ARTICLE pubs.acs.org/JPCC
Change from Glyme Solutions to Quasi-ionic Liquids for Binary Mixtures Consisting of Lithium Bis(trifluoromethanesulfonyl)amide and Glymes Kazuki Yoshida, Mizuho Tsuchiya, Naoki Tachikawa, Kaoru Dokko, and Masayoshi Watanabe* Department of Chemistry and Biotechnology, Yokohama National University, 79-5 Tokiwadai, Hodogaya-ku, Yokohama 240-8501, Japan
bS Supporting Information ABSTRACT: The physicochemical properties of triglyme (G3) and tetraglyme (G4) solutions containing lithium bis(trifluoromethanesulfonyl)amide (LiTFSA) were investigated. The concentration of LiTFSA was altered from a dilute solution to an extremely concentrated one, forming the complexes [Li(glyme)1][TFSA]. The ionic conductivity of the LiTFSA/glyme solutions changed depending on the LiTFSA concentration and exhibited a maximum at ca. 1 mol 3 dm3. The viscosities of the solutions monotonically increased with increasing concentration of LiTFSA. The self-diffusion coefficient of each species in the solutions, which was measured by the pulsed-field gradient spinecho NMR (PGSE-NMR) method, decreased with increasing concentration. The diffusion coefficient of the glymes in relatively dilute LiTFSA solutions was higher than that of the Li+ cation and the [TFSA] anion. As the concentration of LiTFSA increased, the difference between diffusion coefficients of the glyme and Li+ cation became less pronounced. The identical diffusion coefficient of the glyme and Li+ cation in the equimolar mixture of glymeLiTFSA suggested the formation of a complex cation [Li(glyme)1]+ in the liquid. Unexpectedly, the ionicity Λimp/ΛNMR (dissociativity) of [Li(glyme)1][TFSA] was higher than that of excess glyme solutions of [Li(glyme)x][TFSA] (x > 1). This result is similar to that observed for mixtures of typical ionic liquids and cosolvents and is opposite to that observed for conventional organic electrolyte solutions. The evaporation of glyme from the quasi-ionic liquid [Li(glyme)1][TFSA] took place at ca. 200 °C, and the activation energy for the evaporation rate of the glymes from [Li(glyme)1][TFSA] was estimated to be ca. 63 kJ 3 mol1, which appears to correspond to the activation energy for the desolvation of the glyme from [Li(glyme)1][TFSA]. The electrochemical impedance measurements revealed that the activation energy for the charge-transfer reaction of Li/Li+ at the liquid | Li metal interface was ca. 68 kJ 3 mol1, which is in good agreement with that for desolvation of the glymes.
’ INTRODUCTION Lithium ion batteries (LIBs), with high energy and power densities, have become essential to modern society as power sources for portable electronic devices.13 Furthermore, electric vehicles (EVs) equipped with LIBs are now commercialized. The conventional electrolyte of LIBs is composed of mixed organic solvents (cyclic carbonate and linear carbonate) and LiPF6.4 Linear carbonate solvents are extremely flammable, with flash points below room temperature, and LiPF6 decomposes into HF with the presence of a slight amount of water at temperatures higher than 60 °C.5 Thus, LIBs have thermal instability problems at elevated temperatures. To achieve a high degree of safety of LIBs, the development of thermally stable electrolytes is crucial. Room-temperature ionic liquids (RTILs), which consist entirely of cations and anions, have attracted much attention owing to their unique properties such as low volatility, high thermal stability, high ionic conductivity, wide potential window, and high chemical stability.57 RTILs are called “designer solvents” because their physicochemical properties can be easily tuned r 2011 American Chemical Society
simply by changing the structures of the component cations and anions. For example, ionic liquids possessing active protons are obtained from the proton transfer reactions from Brønsted acids to bases, and the resulting protic ionic liquids have a potential as electrolytes of mesothermal fuel cells.811 Furthermore, aprotic RTILs have been extensively studied as electrolytes for LIBs.1225 To replace the flammable organic solvents, onium cation-based aprotic RTILs such as quaternary ammonium and imidazolium RTILs have been adopted as “solvents” for dissolving Li salts for use in LIBs. However, the viscosities of the RTILs increase with increasing lithium salt concentration. Moreover, concentration polarization takes place in the RTILs during chargedischarge of battery because there exist two cationic species, a Li+ cation and an onium cation, in the electrolytes. The Li+ cation transference number is generally very low, less than 0.1,25 and maximum Received: July 19, 2011 Revised: August 7, 2011 Published: August 10, 2011 18384
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Table 1. Concentration Dependence of Density for LiTFSA/ Glyme Solutions at 30 °C C, mol 3 dm3
F, g 3 cm3
[Li(G3)1][TFSA]
3.14
1.46
[Li(G3)2][TFSA] [Li(G3)4][TFSA]
2.00 1.16
1.29 1.16
[Li(G3)8][TFSA]
0.63
1.08
[Li(G3)30][TFSA]
0.18
1.01
[Li(G3)90][TFSA]
0.06
0.99
[Li(G3)400][TFSA]
0.01
0.98
[Li(G4)1][TFSA]
2.75
1.40
[Li(G4)2][TFSA] [Li(G4)4][TFSA]
1.73 0.98
1.26 1.16
[Li(G4)8][TFSA]
0.52
1.08
[Li(G4)30][TFSA]
0.15
1.03
[Li(G4)90][TFSA]
0.05
1.01
[Li(G4)400][TFSA]
0.01
1.00
