Anionic Effects on Solvate Ionic Liquid Electrolytes in Rechargeable

The solvate IL [Li(G4)][BETI] was found to be electrochemically stable in the Li–S cell and allowed a stable operation with a capacity of 600–700 ...
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Anionic Effects on Solvate Ionic Liquid Electrolytes in Rechargeable Lithium−Sulfur Batteries Kazuhide Ueno, Jun-Woo Park, Azusa Yamazaki, Toshihiko Mandai, Naoki Tachikawa, Kaoru Dokko, and Masayoshi Watanabe* Department of Chemistry and Biotechnology, Yokohama National University, 79-5 Tokiwadai, Hodogaya-ku, Yokohama 240-8501, Japan S Supporting Information *

ABSTRACT: A series of equimolar mixtures of Li salts (LiX) and glymes (triglyme (G3) and tetraglyme (G4)), [Li(glyme)]X with different anions (X: [N(SO2C2F5)2] = [BETI]; [N(SO2CF3)2] = [TFSA]; [CF3SO3] = [OTf]; BF4; NO3), were used as electrolytes to study the anionic effects of [Li(glyme)]X on the performance of lithium−sulfur (Li−S) batteries. The dissolution of lithium polysulfides (Li2Sm), which are discharge products of elemental sulfur, was significantly suppressed in the solvate ionic liquid (IL) electrolytes, as seen in [Li(G4)][BETI] and [Li(glyme)][TFSA], wherein all of the glymes participated in the formation of the complex cation [Li(glyme)]+. It was found that NO3 anions were irreversibly reduced at the composite cathode during discharge and BF4 anions formed unexpected byproducts through a chemical reaction with the polysulfide anions. Successful charge/discharge of Li−S cell could not be performed in [Li(glyme)]X in the presence of these anions because of the undesired side reactions. The solvate IL [Li(G4)][BETI] was found to be electrochemically stable in the Li−S cell and allowed a stable operation with a capacity of 600−700 mAh·g−1 and a Coulombic efficiency of 98.5% over 100 cycles, similar to that achieved by [Li(glyme)][TFSA]. In contrast, the Li−S cell with a concentrated electrolyte solution, [Li(G3)][OTf], showed a much lower capacity and Coulombic efficiency.

1. INTRODUCTION Rechargeable lithium−sulfur (Li−S) batteries have attracted much attention lately owing to their greater theoretical capacity of 1672 mAh·g−1 for elemental sulfur (S8) cathodes, compared to that of cathode materials in the well-established Li-ion batteries. Also, a Li−S battery is a promising candidate for lowcost, environmentally benign, and high-density energy storage using abundantly available materials.1−3 Although the insulating property of the sulfur cathode and the dissolution of the reaction intermediates (lithium polysulfides, Li2Sm) into the electrolyte have long been recognized as two major challenges, poor current collection and the resulting low utilization of sulfur have been greatly improved by recent efforts to fabricate composite sulfur cathodes with nanostructured carbon materials2,4−7 and conducting polymers.8,9 However, the critical problem of Li2Sm dissolution still remains unsolved. Li2Sm dissolution causes undesired side reactions as well as a loss of the active materials from the cathode. The solubilized Li2Sm can diffuse through the electrolyte and be reduced at the Li metal anode, and the resulting further-reduced intermediates can be oxidized again at the cathode during cell charging. This is known as a “redox shuttle mechanism,”10−12 which leads to a low Coulombic efficiency and rapid fading of the capacity. Furthermore, nonaqueous liquid electrolytes in common Li-ion batteries, which contain carbonate solvents such as ethylene carbonate (EC) and diethylcarbonate (DEC), cannot be used in Li−S batteries because the polysulfide anions S m 2− irreversibly react with the carbonates.13 These problems closely © XXXX American Chemical Society

