Investigation of an Intermediate Temperature Molten Lithium Salt

Aug 16, 2013 - was the lowest among the lithium single molten salts composed of conventional ... The conventional lithium ion battery electrolyte is a...
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Investigation of an Intermediate Temperature Molten Lithium Salt Based on Fluorosulfonyl(trifluoromethylsulfonyl)amide as a SolventFree Lithium Battery Electrolyte Keigo Kubota and Hajime Matsumoto* Advanced Industrial Science and Technology (AIST) 1-8-31 Midorigaoka, Ikeda, Osaka 563-8577, Japan ABSTRACT: The asymmetric amide anion, such as (fluorosulfonyl)(trifluoromethylsulfonyl)amide (fTfN−), formed intermediate temperature molten lithium salts at 100 °C, which was the lowest among the lithium single molten salts composed of conventional halogenide, nitrate, perchlorate, and perfluoro anions. The molten Li[fTfN] was highly viscous (17 Pa·s at 110 °C, however, it was classified as a good ionic liquid by Walden plot analysis, which indicates the good-dissociation between Li+ and fTfN−. Due to the high transport number of the lithium cation (0.94), the Li[fTfN] behaves as a single lithium-ion conducting material. The Li[fTfN] possesses an electrochemical window (5.0 V) due to the oxidative resistance of the perfluoro anion. The charge and discharge for a conventional composite positive electrode containing LiCoO2 or LiFePO4 was successful. In particular, LiFePO4 showed a high capacity and Coulombic efficiency at 0.1−5 C.

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([(FSO2)2N]−, f2N−) and Tf2N−. Especially, Li[fTfN] exhibits an especially low melting point (100 °C) among the lithium salts composed of conventional anions and perfluoroanions. The Li[fTfN] does not melt at room temperature like the already reported other lithium ionic liquids composed of large borates and aluminates with an oligoether side chain6,7 and glyme−Li complexes;8 however, we selected this salt as a model of the pure lithium ionic liquids without any solvent molecule and also without a long alkyl and/or alkoxy moiety in the anionic species in order to focus on the effect of such a perfluorosulfonylamide anion on the physicochemical properties, such as viscosity, conductivity, and Li+ transport number, and on the electrochemical properties, such as the electrochemical window and Li+ insertion and extraction reaction on a conventional composite positive electrode without using any specific materials for the molten Li[fTfN]. To the best of our knowledge, there are few reports on lithium single ionic liquids and/or molten salt, which can be applied to modern 4 V class lithium secondary batteries at a high operating temperature except for binary or ternary alkali metal molten salts, such as the LiCl−KCl eutectic melt9 and Li[Tf2N]−K[Tf2N]−Cs[Tf2N] mixture.10 In contrast, the Li[fTfN] can be used as a single molten salt without other alkali metal salts due to its low melting point.

INTRODUCTION The conventional lithium ion battery electrolyte is a mixture of organic solvents, such as carbonates, and lithium salts with perfluoroanions, such as hexafluorophosphate, because their electrochemical stability permits the use of 4 V-class positive electrodes, such as LiCoO2 and a carbon-based negative electrode. However, the flammability of the organic solvents has been one of the critical issues when using these materials in a much larger battery system in an electric vehicle or stationary use.1 Ionic liquids (ILs) composed of the perfluoro anion, such as BF4− and bis(trifluoromethylsulfonyl)amide ([(CF3SO2)2N]−, Tf2N−), and onium cations, such as 1ethyl-3- methylimidazolium (C2mim+) and aliphatic quaternary ammonium, have been already significantly studied as one of the candidates for such less-flammable electrolytes. For example, aliphatic quaternary ammonium based ILs exhibited a large electrochemical window (over 5 V), which allows us to use a lithium metal anode without any additives and a high voltage positive electrode, such as LiCoO2.2 However, the onium cations in the ILs have no significant role in the charge and/or discharge process in a lithium battery system except for maintaining the ILs in the liquid state. Furthermore, the transport number of Li+ in the ILs is much lower than that in a conventional organic solvent electrolyte because the molar ratio of Li+ in the ILs is quite low.3 Recently, we reported that the (fluorosulfonyl)(trifluoromethylsulfonyl)amide anion ([(FSO2)(CF3SO2)N]−, fTfN−) has a superior ability to decrease melting points of not only its onium salts,4 but also its alkali metal salts5 compared for a symmetric amide, which means that the two contained side chains are the same, such as bis(fluorosulfonyl)amide © 2013 American Chemical Society

