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Role of Solvent Bulkiness on Lithium-Ion Solvation in Fluorinated Alkyl Phosphate-Based Electrolytes: Structural Study for Designing Nonflammable Lithium-Ion Batteries Michiru Sogawa, Yanko M. Todorov, Daisuke Hirayama, Hideyuki Mimura, Nobuko Yoshimoto, Masayuki Morita, and Kenta Fujii J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.7b06910 • Publication Date (Web): 10 Aug 2017 Downloaded from http://pubs.acs.org on August 15, 2017

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The Journal of Physical Chemistry C is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

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The Journal of Physical Chemistry

Role of Solvent Bulkiness on Lithium-Ion Solvation in Fluorinated Alkyl Phosphate-Based Electrolytes: Structural Study for Designing Nonflammable Lithium-Ion Batteries Michiru Sogawa, a Yanko M. Todorov, a Daisuke Hirayama, b Hideyuki Mimura, b Nobuko Yoshimoto, a Masayuki Morita, a and Kenta Fujii a,* a

Graduate School of Sciences and Technology for Innovation, Yamaguchi University, 2-16-1

Tokiwadai, Ube, Yamaguchi 755-8611 b

TOSOH F-TECH, Inc., 4988 Kaisei-cho, Shunan, Yamaguchi 746-0006, Japan

1 Environment ACS Paragon Plus


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ABSTRACT

We report the characteristics of nonflammable tris-(2,2,2-trifluoroethyl) phosphate (TFEP) as a nonaqueous solvent for lithium (Li)-ion battery (LIB) electrolytes at a molecular level, particularly focusing on the solvent parameters, such as the electron pair-donating ability and molecular bulkiness of TFEP, and the solvation behavior around Li-ions. The binding energy for the interactions between Li-ion and several solvents using density functional theory (DFT) calculations indicated that TFEP can be categorized as a nonaqueous solvent with weak electron pair-donating ability, resulting in the prediction of a small Gutmann’s donor number (the DN value is approximately 12.9). The Walden plots based on ionic conductivity and viscosity suggested that Li-ions tend to exist as contact ion pairs (Li+⋅⋅⋅TFSA−) in the TFEP-based solutions, irrespective of the Li salt concentration cLi. Infrared and Raman spectroscopic studies demonstrated that Li-ions are solvated by two TFEP molecules and one TFSA anion to form the [Li(TFEP)2(TFSA)] complex as a major species in the solutions. This result was in good agreement with the DFT calculations for Li-ion solvation complexes: (1) the solvation energy for the [Li(TFEP)n] complexes linearly decreased with increasing n up to 2. (2) However, the steric repulsion among the solvated TFEP molecules occurs in complexes with n > 3 due to the molecular bulkiness of TFEP. We found that in the current system, [Li(TFEP)2(TFSA)] is a more energy-stable

complex

than

the

solvated


complexes

[Li(TFEP)n].

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INTRODUCTION

Nonaqueous electrolytes are crucial liquid materials for developing high-performance lithium (Li)-ion batteries (LIBs) with high operation voltage as well as solid electrode materials.2-6 To practically use nonaqueous electrolytes in LIBs, one major requirement is a well-designed electrode/electrolyte interphase.5, 7 It is well known that in conventional LIBs using a graphite negative electrode, carbonate-based electrolytes, including ethylene carbonate (EC), and linear carbonates, such as dimethyl carbonate (DMC), are required for working stable and reversible LIBs. The carbonate components, particularly EC, can form a stable passivation film [i.e., solid electrolyte interphase (SEI)] based on its reductive decompositions on the graphite anode during the first charging process to suppress further decomposed products.8-10 The resulting SEI plays a crucial role in reversible Li-ion insertion/desertion into/from the graphite. Recently, a new and attractive LIB electrolyte concept has been proposed: superconcentrated electrolytes.11-17 Superconcentrated electrolytes exhibit excellent LIB performance, which is comparable to the current commercialized systems (LiPF6 in EC + DMC) even in “carbonate solvent-free conditions.”