current density based on the Li+ migration in the RTILs does not meet the requirements for practical use in LIBs. To achieve a high transference number of Li+ cation in RTILs, we prepared roomtemperature lithium ionic liquids consisting of Li+ cations and borate anions having electron-withdrawing groups, which reduces the anionic basicity, as well as lithium-coordinating ether ligands to dissociate the lithium cations from the anionic centers.26,27 However, the ionic conductivity of the lithium ionic liquid at room temperature was low because of its high viscosity and low degree of dissociation. Glymes are aprotic solvents, and they possess high thermal and chemical stability. They are often used as solvents in organic synthesis. In addition, glymes can easily dissolve alkali metal salts because glymes have high donor numbers. It is well-known that certain molar ratio mixtures of Li salts and oligo-ethers such as crown ethers, triglyme (G3), and tetraglyme (G4) form complexes.2831 The crystallographic structures, thermal properties, and ionic conductivity of solid-state glymeLi salt complexes have been reported.3235 Among the series of glymeLi salt complexes, [Li(G3)1][TFSA] and [Li(G4)1][TFSA] remain in a liquid state at room temperature. The melting point of [Li(G3)1][TFSA] is 23 °C and the glass transition temperature of [Li(G4)1][TFSA] is 54 °C.36 We reported the physicochemical properties of molten complexes of glyme (G3 and G4) with Li salts.36,37 The equimolar complexes of glyme (G3 or G4) and Li salts show unique physicochemical properties similar to those of RTILs, such as high thermal stability, low volatility, low flammability, and high ionic conductivity.36 Significantly, a 4 V-class LiCoO2 cathode can be successfully operated in the molten glymeLi salt complexes, making these complexes promising electrolytes for LIBs.37,38 In this study, in order to understand the fundamental properties of molten complexes of glymeLi salt, we investigated the physicochemical properties of glyme solutions containing LiTFSA. It is interesting to study the glyme solutions as model electrolytes because the concentration of LiTFSA can be altered from a dilute solution to an extremely concentrated one; in other words, to the composition of [Li(glyme)1][TFSA]. The solvated structure of the Li+ cation appears to change depending on its concentration.39,40 The dependence of transport properties and
thermal properties of the LiTFSA/glyme binary system on concentration were examined. Furthermore, the interfacial electrochemical process of [Li(glyme)]+ was explored and correlated with the desolvation process of the glyme from [Li(glyme)1][TFSA].
’ EXPERIMENTAL SECTION Purified G3 or G4 (Kishida Chemical Co., Ltd.) and LiTFSA (Morita Chemical Industries Co., Ltd.) were mixed and stirred for 24 h at room temperature, and homogeneous liquids were obtained. The mixtures were stored and handled in an argonfilled glovebox ([H2O] < 1 ppm). The residual water in the mixture was measured with a Karl Fischer moisture meter (Mitsubishi CA-07) and was determined to be less than 30 ppm. The glymeLiTFSA mixtures are denoted as [Li(glyme)x][TFSA], depending on the molar ratio x = [glyme]/[LiTFSA] as shown in Table 1. In certain experiments, diethyl carbonate (DEC, Kishida Chemical Co., Ltd.) was added to the glymeLiTFSA mixtures. Ionic conductivities (σ) of [Li(glyme)x][TFSA] were determined by the complex impedance method using an ac impedance analyzer (Princeton Applied Research, VMP2) in the frequency range of 500 kHz1 Hz with a sinusoidal alternating voltage amplitude of 10 mV root-mean-square (rms). A cell equipped with two electrodes of platinized platinum for conductivity measurements (TOA Electronics, CG-511B) was utilized, and the cell constant was determined by a standard 0.01 mol 3 dm3 KCl aqueous solution (Kanto Kagaku) at 25 °C. The cell was placed in a temperature-controlled chamber, and conductivity was measured at different temperatures with heating from 10 to 90 °C. The density measurements were performed with a thermoregulated density/specific gravity meter DA-100 (Kyoto Electronics Manufacturing Co. Ltd.). The density was measured in the range of 1540 °C. Viscosity measurements were performed with a rheometer (Physica MCR301, Anton Paar) with a coneplate under dry atmosphere. The viscosity (η) was measured at various temperatures with heating from 10 to 90 °C. Pulsed-field gradient spinecho (PGSE) NMR measurements were carried out to determine the self-diffusion coefficients of [Li(glyme)x ][TFSA]. A JEOL-AL 400 NMR spectrometer with a 9.4 T narrow-bore superconducting magnet equipped with a pulsed-field gradient probe and a current amplifier was used for this purpose. 1H, 7Li, and 19F spectra were measured for glymes, Li+ cations, and TFSA anions, respectively. The self-diffusion coefficients were measured by use of a simple Hahn spinecho sequence. Detailed measurement procedures for PGSE-NMR are described elsewhere.4144 The free diffusion echo signal attenuation E is related to the experimental parameters by the Stejskal equation: ln ðEÞ ¼ ln ðS=Sδ ¼ 0 Þ ¼
γ2 g 2 Dδ2 ð4Δ δÞ π2
ð1Þ
where S is the spinecho signal intensity, δ is the duration of the field gradient with magnitude g, γ is the gyromagnetic ratio, and Δ is the interval between the two gradient pulses. The value of Δ was set at 50 ms. The sample was inserted into a NMR microtube (BMS-005J, Shigemi) to a height of 3 mm to exclude convection. PGSE-NMR measurements were performed in the temperature range 1090 °C. Thermal stabilities of the glymeLi salt complexes were evaluated by thermogravimetry (TG/DTA 6200, Seiko) under N2 atmosphere. For measuring the temperature dependence of 18385
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Figure 1. Concentration dependency of ionic conductivity for LiTFSA/glyme solutions at 30 °C. Lines in the figure are just guides for the eye.