Received: May 22, 2013 Revised: July 31, 2013 Published: August 16, 2013 18829

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Figure 1. The thermal windows of lithium halides, nitrate, perchlorate, and perfluoro salts.

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versus Li+/Li). The cyclic voltammetry and charge−discharge tests were performed using the LiCoO 2 and LiFePO 4 composite electrodes. These electrodes (ϕ 16 mm) were composed of polyvinylidene fluoride (PVDF) as the binder and acetylene black (AB) as the conductive additive (LiCoO2:PVDF:AB = 86:7:7, LiFePO4:PVDF:AB = 84:8:8 in weight).

EXPERIMENTAL SECTION The Li[fTfN] was synthesized by cation-exchange from the K[fTfN]11 according to our previous study.12 The residual water was confirmed to be below 50 ppm by a Karl Fischer moisture meter (MITSUBISHI, CA-07). For the following measurements, Li[fTfN] was weighed in a dry chamber (DAIKIN, HRG-50AR, dew point < −50 °C). The melting point of the sample, sealed in an aluminum sealing pan in the dry chamber, was measured by a differential scanning calorimeter (Perkin-Elmer, Pyris 1) under flowing helium gas at a rate of 10 °C·min−1, equipped with a liquid nitrogen cooling system. The thermal decomposition temperature of the sample in a platinum pan was determined by a thermogravimeter (Seiko Instruments, TG/DTA-6200) under flowing nitrogen gas at the rate of 10 °C·min−1. The density was calculated by the volume of a weighed sample visually observed at constant temperature controlled by a thermostatic oil bath. The viscosity was measured by a viscoelastometer (TA Instruments, AR-G2) under flowing dry air. For the electrochemical measurements, in order to prevent absorption of moisture and oxygen, the molten salts impregnated in a glass separator (Nippon Sheet Glass, TGP008F, 80 μm thickness, ϕ20 mm) and each electrode was sealed in a stainless-steel holder (Hohsen Corp., HS-2 test cell) in an argon filled glovebox (MIWA MFG, MDP-1.58-T1000, gas recycling purification system: MP-H60W, [H2O], [O2]: 5 ppm >). Measurements were performed by an electrochemical analyzer (Bio-Logic, VMP3). The specific conductivity was measured by a disk-shaped nickel (diameter = 18 mm) symmetrical cell. The transport number of the lithium cation was measured by a disk-shaped lithium (diameter = 18 mm) as the working electrode, instead of nickel. Cyclic voltammetry was done using the stainless-steel cell holder with the disk-shaped lithium as both the counter electrode and reference electrode. The electrochemical window was measured using a nickel working electrode in the negative potential region (−0.2∼+2.6 versus Li+/Li) or a platinum working electrode in the positive potential region (+3.2∼+5.7