The

noticeable

characteristics

can

be

found

in

both

the

bulk

and

electrode/electrolyte interphase: (1) There are no free solvent molecules in the bulk phase. In other words, all solvent molecules coordinate to the Li-ions in solution. (2) The Li-ions can thus form complicated multiple ion pairs due to the lack of a solvent molecule in the system. (3) The anion species in the ion-pair complexes contribute to formation of an SEI on the graphite electrode to enhance the reductive stability. The other key for the practical application of nonaqueous electrolytes LIBs is “safety”. By using highly volatile nonaqueous solvents, the flammability of conventional electrolyte systems



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(also including super-concentrated systems) mentioned above represents a safety concern. To solve this problem, the development of nonflammable nonaqueous electrolytes has attracted considerable attention. One nonflammable LIB electrolyte approach involves the application of ionic liquids (ILs) having negligible volatilities.18-21 In particular, ILs containing the bis(fluorosulfonyl)amide (FSA) anion (e.g., 1-ethyl-3-methylimidazolium FSA22-23) allow for stable and rapid charging/discharging with a graphite electrode, resulting in promising rechargeable IL-based LIBs.24-26 However, there is an essential problem, i.e., dissolving Li salt in an IL results in a considerably viscous IL solution.27-28 This leads to a decrease in the ionic conductivity and thus leads to poor LIB performances. To avoid this problem while maintaining safety, we focused on the addition of nonflammable solvents into conventional carbonate-based electrolytes. In our previous research, we chose tris-(2,2,2-trifluoroethyl) phosphate (TFEP), which was first applied as flame retardant for LIB electrolytes by Xu et al.29-30. The resulting nonflammable electrolytes [i.e., 1.0 mol dm−3 LiPF6/EC + diethyl carbonate (DEC) containing TFEP] were characterized from both physicochemical and electrochemical viewpoints.31-32 In the EC:DEC:TFEP = 53:27:20 electrolyte solution, (1) the electrolyte exhibited a high ionic conductivity (7.5 mS cm−1 at 298 K), which exceeds the requirement for LIBs (1 mS cm−1). (2) This was mainly due to the low solution viscosity (7.8 mPa s), which is one order of magnitude lower than that for an IL solution containing Li salt. (3) The electrolyte showed good charge/discharge capacities and reversibilities at the graphite electrode. Furthermore, we found that the use of TFEP as a main solvent showed an excellent thermal stability even at elevated temperature and also even with/without a charged electrode. In the Li salt/TFEP solutions containing 1.0 mol dm−3 EC as a co-solvent, the electrolytes show no significant exothermic response up to 673 K, which was the same in the presence of 4.3 V-charged positive electrode


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(LiCoO2). This advantage is a new key for designing much safer LIBs. However, in the TFEPbased electrolyte system, Li-ion solvation at the molecular level, which plays an important role for understanding the electrode reaction mechanism in LIBs, has not yet been investigated to date. In general, the solvation of the metal ion controls the ion-pair formation, complexation, and electrode reaction in electrolyte solutions and strongly depends on the solvent species; i.e., solvent parameters such as permittivity, electron pair-donating/accepting ability, and molecular bulkiness (solvation steric effect).33-34 However, such solvent properties for the TFEP molecule, deeply related to an ion solvation behavior, have not yet been researched. We report on the Liion solvation in TFEP electrolyte solutions by experimental and theoretical techniques to understand and control the electrochemical reaction in LIBs using the TFEP-based electrolytes. First, we characterized the TFEP molecule in terms of Gutmann’s donor number (DN), which corresponds to an electron pair-donating ability. Then, the microscopic solvation structure of Liions was investigated via vibrational spectroscopy [infrared (IR) and Raman spectra] with the aid of density functional theory (DFT) calculations. Herein, we reported that the bulkiness of the TFEP molecule is an important factor for understanding Li-ion solvation in the TFEP system, i.e., the steric repulsion among the solvated TFEPs in the first solvation sphere of an Li-ion is predominant in the solvation, which is an unusual case among conventional LIB electrolytes.