depend on the design of electrolytes; in other words, the electrolyte plays a critical role in improving the performance of Li−S batteries. The most commonly used electrolytes for Li−S batteries are mixed ether electrolytes such as dimethoxyethane (DME) and 1,3-dioxolane (DOL) containing approximately 1 mol·dm−3 of a dissociative lithium salt, typically lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA]),14 and LiNO3 as an additive.15 However, this type of electrolyte containing excess solvents readily solubilizes the Li2Sm species (especially, m ≥ 4) and in fact functions as a catholyte after the first discharge in the Li−S cells.16 The redox shuttle can be prevented by an effective passivation layer on the Li metal anode, which is formed by the reaction of the Li metal surface with the electrolyte mixture in the presence of LiNO3.15 Hence, the inhibition of the redox shuttle depends on the stability of the passivation layer during repeated charge/discharge cycles. The design of electrolytes with a potentially lower solubility of Li2Sm would rather be more straightforward. To this end, the solvation power toward the Li+ cation in Li2Sm (i.e., donor property), which primarily dominates the solubility in aprotic solvents, must be minimized while ensuring adequate Li+ conduction through the originally dissolved Li salts in the electrolytes. Unlike the previously reported strategies using Received: July 18, 2013 Revised: September 14, 2013

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polymer electrolytes17 and inorganic solid electrolytes18,19 that offered physical and kinetic barriers to Li2Sm dissolution, there is an alternate simple route to address the above requirements by suitably designing the liquid electrolytes using the unique solvating property of ionic liquids (ILs).20,21 Because typical ILs consist of charge-delocalized bulky ions with no solvent, their solvation power for the ionic species of Li2Sm is expected to be lower than that of organic electrolytes. For Li−S batteries, we focused on a new family of ILs, “solvate” ILs,22 instead of binary mixtures of aprotic ILs and Li salts,23 which are commonly employed for the IL-based Li battery electrolytes.24 Equimolar mixtures of Li[TFSA] and glymes (abbreviated as [Li(glyme)][TFSA]) are the representative of lithium solvate ILs, wherein the long-lived complex cation [Li(glyme)]+ is formed and functions as a weakly coordinating cation, like the one found in common aprotic ILs.25−27 [Li(glyme)][TFSA] showed a faster transport of Li+ ions than the binary systems of aprotic ILs and Li salts,28 while its donor property was comparable to that of the aprotic ILs.29 The suppression of Li2Sm dissolution was greater in the solvate ILs compared to that in organic electrolytes. As a result, a stable charge/discharge of Li−S batteries with high capacity and high Coulombic efficiency was achieved for over 400 cycles using lithium solvate ILs as electrolytes.30 As in the case of aprotic ILs,31 the ionic structure of the solvate ILs, especially the structure of anions, should have a major impact on their electrolyte properties. Herein, a series of equimolar mixtures with different lithium salts (LiX) and glymes ([Li(glyme)]X) were investigated as potential electrolytes for Li−S batteries. This was motivated by our previous finding that [Li(glyme)]X could be divided into two classes of liquid states, that is, ordinary concentrated solutions and the solvate ILs, depending on the anionic structure,29 and by expectation that the battery performance is also affected by factors such as the class of liquid states and intrinsic electrochemical characters of the anions. The solubility limits of the active materials such as S8 and Li2Sm in [Li(glyme)]X with different anions, including nitrate (NO 3 − ), trifluoromethanesulfonate ([OTf]−), tetrafluoroborate (BF4−), and bis(pentafluoroethenesulfonyl)amide ([BETI]−), were quantitatively studied, and their effect on the battery performance of Li−S cells was explored in detail. A comparison with previously reported data on [Li(glyme)][TFSA]30 was also performed.