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RESULTS AND DISCUSSION Thermal Window. Application of a molten salt is mostly affected by the temperature region of the molten salt maintained in the liquid phase, which is termed the “thermal window” in this work. The lower limit is simply the melting point (Tm) measured by differential scanning calorimetry (DSC), and the upper limit is usually estimated by thermal gravimetry (TG). Figure 1 shows the thermal windows of lithium single salts composed of not only perfluoroanions, but also non-fluorine containing conventional anions such as halides, nitrate, and perchlorate. The lithium halides exhibit a large thermal window because they are composed of a monoatomic ion, which is never thermally decomposed.13 Such thermally and chemically stable salts were already investigated as an electrolyte for a “thermal battery”, which was operated at over 350 °C using a LiCl−KCl eutectic mixture, not a single salt.9 However, the output voltage of the thermal batteries was about 2 V due to the low oxidation potential of the halides, which means that the electrochemical window of these melts were quite narrow for use with modern positive electrodes, such as LiCoO2. LiNO3 and LiClO414 have been studied as substitutes for the halides due to their lower melting points than the halides. However, the lithium secondary battery using the LiNO3 or LiClO4 does not appear practical due to the exothermic reaction with the Li alloy negative electrode.15 In order to consider the use of a high voltage positive electrode, the perfluoro anion, such as PF6− and Tf2N−, seems a much more reasonable choice considering their higher oxidation potential than non-fluorine containing anions, such as halides, nitrate, and perchlorate. Although LiPF6 has been used as a supporting electrolyte dissolved in a 18830

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viscosity of the low melting lithium perfluorosulfonylamides, such as Li[f2N], Li[fTfN], and Li[Tf2N],21 could also be described by the Arrhenius equation. However, they are different from the conventional salts, that is, their values are 102∼103 orders higher than the conventional salts, and their Eη values, which indicate gradients in the plots in Figure 3, are several orders higher (Table 1). On the other hand, the molar

commercially available organic solvent for a lithium ion battery, not only LiPF6, but also LiAsF6, LiCF3SO3, and LiBF4, which are also used as typical supporting electrolytes, have no thermal window, that is, they thermally decompose before melting.16−18 As shown in the lower half of Figure 1, the lithium perfluorosulfonylamides have wide thermal windows, from clear melting behavior during the DSC measurement5,19,20 and also visual observation at much lower than the Td of over 300 °C, which was determined by the 5 % loss in weight in the TG measurement (Figure 2). In particular, Li[fTfN] has the widest

Table 1. Melting Point (Tm), the Activation Energy for the Viscous Behavior (Eη) and Ionic Conductance (EΛ) for Li[fTfN] and Various Lithium Salts LiF LiCl LiBr Lil LiNO3 LiClO4 Li[f2N] Li[fTfN] Li[Tf2N]

Figure 2. TG curves of Li[f2N], Li[fTfN], and Li[Tf2N] at the rate of 10 °C·min−1 under flowing N2 gas.

Eη (kJ mol−1)

EΛ (kJ mol−1)

845 610 550 449 254 236 140 100 233

21.8 17.2 17.2 17.4 18.6 19.5 61.6 78.9 73.3

5.39 6.10 6.97 5.90 14.1 13.0 46.4 64.6 53.4

conductivity of the conventional lithium salts are also expressed by the Arrhenius equations, as described by the following: ⎛ −E ⎞ Λ = Λ 0 exp⎜ Λ ⎟ ⎝ RT ⎠

thermal window, over 200 °C from the lowest Tm of 100 °C, among the perfluorosulfonylamides. As stated later, the Li[fTfN] also possesses the wide electrochemical window of 5.0 V, allowing the charge-discharge of 4 V class electrodes, such as LiCoO2. Therefore, the Li[fTfN] can be the first single lithium molten salt containing no solvent, no complex, and no other alkali metal salt. Viscosity and Molar Conductivity. Viscosities of the lithium salts appear to be essentially governed by the anion species. The viscosities of conventional lithium salts, such as halides, nitrate, and perchlorate, range from 100∼101 mPa·s

(2)

where Λ0 is the limiting conductivity and EΛ is the activation energy. In contrast to the viscosity, the conventional lithium salts possess the high value of 101∼102 S·cm2·mol−1,13 which is one of the reasons for their use as electrolytes in thermal batteries,15 and the perfluorosulfonylamides have 102∼103 lower values than the conventional salts (Figure 4). Similar to Eη, the EΛ values of the perfluorosulfonylamides are several orders higher than the conventional ones (Table 1).