EXPERIMENTAL SECTION

Materials. LiTFSA salt (Kanto Chemical, Battery grade) was vacuum dried at 373 K for 100 h. Solvent TFEP (Tosh F-Tech, Battery grade) was used without further purification. The



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physicochemical properties of TFEP (at 298 K) are as follows: the dielectric constant is 10.5,35 and the density and viscosity were experimentally determined in this work to be 1.5840 g cm−3 and 4.4057 mPa s, respectively. The water content was determined to be less than 50 ppm by Karl Fischer titration. The sample electrolyte solutions were prepared at the required molarities cLi in an Ar-filled glove box. The density values for the sample solutions examined in this research are listed in Table S1 in the Supporting Information. Ionic conductivity and viscosity. The ionic conductivity was measured by an AC impedance method in a frequency range from 100 kHz to 10 mHz using a frequency response analyzer (Solartron 1260). A cell with platinum electrodes was used for performing the measurements. The cell constant was determined by using an aqueous KCl solution (0.01, 0.1, and 1.0 mol dm−3). The viscosity was measured using a DV-I Prime viscometer (Brookfield). Both ionic conductivity and viscosity measurements were conducted at 298 K. ATR-IR and Raman spectroscopy. Attenuated total reflection (ATR)-infrared (ATR-IR) spectra were measured using an FT-Raman/IR spectrometer (FT/IR-6100; Jasco) equipped with a KBr beam splitter and a single-reflection ATR cell. The sample solution was placed on a diamond prism with an incident angle of 45°. The optical resolution was 4.0 cm−1, and the spectral data were accumulated 64. In ATR-IR spectroscopy, the penetration depth dp of the evanescent wave per reflection by a diamond prism was estimated by the following equation: dp = λ/[2πn1(sin2θ − n22/n12)1/2], where λ represents the wavelength, θ is the incident angle, n1 and n2 are the refractive indexes of the diamond prism (2.42) and sample solutions, respectively. The reflective indexes of the samples at 298 K were determined using an Abbe refractometer (NAR1T Solid, Atago), which are listed in the Supporting Information (Table S1). The observed absorbance A was thus corrected using the estimated dp value, i.e., A/dp. Raman spectroscopic


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measurements were conducted at 298 K using the FT-Raman/IR spectrometer equipped with an Nd:YAG laser. Raman spectra were obtained with an optical resolution of 4.0 cm−1 and were accumulated 2048 times to achieve a sufficiently high signal-to-noise ratio. The sample solution was filled in a quartz cell. To deconvolute both the Raman and IR spectra and extract single bands, we performed curvefitting analysis using the sum of pseudo-Voigt functions. The procedure is comprehensively described in our previous research.36-37 With regard to the IR spectra, the integrated intensity of the single band for free TFEP was represented as If = εfcfdp, where εf and cf are the molar extinction coefficient and the concentration, respectively. Considering a mass balance equation, cf = cT − cb = cT − ncLi, where cT, cb, and cLi represent the concentrations of total TFEP, bound TFEP, and Li-ion, respectively, and n is the solvation number of TFEP molecules around Li-ion ions. Thus, we can obtain the following equation: If/cT = −nεf (cLi/cT) + εf. Plots of If/cLi against cT/cLi give a straight line with the slope α (= −nεf) and the intercept β (= εf), and the n value is thus obtained according to n = −α/β. The detailed analysis procedure is described in the literature.36, 38 With regard to the Raman spectra, in order to quantitatively determine the n and J values, the measured spectra were normalized by using the band intensities at 1339.0 cm−1 of the TFSA− as an internal standard. DFT calculations. DFT calculations were performed using the Gaussian 09 software package.39 The geometries of the isolated TFEP and the Li-ion complexes were fully optimized at the B3LYP/6-311G** level, followed by their normal frequency analysis. The binding energy ∆Ebind was calculated as the SCF energy difference between the optimized [Li(TFEP)n(TFSA)m] complex and its individual components (TFEP, Li+, and TFSA−) according to ∆Ebind =



ESCF(complex) − ESCF(Li+) − nΕSCF(TFEP) − mΕSCF(TFSA−) and then corrected by basis set superposition error (BSSE) using the counterpoise method.40-41