3:8, and 1:8 resulting in nominal chemical structures Li2S8, Li2S4, and Li2S2, respectively. The solutions were stirred and heated at 60 °C for 5 days and then maintained for 2 days at room temperature. Each supernatant solution was then diluted with 1 mol·dm−3 Li[TFSA] in either a G3 or G4 solution. The diluted Li2Sm solutions were electrochemically oxidized to S8 in an H-type two-electrode cell. The maximum absorption of S8 oxidized from the polysulfides was recorded using a UV−vis spectrophotometer (UV-2500PC, Shimadzu, Kyoto, Japan). The sulfur composite cathode consisted of elemental sulfur (S8, 99.99%, Wako, Osaka, Japan), Ketjenblack (KB, EC600JD, Lion Corporation, Tokyo, Japan), and poly(vinyl alcohol) (PVA, saponification degree 86−90 mol %, average degree of polymerization 3100−3900, Wako) with a weight ratio of 6:3:1. Before preparing the cathode, KB and S8 were thoroughly mixed using an agitating mortar and then heated to 155 °C for 8 h in a sealed bottle to form a S8/KB composite.2 The S8/KB composite and PVA were thoroughly dispersed in Nmethylpyrrolidone (NMP) to form an ink, which was then spread on an Al foil using a bar-coater, followed by the evaporation of the solvent and pressing it at a pressure of 100 kgf·cm−2. The coin-type cell was assembled in an Ar-filled glovebox using the S8/KB composite cathode (electrode area: ca. 2 cm2, weight of sulfur on the composite electrode: ca. 1.0 mg), a porous glass separator (GA-55, Advantec, Ehime, Japan), a Li metal anode, and an electrolyte. The galvanostatic charge/ discharge measurements (HJ1001SD8, Hokuto Denko, Tokyo, Japan) were performed with cutoff potentials of 1.5 and 3.3 V for the discharge and charge steps, respectively, at 30 °C. Because the Li−S cell was prepared in the fully charged state, the charge/discharge cycle was defined as follows: first discharge → second charge → second discharge → third charge → third discharge, and so on. The Coulombic efficiency was defined as: Nth discharge capacity/Nth charge capacity. The specific capacity of the cell was calculated based on the mass of S8. To study the decomposition of the electrolytes in the presence of Li2Sm, multinuclear NMR measurements were performed using a JEOL ECX400 spectrometer. A sample tube with a double tube structure (SC-008, Shigemi, Tokyo, Japan) was used for the measurements.

3. RESULTS AND DISCUSSION Transport Properties and Solubility of S8 and Li2Sm in [Li(glyme)]X. The transport properties of [Li(G3)]NO3, [Li(G3)][OTf], [Li(G4)]BF4, [Li(G3 and G4)][TFSA], and [Li(G4)][BETI] have been reported earlier;29 we revisited them to reveal their correlation with the solubility of sulfur derivatives. These equimolar mixtures did not readily crystallize and maintained in the liquid state (or supercooled liquids), even though the Tm values of certain complexes were slightly higher than the room temperature.29 Table 1 summarizes the transport properties such as conductivity (σ), viscosity (η), molar concentration (c), and the ratio of self-diffusion coefficients of glyme and lithium (DG/DLi). As reported earlier,29 [Li(glyme)]X can be classified as either a concentrated solution or a solvate IL, depending on the formation of the stable [Li(glyme)]+ cation. The ratio DG/DLi is a useful indicator of the long-lived [Li(glyme)]+, because the diffusion measurements using pulsed-field-gradient NMR technique were made on the 10−2 s time scale (Table 1).27 For [Li(G3 and G4)][TFSA] and [Li(G4)][BETI], DG/DLi is almost unity; that

2. EXPERIMENTAL SECTION Triglyme (G3) and tetraglyme (G4) were obtained from TCI and were distilled under high vacuum in the presence of sodium metal; the water content was maintained below 50 ppm. Battery-grade Li salts Li[BETI], Li[OTf], and LiBF4 were purchased from Kishida Chemical (Osaka, Japan) and used as received. LiNO3 was dried under high vacuum at 100 °C prior to use. Glyme−lithium equimolar mixtures were prepared by mixing stoichiometric amounts of lithium salt and glyme in an Ar-filled glovebox (VAC, dew point < −80 °C). Typically, the lithium salt and glyme were mixed and stirred magnetically in a vial at 60 °C overnight. When necessary, the samples were further heated at 80 °C for several hours to ensure complete mixing. The saturation solubilities of S8 and Li2Sm in [Li(glyme)]X were determined according to a procedure described elsewhere.32 The saturated solutions of Li2Sm were prepared using stoichiometric amounts of S8 and Li2S with the molar ratios 7:8, B