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ionic conductivity increased as the concentration increased in the range 01 mol 3 dm3 because of the increment of the number of charge carriers in the solution. The ionic conductivity passes through a maximum at ca. 1 mol 3 dm3 and then decreases in the range ca. 1.23.0 mol 3 dm3. The decrease in the ionic conductivity was caused by a decrease in the ionic mobility due to an increase in viscosity of the solution. The dependence of the ionic conductivity of electrolyte on the Li salt concentration is very similar to that of a conventional aprotic electrolyte such as propylene carbonate (PC) solution.45 As shown in Figure 2, the viscosity of the solution increased dramatically as the concentration increased. The self-diffusion coefficients of glyme, Li+ cation, and [TFSA] anion, measured by PGSE-NMR, are shown in Figure 3. We could not measure the self-diffusion coefficients of Li+ at concentrations less than 0.15 mol 3 dm3 because the NMR signal was too weak to obtain reliable data. At concentrations higher than 0.15 mol 3 dm3, each diffusion coefficient decreased monotonically as the concentration increased because the viscosity of the electrolyte increased. Diffusivity might be correlated to the fluidity (1/η) by the StokesEinstein equation: D¼
Figure 2. Concentration dependency of viscosity for LiTFSA/glyme solutions at 30 °C. Lines in the figure are just guides for the eye.
weight loss, the temperature of the sample was increased from room temperature to 550 °C at a heating rate of 10 °C 3 min1. For isothermal weight loss experiments, the temperature was elevated from room temperature to a predetermined value at a rate of 20 °C 3 min1 and kept constant, and then the weight change was monitored. Electrochemical impedance measurements were carried out for a [Li metal | [Li(glyme)1][TFSA] | Li metal] symmetric configuration encapsulated into 2032-type coin cells. Two Li metal foils (Honjo Metal, 200 μm thick) were cut into a disk shape (16 mm in diameter) and used as electrodes. A glass filter (GA-55, Advantec) was inserted to separate the two Li disk electrodes. The electrolyte was inserted into voids in the glass filter during the cell fabrication. All materials were handled in an argon-filled glovebox (VAC, dew point < 80 °C). Electrochemical impedance measurements were performed at open circuit voltage with a sinusoidal alternating voltage amplitude of 10 mVrms and frequency ranging from 500 kHz to 20 mHz by using the ac impedance analyzer. The temperature was controlled in the range 3050 °C.
’ RESULTS AND DISCUSSION Transport Properties of Lithium Bis(trifluoromethanesulfonyl)amide and Glyme Binary Mixtures. Figure 1 shows
the isothermal ionic conductivity of [Li(glyme)x][TFSA] at 30 °C as a function of molar concentration of LiTFSA. The
kT απηrs
ð2Þ
where k is the Boltzmann constant, T is the absolute temperature, α is a constant, and rs is the effective hydrodynamic (Stokes) radius. The constant α in the case of a large solute in a small solvent can attain a value of 6. If the ratio of the solute size to solvent is increased, especially for highly viscous media, the correlation breaks down, and the value of α in eq 2 is reduced to ca. 4.46 In the concentration range 02.0 mol 3 dm3, the selfdiffusion coefficients of glyme (DG), anion (D), and Li+ cation (DLi+) are in the order DG > D > DLi+, which agrees well with the order of conventional organic electrolytes.47,48 Although the radius of the Li+ cation in typical ionic crystals is between 0.073 and 0.090 nm and is the smallest among the species in the electrolyte, the Li+ cation solvated by glyme results in the largest hydrodynamic volume among the components, leading to the order DG > D > DLi+. The difference between DG and DLi+ becomes less pronounced with increasing LiTFSA concentrations. Glyme molecules are relatively strong Lewis bases and strongly solvate Li+ cations with strong Lewis acidity. G3 and G4 have 4 and 5 oxygen atoms in their molecular structures, respectively, and G3 and G4 coordinate to Li+ to form complex cations [Li(glyme)1]+ with coordination numbers of 4 and 5, respectively. NMR cannot distinguish the free glyme from the Li+-coordinated glyme in the solutions, which indicates that the ligand exchange takes place very quickly in the solution.38 glyme þ ½LiðglymeÞ1 þ f ½LiðglymeÞ1 þ þ glyme
ð3Þ
Thus, the diffusion coefficient of glyme determined by PGSENMR is the average of values for free and Li+-coordinated glyme in the system. As the concentration of LiTFSA increases, the number of free glymes with greater diffusivity decreases and the number of Li+-coordinated glymes with lesser diffusivity increases. In the case of an equimolar mixture of LiTFSA and glyme, all of the glyme molecules coordinate with the Li+ cation to form a [Li(glyme)1]+ complex cation, the concentration of free glyme is close to zero, and consequently the ligand exchange scarcely occurs. Indeed, the DG and DLi+ are identical in the 18386
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Figure 3. Concentration dependency of self-diffusion coefficient of glyme, [TFSA] anion, and Li+ cation at 30 °C for LiTFSA/glyme solutions. (Left) G3 system; (right) G4 system.