RESULTS AND DISCUSSION Binding energy for Li+–solvent complex. It is well established that the Gutmann’s donor number DN for solvents, i.e., the electron pair-donating ability to bind to a metal ion, is a good indicator to evaluate metal ion–solvent interactions or solvation power. DN values have been reported and accumulated for various conventional aprotic and protic solvents;34 however, the DN value of TFEP applied as a battery electrolyte has not yet been reported. To confirm the electron pair-donating ability of aprotic TFEP, we first performed DFT calculations for the Li+:TFEP 0

-1

-50 ∆Ebind / kJ mol

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-100 -150 TFEP -200

DMC THF

AN PC

EC

DMF DMSO

-250

HMPA

-300 -350 0

10

20 DN

30

40

Figure 1. The binding energy ∆Ebind values calculated for the 1:1 (Li+:solvent molecule) complexes plotted against the corresponding Gutmann’s donor number DN values34 for tris(2,2,2-trifluoroethyl) phosphate (TFEP), acetonitrile (AN), ethylene carbonate (EC), dimethyl carbonate (DMC), propylene carbonate (PC), tetrahydrofran (THF), N,N-dimethylformamide (DMF), dimethylsulfoxide (DMSO) and hexamethylphosphoric triamide (HMPA).


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(1:1) complex to estimate ∆Ebind. Figure 1 shows the ∆Ebind values for various aprotic solvent molecules acting as a monodentate ligand in the Li-ion coordination, including TFEP, plotted against the corresponding DN values. We found that the ∆Ebind values for conventional aprotic solvents [e.g., acetonitrile (AN), EC, DMC, propylene carbonate (PC), tetrahydrofran (THF), N,N-dimethylformamide (DMF), dimethylsulfoxide (DMSO), and hexamethylphosphoric triamide (HMPA)] having no electron pair-accepting site within each molecule, fall on a straight line. Here, note that TFEP is also a typical aprotic solvent with one electron pair-donating site at the P=O group and no accepting site; therefore, the ∆Ebind value for the TFEP system can be used to predict the DN value using the straight line shown in Figure 1. The TFEP molecule showed ∆Ebind = −175.4 kJ mol−1, i.e., the energy stability of the 1:1 complex was the lowest among all the aprotic solvent systems examined in this research. The DN value of TFEP was thus estimated to be approximately 12.9. This result implies that TFEP can be categorized as a weak donor solvent that should give weak ion–solvent interactions while in solution. Walden plots: Ionic conductivity and viscosity. It is well established that a plot based on the Walden rule (Ληα = constant) proposed by Angell et al. is useful for discussing the dissociativity of salts in electrolyte solutions, where Λ and η are the molar conductivity and the viscosity, respectively, and α represents the adjustable parameters.42 Figure 2 shows the Walden plots as a form of log Λ vs. log η−1 for the LiTFSA/TFEP solutions (cLi = 0.2 and 0.5 mol dm−3). Here, note that the vertical deviation of experimental plots from an ideal line of aqueous KCl solution indicates the dissociativity (or ionicity) of the salt in the solution; i.e., the Walden plots for electrolytes with highly dissociated salts are located close to an ideal KCl line, and the distance is more when ions are associated to form ion pairs.42-43 As shown in Figure 2, in conventional electrolyte systems [i.e., LiTFSA/DMF, LiTFSA/H2O, and LiTFSA/EC+DMC (1:2 by volume)1],



-1

4 3

-1

log Λ / Scm mol

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LiPF6/EC+DMC LiTFSA/EC+DMC H2O DMF