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(pure S8) and final (Li2S) species in the discharge reaction of sulfur is less than 10 mM, independent of their anionic structures in [Li(glyme)]X, except for Li2S in [Li(G3)]NO3. However, the solubility of the intermediate species, L2Sm (8 ≥ m ≥ 2), is much higher than that of S8 and Li2S, and greatly depends on the ionic structure of the electrolyte. The solubility limits for [Li(G3)]NO3 and [Li(G3)][OTf] are 1−2 orders of magnitude higher than those for [Li(G3 and G4)][TFSA] and [Li(G4)][BETI]. For [Li(G4)]BF4, the solubility of Li2Sm could not be determined because of an unexpected chemical reaction between Li2Sm and [Li(G4)]BF4; thus, only the data on S8 is shown in Figure 1. This will be discussed later in this paper. Previous works35,36 have reported that the higher the concentration of Li salt, the lower is the solubility of Li2Sm in the concentrated electrolytes. However, [Li(G3)]NO3 and [Li(G3)][OTf] have higher Li salt concentrations (c) than the others (Table 1) but nonetheless show a much higher solubility of Li2Sm. For example, in [Li(G3)]NO3 having an extremely high concentration (c = 4.77 M), the solubility of Li2S8 exceeded 4000 mM in atomic S concentration, which is almost comparable level to that for the more dilute electrolyte of 0.98 M Li[TFSA] in G4.21 Therefore, the Li salt concentration is not only the factor affecting the solubility of Li2Sm. The liquid states of [Li(glyme)]X would rather dominate Li 2 S m dissolution. It can be seen that [Li(glyme)]X can be classified into two groups based on the solubility of Li2Sm: less soluble [Li(glyme)]X and highly soluble [Li(glyme)]X (Figure 1). This classification can be done using the boundary of the solubility limit of ∼1000 mM and can be clearly associated with the criterion for solvate ILs (vide supra). In the “concentrated solutions” such as [Li(G3)]NO3 and [Li(G3)][OTf], glyme molecules that do not participate in the formation of the stable [Li(glyme)]+ remain in the solution and are capable of solvating the additional Li salts (Li2Sm). On the other hand, in the solvate ILs, all of the glyme molecules coordinate with the initially doped Li salts to form the stable solvates. Although the anions themselves potentially serve as a solvent for Li2Sm in the solvate ILs, the coordinating property of [TFSA] and [BETI] is too weak to dissolve Li2Sm. NO3 and [OTf] anions have a stronger interaction with Li+ ion than [TFSA], [BETI], and even the glyme solvents as evidenced by DG/DLi >1 (Table 1). This strong interaction between the Lewis basic anions and Li+ may also be responsible for the higher solubility of Li2S, the least soluble salts among lithium polysulfides, in [Li(G3)]NO3 and [Li(G3)][OTf]. Consequently, the solubility of Li2Sm is much lower in the solvate ILs. The liquid states of [Li(glyme)]X affect the chemical species and proportion of the solubilized Li2Sm. The actual Li2Sm solutions prepared by the reactions of S8 and Li2S contained not only the intended Li2Sm but also other fragments of polysulfides with different chain lengths because of the disproportionate reactions of Sm2−.37 Although it is difficult to analyze the exact composition of the polysulfide species from the UV−vis spectra only, the assignment of the UV−vis absorption bands in organic solvents has been reported:38,39 490−500 nm for S82−, 450−470 nm for S62−, ∼420 nm for S42−, ∼340 nm for S32−, and ∼280 nm for S22−. Figure 2 shows the UV−vis spectra of saturated Li2Sm in [Li(glyme)]X. In Figure 2a, the spectral shapes of S8/Li2S = 7:8 (Li2S8) and S8/Li2S = 1:8 (Li2S2) are almost the same, despite the different mixing ratios of S8 and Li2S before the reactions, implying that identical compositions of Sm2− species are stabilized in

Table 1. Viscosity (η), Ionic Conductivity (σ), and Molar Concentration (c) of [Li(G3)]X and [Li(G4)]X at 30 °Ca glyme−Li salt mixture

η (mPa·s)

σ (mS·cm−1)

c (mol·dm−3)

DG/DLi

[Li(G3)][TFSA] [Li(G3)][OTf] [Li(G3)]NO3 [Li(G4)][TFSA] [Li(G4)][BETI] [Li(G4)]BF4

169 381 206 81 130 314

1.1 0.30 0.31 1.6 0.91 0.50

3.06 3.90 4.77 2.75 2.39 3.87

0.94 1.24 1.79 0.98 1.00 1.11

a

Data were obtained from ref 29 and reorganized in this paper.