Table 2. Ionic Conductivity and Diffusion Coefficient of Glyme, DEC, Li+ Cation, and [TFSA] Anion at 30 °C in the Mixtures of [Li(G3 or G4)][TFSA]:DEC = 1:4. diffusion coefficient, 107 cm2 3 s1 σ, mS 3 cm1 Dglyme
DDEC
DLi+
D
G3:LiTFSA:DEC = 1:1:4
5.5
18
48
18
19
G4:LiTFSA:DEC = 1:1:4
6.1
18
48
18
19
equimolar mixture of glymeLiTFSA ([Li(glyme)1][TFSA]). The equimolar mixture of glyme and LiTFSA can be regarded as a quasi-RTIL consisting of a [Li(glyme)1 ]+ cation and a [TFSA] anion. To confirm the complex cation formation of Li+ with glyme in a 1:1 ratio, the equimolar mixture was diluted with diethyl carbonate (DEC), which has a much lower donor number than glymes. DEC is often mixed with other carbonate solvents such as PC and ethylene carbonate to reduce the viscosity of the electrolyte in practical lithium batteries. The self-diffusion coefficients of glyme, DEC, Li+ cation, and [TFSA] anion in the solution measured by PGSE-NMR are shown in Table 2. The diffusion coefficients of DG and DLi+ are identical, while DEC shows a much higher diffusion coefficient than the other species, suggesting that the Li+ cation and glyme diffuse together in the solution. This result strongly supports the observation that the glyme molecules preferentially coordinate with Li+ cations to form the complex cation [Li(glyme)1]+ and DEC does not participate in the solvation. Furthermore, the perfect accord of DLi+ with DG indicates that the ligand exchange between the complex cations scarcely occurs in the solution. Therefore, it is considered that the complex structure of [Li(glyme)1]+ is also retained in the equimolar mixture of glyme (G3 or G4) LiTFSA (without DEC), and the equimolar mixture behaves like an RTIL. The dissociativity of LiTFSA in the glyme solutions is explored here by means of impedance and PGSE-NMR measurements. The molar conductivity (Λimp) of the solution is defined as follows: σ ð4Þ Λimp ¼ c where c is the molar concentration of the electrolyte and σ is the conductivity measured by the electrochemical impedance
Figure 4. Concentration dependency of ionicity (Λimp/ΛNMR) at 30 °C for LiTFSA/glyme solutions. Lines in the figure are just guides for the eye.
method. On the other hand, molar conductivity (ΛNMR) of the electrolyte can be calculated from the PGSE-NMR self-diffusion coefficients of the cation and anion by use of the NernstEinstein equation: ΛNMR ¼
F2 ðDþ þ D Þ RT
ð5Þ
where F is the Faraday constant, R is the gas constant, T is the absolute temperature, and D+ and D are the self-diffusion coefficients of cation and anion, respectively. There have been plenty of studies on the molar conductivity ratio Λimp/ΛNMR for electrolyte solutions47,48 and RTILs.4144 In the case of electrolyte solutions, the Λimp/ΛNMR ratios reflect the degree of dissociation of salt in the solution.47,48 When a low-permittivity solvent such as dimethyl carbonate is used, its Λimp/ΛNMR is low.47 In contrast, high Λimp/ΛNMR ratios are obtained in high-permittivity solvents.47 Hence, by measuring Λimp/ΛNMR values, the dissociativity of electrolyte solutions can be quantitatively discussed. Such Λimp/ΛNMR ratios have also been reported for common RTILs.4144 The Λimp/ΛNMR ratios as well as the deviations of Λimp from the Walden plots of an ideal KCl aqueous solution49,50 are called “ionicity.” The ionicity values of RTILs are also influenced by many factors. We found that the Λimp/ΛNMR ratios decrease when Lewis basicity and acidity of anions and cations, respectively, increase and 18387
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Figure 5. TG curves for [Li(G3 or G4)x][TFSA] (x = 1, 2, 4, 8) and neat glymes at a heating rate of 10 °C 3 min1. (Left) G3 system; (right) G4 system.
Table 3. Density Equation (G = b aT) Parameters for [Li(glyme)1][TFSA]
Figure 6. Temperature dependence of density for [Li(G3)1][TFSA] and [Li(G4)1][TFSA]. Lines in the figure are based on best-fit parameters and the equation in Table 3.
when the alkyl chain length on the imidazolium cations becomes longer.4144 Although the physicochemical implication of the ionicity of RTILs, Λimp/ΛNMR, remains a matter of controversy,4144,4953 it should be of great interest to present reliable Λimp/ΛNMR values for the LiTFSA/glyme binary mixtures since we can alter the concentration from a dilute solution to a quasi-RTIL level ([Li(glyme)1][TFSA]). Unexpectedly, the Λimp/ΛNMR ratios increased with an increase in the LiTFSA concentration in the range 0.153.2 mol 3 dm3, as shown in Figure 4. The highest concentrations in the figure correspond to [Li(glyme)1][TFSA]. It should be noted that similar phenomena are observed for mixtures of RTILs and cosolvents.54,55 On the contrary, in PC solutions of LiPF6 and LiBF4, the Λimp/ ΛNMR ratios are found to decrease with increasing salt concentrations (0.13.0 mol 3 dm3).56 Although we have not clearly understood these quite opposite phenomena against an increment of salt concentration, the Λimp/ΛNMR decrease in the PC solutions appears to be typical because of enhanced Coulombic interaction between cations and anions with increasing salt concentrations. The unique behavior in RTILs might be interpreted in terms of nanoinhomogeneity in the mixtures.57,58 Even if the appearance of RTIL/solvent mixtures is homogeneous, the persistence of ionic clusters from the RTIL has been pointed out by several authors, especially in imidazolium-based RTILs.57,58 The similar behavior of [Li(glyme)1][TFSA]/glyme mixtures to RTIL/solvent mixtures in terms of the Λimp/ΛNMR change may
a, 103 g 3 cm3 3 K1
b, g 3 cm3
R2
[Li(G3)1][TFSA]
1.000
1.728
1.000
[Li(G4)1][TFSA]
1.074
1.724
0.9996