Figure 4. Arrhenius plots of molar conductivity for Li[fTfN] and various lithium salts.

Figure 3. Arrhenius plots of viscosity for Li[fTfN] and various lithium salts.

The relationship between the viscosity and conductivity can be considered in terms of a Walden plot, as described by Angell et al.22 In this work, adjusted Walden plot, in which the differences in the ion size were considered,23 was applied to the lithium molten salts containing the small Li+. As shown in the Walden plot of the lithium salts (Figure 5), the conventional lithium salts are located near the ideal line (1 mol·dm−3 of KCl aqueous solution), which is called a “good ionic liquid”. This is

(Figure 3), which is expressed by the well-known Arrhenius equation,13 as described by the following: ⎛ Eη ⎞ η = η0 exp⎜ ⎟ ⎝ RT ⎠

Tm (°C)

(1)

where η0 is the limiting viscosity, Eη is the activation energy, and R is the gas constant (8.314 J·K−1·mol−1). In this study, the 18831

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Figure 6. Time dependence of polarization currents at the constant bias of 10 mV in Li | Li[fTfN] | Li at 140 °C (solid line) and Li | LiPF6 in EC-DMC (1:1 in mol) | Li at 45 °C (dotted line). Figure 5. Adjusted Walden plots of temperature-dependent molar conductivities and viscosities for Li[fTfN] and various lithium salts (r+: cation size, r−: anion size).

explained by the fact that the small Li+ cations can easily move among the large counteranions. Alternatively, Li[Tf2N] is apparently located below the ideal line because Li+ are well associated with Tf2N−, and large anion clusters, such as Li[Tf2N]2−, are formed.24 However, the Li[f2N] and Li[fTfN] are located above the ideal line though they are perfluorosulfonylamides as the Li[Tf2N]. The f2N− and fTfN− are well dissociated with Li+, because their structures, which contain the FSO2 side chain, may reduce the interaction with Li+. For the f2N−, which forms anion clusters such as Li[f2N]32−,25 the interaction with Li+ is lower than the Tf2N−.26,27 Therefore, the exchange of CF3SO2 side chain of the Li[Tf2N] to FSO2 of the Li[f2N] and Li[fTfN] would reduce the interaction with Li+. Li Ion Transport in Li Symmetrical Cell. A lithium symmetrical cell has been used in order to investigate transport behavior of Li+ in the electrolyte and lithium-electrolyte interface. The electrochemical measurements using the Li symmetrical cell have been applied to organic solvents,28 RTIL,3,29 and polymer electrolytes.30,31 However, for the single lithium salt, the Li symmetrical cell could not be used because their Tm were too high to combine the lithium metal. In this study, the dc polarization and ac impedance measurement were performed for the Li[fTfN] by the Li symmetrical cell due to its lower Tm (100 °C) compared to lithium metal (180 °C). To compare the electrochemical response of the Li[fTfN] with that for a conventional organic electrolyte, one of the typical organic electrolytes, such as ethylene carbonate (EC) and dimethyl carbonate (DMC) mixed solvent at a 1:1 molar ratio containing 1 mol dm−3 LiPF6, was examined in the same measurements. For the organic solvent electrolyte, there was a large difference between the initial current (I0) and stable current (IS) (Figure 6) because the I0 was contributed by both Li+ and PF6−, whereas IS was contributed only by Li+. Contrary to the results for the organic solvent electrolyte, no significant difference between I0 and IS was observed for Li[fTfN] under the same experimental conditions as well as the single-cationconductive electrolyte.31,32 An apparent difference between the Li[fTfN] and organic solvent electrolyte was observed in the ac impedance (Figure 7). The solution impedance (Rs) of the Li[fTfN] was much higher than the organic solvent electrolyte, however, the diffusion impedance Zd(0) was smaller. From the results of the ac impedance, the transport number of Li+ (tLi+), which is an essential parameter for the electrolytes, was estimated using the following equation:33,34