is, the glyme and Li+ diffused together, suggesting the presence of the desired complex cations. Thus, they can be regarded as solvate ILs. Their enhanced thermal and electrochemical properties also implied the formation of the complex cations.26,29 On the other hand, DG/DLi is clearly higher than unity for [Li(G3)]NO3 and [Li(G3)][OTf]. This is due to either a fast ligand exchange between the unstable [Li(glyme)]+ cations or the presence of free glymes that do not participate the coordination of Li+ in [Li(glyme)]X. They are classified as concentrated solutions (not as solvate ILs). [Li(G4)]BF4 shows an intermediate value of DG/DLi and appears to be midway between the aforementioned two categories. The order of DG/ DLi is as follows: NO3 > [OTf] > BF4 > [TFSA] ≈ [BETI]. This order is in accordance with the order of Lewis basicity (or donor property) of the anions33 and/or Li-X ion-pair dissociation energy obtained from ab initio calculations.34 It is obvious that the formation of robust [Li(glyme)]+ complex cations, a criterion for solvate ILs, is governed by the competition between glymes and X anions to interact with the Li cation and is achieved only when combined with weakly coordinating anions such as [TFSA] and [BETI]. The saturation solubilities of S8 and Li2Sm in [Li(glyme)]X are shown in Figure 1, wherein the solubility is represented as total atomic S concentration. For Li2Sm, three different mixtures were prepared using the stoichiometric amounts of S8 and Li2S, in accordance to the reaction, 8Li2S + (m − 1)S8 → Li2Sm (m = 8, 4, and 2).32 The saturated concentration of both the starting

Figure 1. Solubility limits of S8 Li2S8, Li2S4, Li2S2, and LiS2, where Li2Sm (m = 8, 4, 2) is the nominal formula, assuming a complete reaction between S8 and Li2S without occurrence of disproportion reactions. [Li(glyme)]X represented by empty and full symbols are categorized as solvate ILs and concentrated solutions, respectively, on the basis of the criteria outlined in ref 29. C

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Figure 2. UV−vis spectra of saturated solutions of S8 and Li2Sm in [Li(glyme)]X. (a) S8, Li2S8, Li2S4, and Li2S2 in [Li(G4)][BETI], where Li2Sm (m = 8, 4, 2) is the nominal formula assuming complete reaction between S8 and Li2S without occurrence of disproportionate reactions; (b) S8/Li2S = 1:8 (Li2S2) in different [Li(glyme)]X; a 1-mm-thick quartz cell was used for measurements.

Figure 3. Charge/discharge curves of Li metal anode |[Li(glyme)]X|S8/KB composite cathode cells at a constant current density of 139 mAg−1 sulfur; (a) [Li(G4)][BETI], (b) [Li(G3)][OTf], (c) [Li(G3)]NO3, (d) [Li(G4)]BF4.

[Li(G4)][BETI]. More interestingly, the shoulder at 617 nm, assigned to the trisulfide anion radical (S3•−),37 is observed in [Li(G3)][OTf] and [Li(G3)]NO3, but not in [Li(G3 or G4)][TFSA]30 and [Li(G4)][BETI] (Figure 2b). Thus, the

[Li(G4)][BETI]. The two shoulders at 280 and 420 nm correspond to S22− and S42−, respectively, as reported previously in literature.38,39 The high absorbance of S8/Li2S = 3:8 (Li2S4) also implies that S42− species are rich in D