2 1 0

TFEP 0

1

2 -1

log η / poise

3

4

-1

Figure 2. Walden plots for 0.2 and 0.5 mol dm−3 LiTFSA/TFEP solution (red triangle and circle), together with 0.5 and 1.0 mol dm−3 LiTFSA/DMF (open and filled diamonds), 1.0 mol dm−3 LiTFSA/water (square), 1.0 mol dm−3 LiTFSA/EC+DMC1 (open triangle), and 1.0 mol dm−3 LiPF6/EC+DMC1 (filled triangle). including a commercially used system for LIBs (LiPF6/EC+DMC)1, the Walden plots were close to the ideal KCl line. In conventional solvents, the Li-ion is solvated by solvent molecules below cLi < 1.0 mol dm−3 to form a tetrahedral [Li(solvent)4]+ complex. In other words, no contact ionpair formation occurs due to interactions between Li-ion and counter-anion. In the current LiTFSA/TFEP system, the Walden plots considerably deviated in the vertical direction from an ideal line even in the diluted solutions (cLi = 0.2 and 0.5 mol dm−3). This strongly suggests that the TFSA− ion directly binds to Li-ion to form a contact ion pair (Li+⋅⋅⋅TFSA−). To elucidate the Li-ion complexes from a structural viewpoint, we performed ATR-IR and Raman spectroscopic studies on the TFEP-based electrolyte solutions to determine the individual solvation (or coordination) numbers of TFEP and TFSA− in the Li-ion complex [Li(TFEP)n(TFSA)m](1 − m).


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Li-ion solvation structure. Figure 3a shows the IR spectra observed for the LiTFSA/TFEP solutions in a range of 755–1000 cm−1. The peak at around 887.0–907.2 cm−1 in neat TFEP (blue line), assigned to the O–P–O asymmetric stretching mode of TFEP in the bulk (hereafter, free TFEP), gradually decreased in intensity with increasing cLi. In contrast, a new peak appeared at the higher frequency side (903.5 cm−1), which intensified with increasing cLi. According to the DFT calculations, the 903.5 cm−1-band is ascribed to TFEP bound to Li-ion (bound TFEP), which will be comprehensively discussed later. We note that the isosbestic point is clearly apparent at 907.0 cm−1, indicating that the TFEP molecules coexisted as two species, free and bound TFEP, while in solution. A similar behavior was observed in the 785–825 cm−1 range. The free band at 803.6 cm−1 decreased and the bound band at 793.7 cm−1 increased in intensity with increasing cLi, which also exhibited an isosbestic point at 798.0 cm−1. We then focused the bands around 887.0–907.2 cm−1 and performed quantitative data analysis using a least-squares curve fitting for the observed IR spectra to deconvolute the signal into individual bands for each

Figure 3. (a) Infrared (IR) spectra observed for the LiTFSA/TFEP solutions with varying cLi and (b) a typical curve-fitting result for the cLi = 1.0 mol dm−3 solution.



component. Here, note that the spectrum observed in neat TFEP (cLi = 0 mol dm−3) was appreciably distorted in shape. Therefore, we assumed herein that the IR band at around 850–950 cm−1 (free TFEP) can be represented by two components. The fitting parameters (half-width at half-maximum and Gauss/Lorentz ratio on pseudo-Voigt functions) and also the ratio of the peak height between the two free bands were fixed for all the cLi during the current fitting analysis. The typical result for cLi = 1.0 mol dm−3 is shown in Figure 3b. The IR spectrum for cLi = 1.0 mol dm−3 could be deconvoluted into free TFEP (blue line: 887.0 and 907.2 cm−1) and bound TFEP (red line: 903.5 cm−1) components, which was also applied to the IR spectra for every cLi. We thus estimated the solvation number of the Li-ions nTFEP by analyzing the sum of integrated intensities for 887.0 cm−1- and 907.2 cm−1-bands (i.e., If = I887 + I907) as a function of cLi. Figure 4 shows a plot of If/cT vs. cLi/cT based on the equation described in the Experimental Section. For every cLi examined here, the plots fall on a straight line with a slope (−nJf) and intercept (Jf) to be nTFEP = 2.0 ± 0.1 according to n = −slope/intercept. It is well known that in conventional

If / cT

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0.00 0.05 0.10 0.15 0.20 0.25 cLi / cT

Figure 4. If/cT vs. ILi/cT plot using the deconvoluted bands (If = I887 + I907) of the free TFEP.