Figure 7. Cole−Cole plots for Li | Li[fTfN] | Li at 140 °C (solid line) and Li | LiPF6 in EC-DMC (1:1 in mol) | Li at 45 °C (dotted line).

t Li + =

Rs R s + Zd(0)

(3)

For the organic solvent electrolyte, the tLi+ was 0.34 (eq 3), which was similar to tLi+ of other organic solvent electrolytes measured by NMR35 and the Hittorf method.36 Contrary to the results for the organic electrolyte, a very high value of 0.94 was calculated for Li[fTfN]. The high value of Li[fTfN] (0.94) is much higher than those for an organic solvent28 and RTIL3,29 electrolytes (below 0.6). Furthermore, among the single lithium salts (LiF: 0.55,37 LiNO3: 0.7338), which were determined by a molecular dynamics calculation, Li[fTfN] also has the highest tLi+. This indicates that the mobility of the counteranions become low with an increase in their size, and thus the mobility of Li+ becomes relatively high. The tLi+ was also estimated by the ac-dc complex method developed by Bruce et al,39 which have been used for organic solvents,28 RTIL,3,29 and polymer electrolytes30,31 according to the following equation: t Li + =

I S(E − I 0R ct0) I 0(E − I SR ctS)

(4)

where Rct is the charge transfer resistance between the Li[fTfN] and the lithium metal in the ac impedance spectra (Figure 7) and the superscripts of 0 and S refer to the initial and steady states, respectively. For the organic solvent electrolyte, the same value (0.34) was obtained. To apply the obtained results to eq 4, the same high value of 0.94 was also calculated for the Li[fTfN] melt. Strictly speaking, the ac-dc complex method, in which migration currents by potential difference and diffusion ones by concentration gradient are considered, seems to be not 18832

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electrode.40 On the oxidative side (above 3.2 V versus Li+/ Li), a current on the platinum electrode was observed at 5.0 V versus Li+/Li, which was determined when the current density exceeded 0.1 mA·cm−2. This reaction is considered to be the oxidative decomposition of the fTfN anion because the Li[fTfN] contains only Li+ and fTfN−. Table 2 summarizes the electrochemical windows for the Li[fTfN] and alkali metal molten salts. In the mixture containing different anions,41−43 the anode limit was determined by oxidation of the easiest anion. The order of the oxidation potential of the lithium halides, which were calculated from thermodynamic data, was reported to be as follows: F > Cl > Br > I.15 The LiNO3−KNO3 showed wider electrochemical window at the lower operating temperature than the halides.44 However, the LiNO3−KNO3 generates exothermic reactions with the Li alloy negative electrodes, which lead to fire disasters.15 The Li[Tf2N] mixtures have wide electrochemical windows of over 5.0 V due to the high oxidation resistance of the Tf2N−.21 The Li[Tf2N]−Cs[Tf2N] is slightly more stable than the Li[Tf2N]−K[Tf2N] due to its lower operating temperature. In the electrochemical experiments in a previous study, the lithium salts were used as mixtures because of their high melting temperatures. Meanwhile, the Li[fTfN] shows the wide electrochemical window of 5.0 V at a temperature equal to the Li[Tf2N] mixtures, although it is a single salt. Charge−Discharge Property. As already stated, due to the thermal window of 100−300 °C and the electrochemical window of 5.0 V, Li[fTfN] has the possibility of combining with a 4 V class positive electrode for a lithium secondary battery. In this work, the compatibility of LiCoO2 and LiFePO4 with the Li[fTfN] has been investigated by cyclic voltammetry using Li | Li[fTfN] | LiCoO2 and Li | Li[fTfN] | LiFePO4 cells, respectively (Figure 9). A pair of oxidation and reduction