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formation of S3•− strongly depends on the liquid states of [Li(glyme)]X (concentrated solution or solvate IL). In fact, the band for S3•− was common in many organic solvents such as dimethylformamide (DMF),40 G439 and dimethyl sulfoxide,41 as well as in the concentrated solutions of [Li(glyme)]X. Interestingly, recent studies revealed the absence of the characteristic absorption (617 nm) of S3•− in aprotic ILs;38 this is consistent with our results in solvate ILs. It seems that the radical form of S3•− cannot be stabilized in IL media consisting of weakly coordinating cations and anions. Li−S Cell Performance. We selected [Li(glyme)]X for this study because of our interest in their electrochemical properties in Li−S cells and their solvating ability for Li2Sm. For instance, the NO3− anion is well-known to form a protective layer on the metallic lithium anode, which prevents the redox shuttle reaction.15 The investigation of the electrochemical stability of the [Li(glyme)]X electrolytes in Li−S cells is also of great importance. To study the performance of [Li(glyme)]X as an electrolyte for the Li−S battery, galvanostatic charge/discharge tests were carried out at a slow charge/discharge current rate of 139 mAg−1 sulfur at 30 °C (Figure 3). Figure 4 shows the cycle

approximately 2.4 V in the charge curves. These voltage plateaus remain unchanged during the repeated charge/ discharge steps even after 100 cycles, while the capacity gradually decreases with the cycle number. Although the details of the electrochemical reactions of sulfur are still unclear,42 the two voltage regions obtained during the discharge step are regarded as the reduction voltages of S8 to Li2Sm (m ≥ 4)30 for the upper voltage plateau and of Li2Sm to Li2S for the lower voltage plateau, respectively. We previously reported that the upper voltage plateau (∼2.3 V) for IL-type electrolytes was slightly lower than that for organic electrolytes (∼2.4 V). This can be interpreted as the “solvent effect” caused by the difference in the solubility of Li2Sm. Li2Sm is more stabilized when it is solvated by the solvent molecules, resulting in a greater positive voltage plateau in organic electrolytes that can easily dissolve Li2Sm. Conversely, in solvate ILs, wherein the solubility of Li2Sm is significantly low, Li2Sm mainly remains trapped as a solid in the composite cathode, and solid-state redox reactions are predominant. The upper voltage plateau of 2.3 V and a high Coulombic efficiency close to 100% over 100 cycles for the [Li(G4)][BETI] electrolyte are indicative of the solid-state reactions. For [Li(G3)][OTf] (Figure 3b), which is not classified as a solvate IL, the plateau voltages as well as the capacity changes with the increasing cycle number. During the second cycle, the first plateau in the discharge curve of 2.4 V is higher than that in [Li(G4)][BETI] (∼2.3 V) and similar to the results in organic electrolytes, thereby suggesting that Li2Sm species are highly solubilized in the [Li(G3)][OTf] electrolyte after the first discharge. This is not in conflict with the extremely high solubility limit, as shown in Figure 1. An increase in the overvoltage with the repetition of the charge/discharge cycles is attributed to the side reactions on the cathode and anode. The dissolved Li2Sm reacts with the metallic lithium anode and partially forms insoluble byproducts on the anode. On the other hand, solubilized Li2Sm that is reduced at the metallic lithium anode diffuses back to the cathode and is reoxidized for the charging reactions (redox shuttle mechanism), in a manner similar to organic electrolytes. In fact, the Coulombic efficiency of [Li(G3)][OTf] is 93% for the second cycle, which is much lower than that for [Li(G4)][BETI] (∼98.5%, Figure 4). The redox shuttle mechanism and the accompanying side reactions results in a capacity loss as well as overvoltage for the charge/ discharge cycle.21,43 With regard to the electrolyte properties (Table 1), [Li(G3)][OTf] is more viscous and less ionconductive than [Li(G4)][BETI], leading to the lower initial discharge capacity. Therefore, [Li(G3)][OTf] is not suitable for use as an electrolyte in a Li−S battery. Unlike [Li(G3)][OTf], [Li(G3)]NO3 was expected to form an effective protective layer on the metallic lithium anode to inhibit the undesired side reactions. However, the results of the charge/discharge test for [Li(G3)]NO3 was worse than that for [Li(G3)][OTf]; a drastic capacity decay over several charge/ discharge cycles is observed (Figure 4a). The discharge capacity of the first cycle exceeds the theoretical value of the S cathode (Figure 3c). The discharge curves have noticeable plateaus at approximately 1.7 V, which is not observed in [Li(G4)][BETI] and [Li(G3)][OTf]. Recently, it was reported that the nitrate anion could be irreversibly reduced on the carbon−sulfur composite cathode during the discharge of the Li−S cell.44 We also confirmed the corresponding reduction current of approximately 1.7 V for Li/Li+ on Ni working electrode in the cyclic voltammetric measurement for [Li(G3)]NO3 (see