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aprotic solvents, an Li-ion is four-coordinated with solvent molecules to form a tetrahedral [Li(solvent)4]+ complex.44-45 However, in the current TFEP system, the result of nTFEP = 2.0 suggested that the Li-ion was solvated with two TFEP molecules. Considering the Walden plots (ionic conductivity and viscosity data) discussed above, we expected that the Li-ion might be coordinated with both the TFEP and TFSA components, resulting in the formation of the [Li(TFEP)2(TFSA)m] ion-pair complex. Therefore, we investigated the TFSA coordination to the Li-ion in TFEP-based solutions using Raman spectroscopy. Figure 5a shows the Raman spectra for the LiTFSA/TFEP solutions with varying cLi in the 735–765 cm−1 range. The peak at 748.5 cm−1 was monotonically intensified with increasing cLi, which was ascribed to the CF3 bending vibration coupled with the S–N–S symmetric stretching vibration of TFSA−.36 Similar to the data analysis for the IR spectra as described above, we performed curve fitting analysis for the observed Raman spectra in this frequency range. The typical result (cLi = 1.0 mol dm−3) is shown in Figure 5b. The Raman spectrum could be

Figure 5. (a) Raman spectra for the LiTFSA/TFEP solutions with varying cLi; (b) a typical curve fitting result for the cLi = 1.0 mol dm-3 solution.



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deconvoluted into two components at 746.1 cm−1 (blue line) and 748.9 cm−1 (red line). Here we also measured the Raman spectrum for the TFEP solution containing 1-ethyl-3methylimidazolium TFSA IL ([C2mIm][TFSA]) as an indicator of “completely dissociated TFSA−” owing to the delocalized C2mIm+ cation with a larger ionic size, which assists in understanding ion-pair formation between the Li-ion and TFSA (see, Figure S1). In the [C2mIm][TFSA]/TFEP solution, the dissociated TFSA (i.e., free TFSA) exhibited a peak at 746.1 cm−1. We thus concluded that in the LiTFSA/TFEP solution, the minor band (746.1 cm−1) corresponds to the free TFSA and the major band (748.5 cm−1) to the bound TFSA. To estimate concentrations of free and bound TFSA species, we performed the following estimation using the observed intensity and Raman scattering coefficient. In previous work, we reported that the ratio of the Raman scattering coefficients, Jf/Jb, is approximately constant (~ 0.9) in aprotic solvent36 and ionic liquid46 solutions containing LiTFSA salt. We applied this value (Jf/Jb = 0.9) to estimate the concentrations of free and bound TFSA (cf and cb, respectively) through a relation: cT = cf + cb, resulting in the following equation, (Ib/If)(Jf/Jb) = cb/cf. The cb/cT (concentration ratio of the bound TFSA to the total TFSA) was determined to be approximately 0.8 for all the cLi solutions examined here, listed in Table S2. This result indicated that 80 % of the total TFSA anions coordinate with the Li-ions to form the Li+⋅⋅⋅TFSA− (1:1) ion pairs, even in dilute cLi solutions. Thus, we concluded that in the LiTFSA/TFEP solutions, the Li-ions mainly exist as [Li(TFEP)2(TFSA)] complexes as a major species, which coexist with ion-pair-free species as a minor species (20%). To gain more insight into the Li-ion solvation structure at the atomistic level in detail, we conducted a theoretical study using DFT calculations for the [Li(TFEP)2(TFSA)] complexes. Figure 6 shows theoretical IR bands calculated for the optimized geometry of the


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[Li(TFEP)2(TFSA)] complex (right-hand side of Figure 6) together with the corresponding experimental bands for the cLi = 1.0 mol dm−3 solution. In the calculations, the TFSA anion, which involves cis and trans conformers in equilibrium,47-49 within the [Li(TFEP)2(TFSA)] complex was treated as a cis conformer because the TFSA anion prefers the cis conformation in the first Li coordination shell according to our previous research.50 It is well known that in general, Li-ions are tetrahedrally four-coordinated with conventional oxygen donor solvents, such as water, alcohol, and DMF, when solvated.44-45 In the optimized [Li(TFEP)2(TFSA)] structure, we found that the Li-ion was coordinated with four oxygen atoms from two monodentate TFEP molecules and one bidentate TFSA anion.