suggest that [Li(glyme)1][TFSA] behaves like an RTIL. [Li(glyme)1][TFSA] dissociates into [Li(glyme)1]+ and [TFSA] by themselves and is stabilized by Coulombic interaction, and no more glyme molecules are necessarily required for the solvation of Li+ and [TFSA], which may induce nanoinhomogeneity in the mixtures. Such an inhomogeneous structure of the mixtures with excess glyme might cause an apparent decrease in Λimp/ΛNMR compared with that of [Li(glyme)1][TFSA]. An alternative explanation might be the formation of a network of Li+ cations cross-linked by the excess glyme molecules, which simultaneously interact with several Li+ cations.59 If the cross-linked structure is formed in the liquid, the hydrodynamic radius of Li+ should be much larger than those of [Li(glyme)1]+, which would lead to high viscosity, small diffusivity of Li+, and a smaller transference number of Li+ than that of [TFSA] anion. However, as shown in Figures 2 and 3, viscosity is decreased and diffusivity of Li+ is increased by excess glyme. On the other hand, the transference number of Li+ in [Li(glyme)1][TFSA]/glyme mixtures is decreased slightly with increasing excess glyme (Supporting Information, Figure S1). The possible formation of the cross-linked structure in [Li(glyme)1][TFSA]/glyme mixtures cannot be ruled out; however, from the results of viscosity and diffusivity measurements, it is postulated that the cross-linked structure is not major in the mixtures. Thermal Stability and Temperature Dependency of Transport Properties for [Li(glyme)1][TFSA]. Thermal stabilities of mixtures of glyme and LiTFSA were investigated by thermogravimetry. Figure 5 shows thermogravimetric (TG) curves for mixtures of glymes and LiTFSA. The weight loss of mixtures with excess glyme ([Li(glyme)x][TFSA] (x > 1)) starts at ca. 100 °C because of evaporation of glymes. The weight loss of equimolar mixtures [Li(G3)1][TFSA] and [Li(G4)1][TFSA] starts at ca. 200 °C. The weight loss of [Li(glyme)1][TFSA] between 200 and 400 °C was attributed to evaporation of glyme, and the rapid decrease at ca. 400 °C was due to decomposition of LiTFSA. 18388
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Figure 7. Temperature dependence of ionic conductivity (left) and viscosity (right) for [Li(G3)1][TFSA] and [Li(G4)1][TFSA]. Lines in the figures are based on best-fit parameters in Tables 4 and 5 and VFT equations.
Table 4. VFT Equation Parameters of Ionic Conductivity Data σ0, 101 S 3 cm1
B , 102K
T0, K
[Li(G3)1][TFSA]
2.6 ( 0.1
7.1 ( 0.2
174 ( 1.8
[Li(G4)1][TFSA]
2.2 ( 0.1
5.3 ( 0.1
196 ( 1.3
Clearly, the thermal stability of [Li(G3)1][TFSA] and [Li(G4)1][TFSA] was improved on comparison with mixtures with excess glyme. Of course, the weight loss temperature of ca. 200 °C for [Li(G3)1][TFSA] and [Li(G4)1][TFSA] is lower than that of typical RTILs, but the temperature is much higher than that of organic electrolyte solutions used in lithium batteries, which would contribute to enhancing the thermal stability of lithium batteries. A higher temperature is needed to evaporate Li+coordinated glyme than that for free and noncoordinated glyme. This is because the desolvation process of Li+ needs to break iondipole interactions and electron donoracceptor interactions between the ether oxygens and the Li+ cation before evaporation of glyme from [Li(glyme)1][TFSA], and a much higher activation energy is needed for desolvation of Li+ than for evaporation of free glyme. Indeed, the weight losses of mixtures of glymeLiTFSA of x > 1 occur in three steps. First, the excess glyme mixtures lose weight at 140 °C due to evaporation of free glyme to yield [Li(glyme)1][TFSA] at ca. 200 °C. Second, desolvated glyme molecules evaporate from [Li(glyme)1][TFSA] in the temperature range 200400 °C to yield LiTFSA at 400 °C. Finally, decomposition of LiTFSA occurs at 400 °C. From these results, the formation of quasi-ionic liquid [Li(glyme)1][TFSA] can also be confirmed from the thermal analysis. Hereafter, the fundamental transport properties of neat [Li(glyme)1][TFSA] are presented. The temperature dependence of density (F) for the [Li(G3 or G4)1][TFSA], as depicted in Figure 6, shows a linear change with temperature. Table 3 lists the best linear-fitting parameters for densitytemperature profiles with the following equation:
Table 5. VFT Equation Parameters of Viscosity Data η0, 101 mPa 3 s
B , 102K
T0 , K
[Li(G3)1][TFSA]
1.2 ( 0.1
10.2 ( 0.1
162 ( 0.8
[Li(G4)1][TFSA]
6.5 ( 0.4
5.3 ( 0.1
194 ( 1.2
interaction between cation and anion in [Li(G3)1][TFSA] than that in [Li(G4)1][TFSA]. Figure 7 shows the temperature dependence of ionic conductivity and viscosity of [Li(G3)1][TFSA] and [Li(G4)1][TFSA]. The conductivities of [Li(G3 or G4)1][TFSA] increase with increasing temperature. Arrhenius plots of the conductivities showed convex-curved profiles; therefore, the experimental data were fitted with the VogelFulcherTamman (VFT) equation: B ð7Þ σ ¼ σ 0 exp T T0 where σ0 (S 3 cm1), B (kelvins), and T0 (kelvins) are constants. The best-fit parameters are summarized in Table 4. The Arrhenius plots of the viscosities showed concave-curved profiles and were also fitted with the VFT equation: B ð8Þ η ¼ η0 exp T T0
ð6Þ
where η0 (mPa 3 s), B (kelvins), and T0 (kelvins) are constants. The best-fit parameters of the viscosity are summarized in Table 5. The VFT-type parameters σ and η for [Li(G3)1][TFSA] and [Li(G4)1][TFSA] are similar, indicating that the temperature dependency of σ is dominated by that of η and is also similar to those observed for conventional RTILs.4144 Translational dynamics of the chemical species in the liquid may be correlated to the fluidity (1/η), and thus, the selfdiffusion coefficients, determined by PGSE-NMR, also showed VFT-type behaviors (Figure 8 and Table 6) as expected: B ð9Þ D ¼ D0 exp T T0
where T is the absolute temperature and a and b are constants. The molar concentration of [Li(G3)1][TFSA] is higher than that of [Li(G4)1][TFSA] (Table 1), resulting in a stronger Coulombic
where D0 (cm2 3 s1), B (kelvins), and T0 (kelvins) are adjustable parameters. In this temperature range, the diffusion coefficients of G3, [TFSA] anion, and Li+ cation in [Li(G3)1][TFSA]
F ¼ b aT
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Figure 8. Self-diffusion coefficients of glyme, Li+ cation, and [TFSA] anion of [Li(G3)1][TFSA] (left) and [Li(G4)1][TFSA] (right) as a function of temperature.