applicable to Li[fTfN] because no concentration gradient existed in the single lithium molten salt; however, the same high value was derived not only by the ac-dc complex method, but also by the ac impedance measurement at low frequency, thus, we would like to emphasize here that the Li+ might be a dominant mobile species in the Li[fTfN] and also such a singlecation-conducting nature of Li[fTfN] melt might not be observed in all of the single lithium salts. We are investigating the transport property of the Li[fTfN]−Cs[fTfN] mixtures, which is submitted in the next work, for example, a mixture at high concentration of Li[fTfN], such as Li0.9Cs0.1[fTfN], was found to have high tLi+ (0.80) by the ac-dc complex method, which is near the value of the single Li[fTfN]. It will be necessary to measure diffusion coefficients for the Li[fTfN] and its mixtures, in order to prove the dominant mobility of the Li+. This measurement using a high temperature nuclear magnetic resonance (NMR) is now underway. Electrochemical Window. The electrochemical window, which fixes the limits of the cathode and anode potentials, is one of the main features determining the compatibility of the electrode materials. The cathode and anode limit of the Li[fTfN] has been investigated by cyclic voltammetry (Figure 8). In the negative region (below +2.6 V versus Li/Li+), a pair

Figure 8. A combined cyclic voltammogram of Li[fTfN] with a nickel electrode in the negative region (−0.2∼+2.6 V versus Li+/Li) or a platinum electrode in the positive region (+3.2∼+5.7 V versus Li+/Li) at a scan rate of 10 mV·s−1 and 150 °C.

of reduction and oxidation currents on the nickel electrode was observed at 0 V versus Li+/Li. As for our previous study of Li[Tf2N] mixtures, their reactions were interpreted as the deposition and redissolution of lithium metal.21 Low reduction and oxidation peaks were also observed at +1.5 and +2.0 V versus Li+/Li, respectively. These peaks may be attributed to the under-potential deposition of lithium on the nickel

Figure 9. Cyclic voltammograms in Li | Li[fTfN] | LiCoO2 and Li | Li[fTfN] | LiFePO4 cell at the scan rate of 0.1 mV·s−1 and 110 °C.

Table 2. Electrochemical Windows (E.W.) for Alkali Metal Molten Salt Containing Lithium Salts from the Reduction of Li+ in Each Anodic Reaction, and Their Operating Temperatures (T) molten salt

E. W. (V)

T (°C)

LiCl−KCl41 LiCl−KCl−CsCl42 LiBr−KBr−CsBr43 LiF−LiCl−CsI41 LiNO3−KNO344 Li[fTfN] (this study) Li[Tf2N]−K[Tf2N]21 Li[Tf2N]−Cs[Tf2N]21

3.6 3.8 3.4 2.7 4.5 5.0 5.0 5.2

450 300 250 425 180 150 170 130

method CV CV CV CV CV CV CV CV

at at at at at at at at

250 mV s−1 W.E.: tungsten, C.E.: tungsten, R.E.: Li−S alloy 10 mV s−1 W.E.: nickel or GC, C.E.: aluminum or GC, R.E.: Li−Al alloy 10 mV s−1 W.E.: nickel or GC, C.E.: aluminum or GC, R.E.: Li−Al alloy 250 mV s−1 W.E.: tungsten, C.E.: tungsten, R.E.: Li−S alloy 50 mV s−1 W.E.: silver, C.E.: platinum,R.E.: silver l0 mV s−1 W.E.: nickel or platinum, C.E.: lithium, R.E.: lithium l0 mV s−1 W.E.: nickel or GC, C.E.: lithium or GC, R.E.: lithium l0 mV s−1 W.E.: nickel or GC, C.E.: lithium or GC, R.E.: lithium