Figure 6. Theoretical IR bands for (a) isolated TFEP and (b) [Li(TFEP)2(TFSA)] complexe optimized in the gas phase by DFT calculations, together with the experimental bands for cLi = 1.0 mol dm−3. (c) The optimized geometry of the [Li(TFEP)2(TFSA)] complex.

As mentioned above, the experimental IR spectrum exhibited three bands: two of free TFEP [887.0 and 907.2 cm−1; blue lines in Figure 6 (obs.)] and one of bound TFEP (903.5 cm−1; red



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line). With regard to the free TFEP, the 887.0 cm−1-band was successfully reproduced by the theoretical band originated from the O−P−O asymmetric stretching mode of TFEP molecule (Figure 6a). Here, note that the TFEP has three O−P−O groups within a molecule; however, the normal frequency analysis based on the optimized TFEP structure gave only one theoretical band (actually, very near two theoretical bands at 884.2 and 884.8 cm−1). This is because the geometrical symmetry of the optimized TFEP molecule is so high to give essentially similar frequencies among the three O−P−O groups. In the current DFT study, we could not assign the 907.2 cm−1-band (free TFSA). In speculation, it seemed that the 907.2 cm-1-band is due to conformational TFEP-isomers in solutions. In solution system, it is possible that the TFEP molecules show some stable conformers because of its complexed molecular structure and molecular flexibility. Hence, the frequency position of the O−P−O vibration might be distributed widely, resulting in the additional observed band around 907.2 cm-1; however, it was not found in the theoretical band for the symmetrical TFEP structure in gas phase calculations. In Figure 6b, the [Li(TFEP)2(TFSA)] complex showed the theoretical bands in two frequency regions at around 880 and 900 cm−1. Note that in the observed spectrum, bound TFEP exhibited a broad band, and therefore, the frequency was widely distributed around 903.5 cm−1. This observation is consistent with the theoretical results of this research; i.e., the coordination of TFEP molecules with the Li-ion makes their vibrations (O–P–O asymmetric stretching modes) more complicated. This separates their bands into several components that locate over a wide frequency region, resulting in a single broad band for bound TFEP in the observed IR spectrum. Based on our experimental and theoretical investigations, we propose that Li-ions are coordinated with both TFEP (solvent) and TFSA (counteranion) to mainly exist as the contact ion pair [Li(TFEP)2(TFSA)], which is unusual in typical nonaqueous dipolar solvents, such as


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DMF, DMSO, and AN. This might be because the TFEP molecule is a significantly bulky solvent molecule with weak electron pair-donating abilities. To clarify this issue, we performed additional DFT calculations for the [Li(TFEP)n] complexes with varying solvation number (n) to estimate their ∆Ebind corresponding to the solvation stability using an energy scale. Figure 7 shows the ∆Ebind values calculated for the [Li(TFEP)n]+ complexes (considering monodentatetype coordination) with n = 1–4. The ∆Ebind value monotonically decreased with increasing n up to 2. However, the decrease became appreciably moderate at n = 3, leading to no or little energy difference with further increase in the value of n to 4. The energy variation can be explained in terms of “steric repulsion” among the solvated TFEP molecules in the first solvation sphere as follows: (1) when n ≤ 2, the interactions between Li-ion and oxygen atoms at the O=P group (TFEP) mainly contribute to the ∆Ebind value. In particular, focusing on the n = 2 system, no steric repulsion among the two solvated TFEP molecules occurs in the [Li(TFEP)2]+ complex

0

-1

-200 ∆Ebind / kJ mol

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-400 -600 -800 0

[Li(TFEP)n] [Li(TFEP)2(TFSA)]

1

2 3 Solvation number, n

4

Figure 7. ∆Ebind plotted against n for the [Li(TFEP)n]+ complexes and the [Li(TFEP)2(TFSA)] complex.