Table 6. VFT Equation Parameters of Self-Diffusion Coefficient Data D0, 105 cm2 3 s1
B , 102 K
T0 , K
9.1 ( 1.3
9.6 ( 1.1
166 ( 9.8
9.9 ( 1.3
10.1 ( 1.2
165 ( 9.3
7
12.7 ( 2.1
15.4 ( 0.9
127 ( 5.6
1
H (G4) 19 F (TFSA)
6.8 ( 1.5 5.2 ( 1.5
7.0 ( 1.6 6.8 ( 0.9
191 ( 16 189 ( 11
7
6.1 ( 1.4
6.8 ( 0.7
192 ( 7.6
[Li(G3)1][TFSA] 1
H (G3) F (TFSA)
19
Li (Li+)
[Li(G4)1][TFSA]
Li (Li+)
follow the sequence DG3 = DLi+ > D. In addition, the diffusion coefficients of G4, [TFSA] anion, and Li+ cation are almost the same (DG4 = DLi+ ≈ D), suggesting that the complex cation structures of [Li(G3)1]+ and [Li(G4)1]+ are retained even in this temperature range. The conductivities of [Li(G3)1][TFSA] and [Li(G4)1][TFSA] at 30 °C are 1.0 and 1.6 mS 3 cm1, respectively. The conductivity of [Li(G4)1][TFSA] is higher than that of [Li(G3)1][TFSA] in the whole temperature range, although the molar concentration (the number of carrier ions) is higher for [Li(G3)1][TFSA]. This is because the viscosity of [Li(G3)1][TFSA] is higher than that of [Li(G4)1][TFSA]. The higher concentration results in a stronger Coulombic interaction between ionic species in the liquid, leading to higher viscosity of [Li(G3)1][TFSA]. The higher diffusivity of Li+ and [TFSA] for [Li(G4)1][TFSA] than [Li(G3)1][TFSA] also supports the conductivity difference. Evaporation of Glyme from [Li(glyme)1][TFSA] and Its Correlation with Interfacial Electrochemical Reaction. When [Li(glyme)1][TFSA] is used as an electrolyte in lithium secondary batteries, a desolvation reaction of [Li(glyme)1]+ to Li+ occurs at the cathode and anode interfaces with discharging and charging reactions, respectively. The desolvation reaction appeared to be rather difficult because the Li+ ions strongly
coordinate to the glyme molecule, which might slow down the interfacial electrochemical reactions. In order to evaluate the activation energy for desolvation of Li+, the rate of weight loss of during the evaporation of glyme molecule from [Li(glyme)1][TFSA] was measured. The activation energy was compared with that of the interfacial electrochemical reaction at the metallic Li[Li(glyme)1][TFSA] interface in order to obtain insight into the electrochemical reactions. Figure 9 shows the time dependence of weight loss for [Li(glyme)1][TFSA] at different temperatures. When the temperature of [Li(glyme)1][TFSA] was kept at 100 °C, almost no weight loss was observed. It can be said that [Li(G3)1][TFSA] and [Li(G4)1][TFSA] behave like RTILs below 100 °C. At temperatures higher than 100 °C, gradual weight loss was observed and the rate of weight loss increased as the temperature increased. The weight change of [Li(glyme)1][TFSA] in the time scale shown in Figure 9 is almost linear against the measuring time. The evaporation rate of glyme from [Li(glyme)1][TFSA] was thus assumed to be a quasi-zero-order reaction as follows: ð10Þ
N ¼ kt þ N0
where N0 and N are moles of glyme in the liquid at the initial time and at time t, respectively, and k (moles per second) is the evaporation rate of glyme. The evaporation of glyme takes place at the liquidgas interface. Therefore, the evaporation rate k is dependent on the surface area of the liquid. The TG measurements were carried out in the same Al pans (inner diameter 5.0 mm) used in the present study. From the relationship between weight loss and time (Figure 9 and eq 10), the evaporation rates of glyme at various temperatures were estimated. Figure 10 shows Arrhenius plots of the evaporation rate of glyme from [Li(glyme)1][TFSA]. The activation energy for evaporation of G3 from [Li(G3)1][TFSA] was estimated to be 64 kJ 3 mol1 and that for G4 from [Li(G4)1][TFSA] was 61 kJ 3 mol1. The following two steps are assumed to occur at the liquidgas interface during evaporation of glyme from [Li(glyme)1][TFSA]: k1
½LiðglymeÞ1 ½TFSA sf glyme ðliquidÞ þ LiTFSA 18390
ð11Þ
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Figure 9. Relationship between weight loss and time for [Li(G3)1][TFSA] (left) and [Li(G4)1][TFSA] (right) at different temperatures.