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currents corresponding to the deintercalation and intercalation of Li+ by each electrode was observed, and no other currents due to the chemical reactions of the Li[fTfN] and the electrodes were observed. Thus, the Li[fTfN] molten salt can be used as the electrolyte of a lithium secondary battery in combination with the LiCoO2 and LiFePO4 composite electrodes. The lithium half cells have been evaluated at the rate of 0.1 C, which means 0.023 mA·cm−2 for the LiCoO2 and 0.015 mA· cm−2 for the LiFePO4. The first and second discharge capacities of the LiCoO2 were 114 mAh·g−1 and 108 mAh·g−1, which showed an apparent decrease, and their coulomb efficiencies are 84.4% and 94.6% (Figure 10). This is due to the thermal Figure 12. Rate performance at 0.1−5 C in Li | Li[fTfN] | LiFePO4 at 110 °C and Li | LiPF6 in EC-DMC (1:1 in mol) | LiFePO4 at 25 °C.

Li | charged LiFePO4 cell (Figure 13), the bulk resistance of the Li[fTfN] is greater than that of the LiPF6 in EC-DMC because

Figure 10. A charge−discharge curve in Li | Li[fTfN] | LiCoO2 cell at the rate of 0.1 C (0.023 mA·cm−2) between +3.2∼+4.2 V versus Li+/Li at 110 °C. Figure 13. Cole−Cole plot of Li | Li[fTfN] | charged LiFePO4 at 110 °C and Li | LiPF6 in EC-DMC (1:1 in mol) | charged LiFePO4 at 25 °C.

brittleness of the constitutional material, such as PVDF or the charged LiCoO2. However, the LiFePO4 showed a reversible charge and discharge with a higher capacity and coulomb efficiency (152 mAh·g−1, 99.9%) (Figure 11). There was little

of the low conductivity, but its charge transfer resistance, which is composed of the Li | Li[fTfN] and Li[fTfN] | charged LiFePO4 interfaces, is lower due to the high temperature. This result is considered to be due to the high Li+ diffusion rate in the LiFePO4 particles at high temperature.45 A detailed discussion, such as dividing the Li | Li[fTfN] and Li[fTfN] | charged electrodes and comparing them with the other electrolytes, such as ionic liquids at the same temperature, are currently underway.

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CONCLUSIONS Li[fTfN] can be used as a single lithium molten salt electrolyte for a lithium secondary battery, due to its thermal window of 100−200 °C and electrochemical window of 5.0 V. The molten Li[fTfN] was highly viscous (17 Pa·s at 110 °C); however, it was categorized as a good ionic liquid and its Li+ transport number is close to unity (0.94). The charge and discharge with the existing composite electrodes, such as LiCoO2 and LiFePO4, was successful in Li[fTfN]. In particular, LiFePO4 showed a nearly ideal capacity and Coulombic efficiency at 0.1 C, and its rate performance at 0.1−5 C exceeded that using an organic electrolyte at room temperature. Although only the single Li[fTfN] was reported in this work, its melting point and transport properties would be improved by additive materials, such as the other alkali metal fTfN salts. The improved Li[fTfN] mixtures have the possibility to be molten salt electrolytes operating at room temperature. Studies

Figure 11. A charge−discharge curve of Li | Li[fTfN] | LiFePO4 cell at the rate of 0.1 C (0.015 mA·cm−2) between +2.5∼+4.2 V versus Li+/Li at 110 °C.

irreversible capacity between the first and second charges. Moreover, the LiFePO4 maintained a high capacity at the high rate of 0.5−5 C (0.075−0.75 mA·cm−2) (Figure 12). When compared to an organic solvent electrolyte under the same conditions at 25 °C, which was 1 mol·dm−3 of LiPF6 added to a mixed organic solvent of ethylene carbonate (EC) and dimethyl carbonate (DMC) at a 1:1 molar ratio, the Li | Li[fTfN] | LiFePO4 possesses a better rate performance in spite of its much higher viscosity and low conductivity (17 Pa·s and 0.11 mS·cm−1 at 110 °C). As shown in the Cole−Cole plot for the 18834

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of the alkali metal fTfN mixtures containing Li[fTfN] are now underway.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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REFERENCES

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