Figure 4. Cycling performance of the Li metal anode |[Li(glyme)]X| S8/KB composite cathode cells at a constant current density of 139 mAg−1 sulfur; (a) capacity, (b) Coulombic efficiency. Closed and open plots represent charge and discharge capacities, respectively.

performance of the Li−S cells using the [Li(glyme)]X electrolyte. Li−S battery performance for [Li(G3 or G4)][TFSA] has been reported in our previous work.30 First of all, we emphasize that successful charge/discharge measurements can be repeated only with the solvate IL electrolytes, [Li(G4)][BETI] (Figure 3a) and [Li(G3 and G4)][TFSA],30 among all of the [Li(glyme)]X tested. For [Li(G4)][BETI] (Figure 3a), there are two distinct voltage plateaus of approximately 2.3 and 2.0 V in the discharge curves, whereas there is only one plateau region of E

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suggesting the decomposition of BF4 anion in the presence of Li2Sm. Zhang also reported the reaction between LiBF4 and Li2Sm.45 Considering the results in Figure 5 and the fact that the oxidation of Li2Sm does not occur during the charging step (Figure 3d), we propose a possible side reaction between BF4− and Sm2− as shown in Scheme 1. Li+, a strong Lewis acid, could eliminate the fluorine atom of BF4 anions to yield LiF in the presence of nucleophilic Sm2−. Simultaneously, the resulting Lewis acid, BF3, could form adducts with Sm2−, for example, [BF3−Sm−BF3]2−. The oxidation of Li2Sm to S8 might be disrupted by the formation of the end-capped polysulfides. Contrary to the failure of the [Li(G4)]BF4 electrolyte, a successful redox reaction of sulfur has been reported in the LiBF4-doped poly(ethylene oxide) polymer electrolyte17 and tetraethyl ammonium BF4 in the DMF electrolyte.46 A charge-screening effect derived from the inherently high ionic strength of [Li(G4)]BF4 (c = 3.87 mol· dm−3) could cause an unfavorable reaction between the anionic species (i.e., BF4− and Sm2−). A similar acceleration of the anion−anion reaction has already been reported in the exchange reaction between iodide/triiodide redox couples in ILs.47

Figure S1 of the Supporting Information). Because a large amount of nitrate anions exists in [Li(G3)]NO3, their cathodic decomposition appears to be more pronounced in the Li−S cell, compared to the cells with typical organic electrolytes containing a small amount of LiNO3 as an additive. The reduction product of nitrates has a negative effect on the reversibility of the Li−S cell. Despite the fact that LiNO3 was supposed to be an effective additive, the highly concentrated electrolyte of [Li(G3)]NO3 cannot be used in Li−S batteries. As mentioned earlier, we failed to estimate the solubility limit of Li2Sm in [Li(G4)]BF4 because the oxidation of Li2Sm to S8 did not occur. Likewise, [Li(G4)]BF4 appeared to hinder the reversible redox reaction of sulfur in the charge/discharge test of the Li−S cell. In this case, the charge curves have voltage plateaus at of approximately 3.2 V (Figure 3d), which are not observed in the typical charging curves of Li−S cells. The cell shows rapid capacity decay and the lowest Coulombic efficiency, ∼65%, at 100th cycle (Figure 4). While determining the saturation solubility of Li2Sm in [Li(G4)]BF4, the solution color changed from reddish brown to yellow upon mixing [Li(G4)]BF4 and S8/Li2S in a ratio of 3:8 (nominal Li2S4), followed by stirring for 5 days at 60 °C (see Figure S2). Figure 5a and b shows 11B and 19F NMR spectra for [Li(G4)]BF4