Figure 10. Arrhenius plot of evaporative rate constant for [Li(G3)1][TFSA] and [Li(G4)1][TFSA]. k2
glyme ðliquidÞ sf glyme ðgasÞ
ð12Þ
From the TG results of mixtures of glyme and LiTFSA of x > 1 (Figure 5), it is clear that the rate of evaporation in eq 12 is faster than that of desolvation in eq 11. Thus, the desolvation step (eq 11) controls the evaporation rate of glyme, and the ratedetermining step of evaporation of glyme from the equimolar mixture [Li(glyme)1][TFSA] is presumed to be the desolvation process of [Li(glyme)1]+. Although G3 and G4 have four and five oxygen atoms, respectively, the activation energies for [Li(G3)1][TFSA] and [Li(G4)1][TFSA] are very close to each other. The similar activation energies for [Li(G3)1][TFSA] and [Li(G4)1][TFSA] indicate that the desolvation processes of [Li(G3)1 ]+ and [Li(G4)1]+ are controlled by the same factor. It has been reported that the first solvation energy of diethyl ether in the gas phase toward K+ is 93.3 kJ 3 mol1.60 Hydration energy to K+ in the gas phase has been reported to depend on the number of solvated water molecules.61 The first, second, third, and fourth solvation energies are 74.9, 67.4, 55.2, and 49.4 kJ 3 mol1, respectively.61 In the cases of G3 and G4, Li+ solvation occurs intramolecularly and a one-to-one complexation has been confirmed with LiTFSA.36,38
Figure 11. Temperature dependence of Nyquist plots for a symmetric [Li | [Li(G3 or G4)1][TFSA] | Li] cell. The interfacial process is represented as single-electrode impedance. See text for details.
On the contrary, the dissociation energy of the LiTFSA ion pair is calculated to be ca. 600 kJ 3 mol1.62 Consequently, it is difficult to consider that the activation energy in Figure 10 corresponds to the desolvation energy for a whole glyme molecule from Li+. It is more reasonable to assume that the desolvation process occurs step by step. The activation energy required to break the final Li+O iondipole interaction at the liquidgas interface appears to correspond to the activation energy observed in Figure 10. The electrochemical charge transfer process at the Li metal | [Li(glyme)1][TFSA] interface was investigated by means of complex impedance measurements. Glyme molecules have good reductive stability and the electrochemical deposition and stripping of Li can be performed reversibly in the glyme solution containing Li salt.36,37 A symmetric cell composed of [Li metal | [Li(glyme)1][TFSA] | Li metal] was fabricated to study the kinetics of the electrochemical reaction. Figure 11 shows Nyquist plots for the symmetric cell measured at various temperatures. A semicircle due to the charge-transfer process at the electrode | electrolyte interface was observed at each temperature and was stable against the duration of time. The interfacial process was 18391
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Figure 12. Nyquist plots for a symmetric [Li | [Li(G3)1][TFSA] | Li] cell. Inset: Schematic diagram of an equivalent circuit model used to simulate the impedance spectra.
Figure 13. Arrhenius plots for interfacial resistance for [Li(G3)1][TFSA] and [Li(G4)1][TFSA].
represented as single-electrode impedance. The charge-transfer resistance was evaluated by use of an equivalent circuit, as shown in Figure 12. Rsol is the solution resistance for ionic conduction between the anode and cathode, Rinterface is the interfacial resistance for the electrochemical reaction, and a constant phase element (CPE) is used instead of double-layer capacitance of the electrodes to analyze the Nyquist plots by nonlinear least-squares fitting. The charge transfer resistances of cathode (Rct1) and anode (Rct2) should be the same, because the surface area of the anode and cathode are the same in the symmetric cell. As temperature was increased, the interfacial resistance decreased because the charge-transfer process accelerated. Rinterface can be correlated with a kinetic constant for the charge-transfer reaction as follows: Rinterface ¼ Rct1 ¼ Rct2 ¼
RT i0 AF
ð13Þ
where A is the electrode area of each Li metal and i0 is the exchange current density for the electrochemical reaction at each Li electrode, which is in proportion to the kinetic constant for the electrochemical reaction at the interface. The Rinterface values at 30 °C for [Li(G3)1][TFSA] and [Li(G4)1][TFSA] were estimated to be 62 and 58 Ω 3 cm2, respectively. It should be noted that the Rinterface values lower than 100 Ω 3 cm2, even around room temperature, are comparable to or lower than those
reported for conventional organic electrolyte solutions,63 although the comparison of absolute Rinterface values between different systems is very hard to make because the Rinterface values are affected by surface roughness and passivation of lithium electrode surfaces. Figure 13 shows Arrhenius plots of (1/ Rinterface). The activation energy for the interfacial reaction in [Li(G3)1][TFSA] is 69 kJ 3 mol1, and that for [Li(G4)1][TFSA] is 67 kJ 3 mol1. The activation energies for [Li(G3)1][TFSA] and [Li(G4)1][TFSA] are close to each other. Thus, the charge-transfer rates in [Li(G3)1][TFSA] and [Li(G4)1][TFSA] are presumed to be controlled by the same activation process. The magnitude of the activation energy is very close to that observed for polymer electrolytes based on poly(ethylene oxide)64,65 and similar to that observed for conventional organic electrolyte solutions.66,67 It is also interesting to note that the activation energy for the interfacial electrochemical reactions is almost identical to that of the evaporation of glyme from [Li(glyme)1][TFSA], which strongly suggests that the desolvation process of Li+ cations at the interface of Li metal[Li(glyme)1][TFSA] is the rate-determining step. The glyme molecules are multidentate ligands and form one-to-one complexes with LiTFSA, which may imply a hard desolvation process. However, the absolute value and activation energy of the charge-transfer resistance at the interface are similar to those for conventional electrolyte solutions,62,65,66 and we did not observe large differences between them. The desolvation reactions of the glymes occur step by step, and the last desolvation process between Li+ and an ether oxygen dipole appears to be the activation barrier for the reactions.