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because there is sufficient space to locate them, as shown in Figure S2. (2) At n = 3, the Li⋅⋅⋅O interactions as well as the steric repulsion due to the bulkiness of TFEP contribute to ∆Ebind for the [Li(TFEP)3]+ complex (Figure S2). (3) Finally, the repulsive effect dominates in the n = 4 system. In other words, significant steric repulsion cancels the stabilization gained from an increase in the value of n, resulting in no change in the ∆Ebind value, even with an increase in the value of n from 3 to 4. It is worth noting that in the case of n = 3, the energy of the solvation complex significantly decreases by replacing the TFEP molecule with the TFSA anion (red circle in Figure 7). [Li(TFEP)2(TFSA)], which is the main species in the LiTFSA/TFEP solutions, as discussed above, is most stable among the complexes examined herein because of two factors: (1) the steric repulsion becomes weaker by replacing bulky TFEP with the smaller TFSA anion and (2) the electrostatic cation–anion interactions between Li+ and TFSA− play a key role in Liion stabilization in addition to the local Li+⋅⋅⋅O interactions. The solvation steric effect, i.e., the repulsive force among the solvated solvent molecules in the solvation sphere, is an important factor for controlling the solvation stability and reactivity of metal ions in solutions. Ishiguro et al. reported that for first-row transition metal ions, an octahedral six-coordination of solvent molecules is sterically hindered in bulky solvents to reduce the solvation number to approximately 4 (e.g., zinc(II) ion in N,N-dimethylpropionamide).51-52 This mechanism can be applied to the current LiTFSA/TFEP system. The significant bulkiness of TFEP decreases the solvation number (generally four-coordinated in conventional solvents). Instead, the smaller TFSA anion coordinates with the Li-ion through ion pairing while in solution.

CONCLUSIONS


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In this study, the solvation structure of the Li-ion in TFEP-based electrolytes was investigated using vibrational spectroscopic experiments and DFT calculations. The experimental results state that (1) TFEP had a weaker electron pair-donating ability and larger molecular size. (2) Thus, the Li-ion complex solvated only by solvent molecules, such as ([Li(TFEP)n]+), is unfavorable even in dilute solutions. (3) The Li-ions are coordinated with both TFEP (monodentate manner) and TFSA anion (bidentate manner) to form the tetrahedral [Li(TFEP)2(TFSA)] complex (contact ion pair) as a main species. It is clear that such ion pairing is not favorable for designing highperformance LIBs due to the resulting decrease in ionic conductivity; however, it possesses a strong advantage (i.e., nonflammability and excellent thermal stability) for developing safer LIB electrolytes. Thus, a strategy for suppressing contact ion-pair formation must be devised, i.e., the Li salt dissociativity in TFEP-based solutions must be increased. To address this issue, an extended study for establishing “ion-pair-free” TFEP-based electrolytes is currently underway.

ASSOCIATED CONTENT Supporting Information. The Supporting Information is available free of charge on the ACS Publication website at DOI: cLi, density, and refractive index values (Table S1); Ib/If and cb/cT (Table S2); Raman spectra observed for the 1.0 mol dm−3 [C2mIm][TFSA]/TFEP (Figure S1); optimized geometries of [Li(TFEP)n] (n = 1–4) obtained by DFT calculations (Figure S2).

AUTHOR INFORMATION Corresponding Author



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*E-mail: [email protected]

ACKNOWLEDGMENT This research has been financially supported by Grant-in-Aids for Scientific Research from the Ministry of Education, Culture, Sports, Science and Technology (No. 15K17877 to K.F.).

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Table of Contents Only

Role of Solvent Bulkiness on Lithium-ion Solvation in Fluorinated Alkyl Phosphate-based Electrolytes: Structural Study for Designing Nonflammable Lithium-ion Batteries Michiru Sogawa,a Yanko M. Todorov, a Daisuke Hirayama, b Hideyuki Mimura, b Nobuko Yoshimoto, a Masayuki Morita, a and Kenta Fujii a,*

[Li(TFEP)2(TFSA)]


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Role of Solvent Bulkiness on Lithium-Ion Solvation ... - ACS Publications

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