4. CONCLUSIONS In this work, the effects of the anionic species of [Li(glyme)]X electrolytes on the solubility of S8 and Li2Sm and Li−S cell performance were studied. The liquid states, categorized into typical concentrated solutions or solvate ILs by the presence or absence of “free” glymes, respectively, dominated the saturation solubility of Li2Sm in [Li(glyme)]X. In the solvate ILs, where excess glymes capable of solvating Li2Sm did not exist, the solubility limits were more than 1 order of magnitude lower than those for the concentrated solutions, despite the lower salt concentration (c) in the solvate ILs. Furthermore, the absence of S3•− was common in aprotic ILs38 and solvate ILs; however, S3•− was detected in the concentrated solutions of [Li(G3)]NO3 and [Li(G3)][OTf]. In contrast to Li−S cells involving organic electrolyte solutions with different Li salts,13 the anion X in [Li(glyme)]X played a substantial role in the performance of the Li−S cell. [Li(G3)][OTf] did not form a solvate IL because of the strong interaction between Li+ and [OTf]−,29 resulting in highly solubilized Li2Sm, which led to the unfavorable redox shuttle mechanism. In the Li−S cell containing [Li(G3)]NO3, a serious side reaction involving the irreversible cathodic decomposition of NO3− on the S8/KB composite electrode resulted in the poor cycle performance. Moreover, we experimentally revealed a chemical reaction between Sm2− and BF4−. The resulting byproducts thwarted the reversible oxidation of Li2Sm to S8 in [Li(G4)]BF4. A recent study on Li−S battery using aprotic ILs48 suggests that bis(fluorosulfonyl)amide anion ([FSA]−) is not also suitable because of the side reaction of [FSA]− with the polysulfides, although we did not test [Li(glyme)][FSA] as the electrolyte for Li−S cell. Consequently, only the solvate IL

Figure 5. NMR spectra of [Li(G4)]BF4 before (black) and after (red) mixing with polysulfides. (a) 11B- and (b) 19F-NMR. Asterisks represent peaks from byproducts.

before and after mixing with polysulfides, respectively. The small peaks appear in the downfield region of the main peaks for the BF4 anion in both the 11B and the 19F NMR spectra, Scheme 1. Possible Reaction between Sm2− and BF4− Anions

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electrolytes[Li(G4)][BETI] and [Li(G3 or G4)][TFSA]30suppressed the dissolution of Li2Sm effectively and were electrochemically stable in Li−S cells. Therefore, they were found to be the best electrolytes for Li−S batteries among the variants of [Li(glyme)]X. Given the markedly positive effect of the solvation property of ILs on the suppression of polysulfide dissolution, Li−S batteries20,21,23,30 using ILs as electrolytes could be named as Li| Ionic Liquid|S (LILS) batteries. Several researchers have recognized the low Li+ transport property, because of the intrinsic high viscosity of ILs, as a serious shortcoming of the IL-based electrolytes. In fact, this problem has prevented the successful operation of Li−S cells with high current density. However, in the case of solvate ILs, the high viscosity can be reduced by dilution with low-viscous organic solvents, wherein the solvate ILs serve as supporting Li salts in the electrolyte solutions. When the solvate IL [Li(G4)][TFSA] was diluted with a nonpolar, nonflammable hydrofluoroether (HFE), the capacity and rate capability of Li−S cells greatly improved.30 In this mixed electrolyte, the dissociation of the constituent ions was accomplished by the weakly coordinating [Li(G4)][TFSA] (i.e., [Li(G4)]+ and [TFSA]−) by itself, while the nonpolar HFE offered nonflammability and low solvation power for Li2Sm in addition to an improvement in the transport properties. Such a thorough role sharing in the design of electrolytes would further enhance the LILS battery performance.

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ASSOCIATED CONTENT

* Supporting Information S

Cyclic voltammogram of [Li(G3)]NO3 on a Ni disk as the working electrode and color photographs of the mixture of [Li(G4)]BF4 and S8/Li2S = 3:8 (Li2S4). This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Telephone/fax: +81-45-339-3955. E-mail: [email protected]. jp. Present Address

J.-W.P.: Korea Electrotechnology Research Institute (KERI), Changwon 641-120, Republic of Korea. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This study was supported in part by the Advanced Low Carbon Technology Research and Development Program (ALCA) of the Japan Science and Technology Agency (JST) and the Technology Research Grant Program of the New Energy and Industrial Technology Development Organization (NEDO) of Japan.

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REFERENCES

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