstrate the applicability of the SIL-based, nonaqueous electrolyte in lithium-ion batteries, the fundamental electrochemical properties and the results on the battery test using the SILs diluted by other solvents will be reported elsewhere.

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optimized structure for [Li(solvent)n]+ calculated by ab initio calculation (PDF).

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ACKNOWLEDGMENTS This study was supported by the Advanced Low Carbon Technology Research and Development Program (ALCA) of the Japan Science and Technology Agency (JST). This study was supported in part by the JSPS KAKENHI (Nos. 15H03874 and 15K13815 for K.D. and No. 15H06439 for K.U.) from the Japan Society for the Promotion of Science (JSPS).

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REFERENCES

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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.jpcc.5b11642. Raman spectra (900−800 cm−1) for the diluted [Li(G4)][TFSA], concentration dependence of the selfdiffusion coefficients for the SILs diluted by AN, and the 15800


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Li+ Solvation and Ionic Transport in Lithium Solvate Ionic Liquids

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