’ CONCLUSION In this study, in order to understand a change from a dilute solution to a quasi-RTIL, transport and thermal properties were explored in detail by using a binary system consisting of LiTFSA and glymes (G3 and G4), covering a wide range of LiTFSA concentrations. The ionic conductivity of LiTFSA/glyme solutions changed depending on the LiTFSA concentration and passed through a maximum at ca. 1 mol 3 dm3. The viscosity of the solutions monotonically increased as the concentration of LiTFSA increased. The self-diffusion coefficient of each chemical species in the electrolyte solutions decreased as the concentration increased. DG in excess glyme solutions was higher than DLi+ and D. As the concentration of LiTFSA increased, the difference between DG and DLi+ became less pronounced. Identical values of DG and DLi+ in the equimolar mixture of glyme LiTFSA suggested the formation of a complex cation [Li(glyme)1]+ in the liquid. The ionicity Λimp/ΛNMR of [Li(glyme)1][TFSA] was higher than that of excess glyme solutions of [Li(glyme)x][TFSA] (x > 1). This result is similar to that observed for mixtures of typical RTILs and cosolvents and is opposite to that observed for conventional organic electrolyte solutions. The excess glyme solutions of [Li(glyme)x][TFSA] (x > 1) started to lose their weight at ca. 100 °C due to the evaporation of excess glyme. On the other hand, evaporation of glyme from the equimolar mixture [Li(glyme)1][TFSA] takes place at ca. 200 °C and [Li(glyme)1][TFSA] behaves like a quasi-RTIL consisting of [Li(glyme)1]+ and [TFSA]. A higher temperature is needed to evaporate Li+-coordinated glyme than that for free and noncoordinated glyme. The activation energy for desolvation of [Li(glyme)1]+ at the liquidgas interface was estimated to be ca. 63 kJ 3 mol1. Electrochemical impedance measurements revealed that absolute value and activation energy for the charge transfer resistance at the liquidLi metal interface was comparable 18392
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The Journal of Physical Chemistry C to those of conventional organic electrolyte solutions. The activation energy was ca. 68 kJ 3 mol1, which is in good agreement with that for the desolvation of [Li(glyme)1]+ at the liquidgas interface. It was concluded that the reaction barrier for the charge-transfer reaction of Li/Li+ at the electrode interface was the last desolvation process between Li+ and an etheroxygen dipole in the glyme molecules.
’ ASSOCIATED CONTENT
bS
Supporting Information. One figure, showing the transference number of Li+ at 30 °C as a function of LiTFSA concentration in G3 and G4 solutions. This material is available free of charge via the Internet at http://pubs.acs.org/.
’ AUTHOR INFORMATION Corresponding Author
*E-mail
[email protected].
’ ACKNOWLEDGMENT This study was supported in part by Technology Research Grant Program from the NEDO of Japan, by a Grant-in-Aid for Scientific Research from the MEXT of Japan in the priority area “Science of Ionic Liquids” (452/17073009), and by the ALCA program of Japan Science and Technology Agency (JST). ’ REFERENCES (1) Tarascon, J.-M.; Armand, M. Nature 2001, 414, 359. (2) Armand, M.; Tarascon, J.-M. Nature 2008, 451, 652. (3) Goodenough, J. B.; Kim, Y. Chem. Mater. 2009, 29, 587. (4) Xu, K. Chem. Rev. 2004, 104, 4303. (5) Welton., T. Chem. Rev. 1999, 99, 2071. (6) Seddon, K. R. Nat. Mater. 2003, 2, 363. (7) Wasserscheid, P.; Keim, W. Angew. Chem., Int. Ed. 2000, 39, 3772. (8) Noda, A.; Susan, M. A. B. H.; Kubo, K.; Mitsushima, S.; Hayamizu, K.; Watanabe, M. J. Phys. Chem. B 2003, 107, 4024. (9) Susan, M. A. B. H.; Noda, A.; Mitsushima, S.; Watanabe, M. Chem. Commun. 2003, 938. (10) Nakamoto, H.; Watanabe, M. Chem. Commun. 2007, 2539. (11) Lee, S.-Y.; Ogawa, A.; Kanno, M.; Nakamoto, H.; Yasuda, T.; Watanabe, M. J. Am. Chem. Soc. 2010, 132, 9764. (12) Garcia, B.; Lavallee, S.; Perron, G.; Michot, C.; Armand, M. Electrochim. Acta 2004, 49, 4583. (13) Matsumoto, H.; Sakaebe, H.; Tatsumi, K. J. Power Sources 2005, 146, 45. (14) Sakaebe, H.; Matsumoto, H.; Tatsumi, K. J. Power Sources 2005, 146, 693. (15) Seki, S.; Kobayashi, Y.; Miyashiro, H.; Ohno, Y.; Mita, Y.; Usami, A.; Terada, N.; Watanabe, M. Electrochem. Solid-State Lett. 2005, 8, A577. (16) Seki, S.; Kobayashi, Y.; Miyashiro, H.; Ohno, Y.; Usami, A.; Mita, Y.; Kihira, N.; Watanabe, M.; Terada, N. J. Phys. Chem. B 2006, 110, 10228. (17) Seki, S.; Kobayashi, Y.; Miyashiro, H.; Ohno, Y.; Usami, A.; Mita, Y.; Watanabe, M.; Terada, N. Chem. Commun. 2006, 5, 544. (18) Matsumoto, H.; Sakaebe, H.; Tatsumi, K.; Kikuta, M.; Ishiko, E.; Kono, M. J. Power Sources 2006, 160, 1308. (19) Seki, S.; Ohno, Y.; Kobayashi, Y.; Miyashiro, H.; Usami, A.; Mita, Y.; Tokuda, H.; Watanabe, M.; Hayamizu, K.; Tsuzuki, S.; Hattori, M.; Terada, N. J. Electrochem. Soc. 2007, 154, A173. (20) Fernicola, A.; Croce, F.; Scrosati, B.; Watanabe, T.; Ohno, H. J. Power Sources 2007, 174, 342.
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