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Sulfone-Based Electrolytes for Nonaqueous Li−O2 Batteries Fanny Bardé,*,† Yuhui Chen,‡ Lee Johnson,‡ Stijn Schaltin,§ Jan Fransaer,§ and Peter G. Bruce‡ †

Toyota Motor Europe, R&D 3, Advanced Technology 1, Hoge Wei 33, B − 1930 Zaventem, Belgium Departments of Materials and Chemistry, University of Oxford, Parks Road, Oxford OX1 3PH, United Kingdom § Department of Materials Engineering (MTM), KULeuven, Kasteelpark Arenberg 44 box 2450, 3001 Heverlee, Belgium ‡

ABSTRACT: We investigated the use of sulfone-based electrolytes for the Li−O2 battery. The study compared the behavior of three commercially available sulfones: ethyl vinyl sulfone (EVS), tetramethylene sulfone (TMS), also called sulfolane, and ethyl methyl sulfone (EMS). First, we carried out a preliminary investigation of the oxygen reduction reaction and oxygen evolution reaction (ORR/OER) as a function of solvent type and Li+ concentration. Then, TMS and EMS were tested (LiTFSI salt) in Li−O2 cells. The cells exhibited initial capacities around 1800 and 2000 mAh.g−1carbon, respectively. The capacity retention on cycling was quite low. We analyzed the reaction products during discharge and charge by means of powder X-ray diffraction, infrared spectroscopy, 1 H-nuclear magnetic resonance, and mass spectrometry. Although EVS was at first sight the most attractive sulfone, since it is a liquid at room temperature, it was the least stable in the presence of oxygen; its vinyl group was attacked by reduced O2 species. On the other hand, both TMS and EMS performed better during the first five cycles; Li2O2 formation and decomposition was the main reaction, although some byproducts formed during cycling. After five cycles, there was still a considerable amount of Li2O2 formed, but decomposition to form Li2CO3 became significant, and it accumulated at the O2 electrode. This was the likely reason for capacity fading.

1. INTRODUCTION The rechargeable nonaqueous Li−O2 battery is receiving intense interest at the present time because of its high theoretical energy density (3505 Wh/kg),1−5 which has the potential to enable battery electric vehicles with a driving range approaching 500 km between charging. However, several scientific and technical challenges must be overcome before this technology becomes viable.1−16 Among the present challenges, the search for a stable electrolyte is of critical importance.1−5,8 Early investigations of nonaqueous Li−O2 batteries focused on the use of carbonate-based electrolytes, but those carbonates decompose irreversibly at the cathode to form side-reaction products, typically identified as lithium formate, lithium acetate, lithium propyl-dicarbonate, and lithium carbonate, with little or no evidence of Li 2 O 2 formation.6−10 Attention then turned to both linear (dimethoxyethane, diglyme, triglyme, tetraglyme) and cyclic (1,3dioxolane, 2-methyl-tetrahydrofuran) ether-based electrolytes. Although ethers are more stable than carbonates and hundreds of cycles could be achieved in ether by limiting the depth of discharge,17,18 electrolyte degradation also occurs during cycling.19−21 Several researchers have proposed the use of ionic liquids (IL) because of their relatively high electrochemical and chemical stability against O2 radicals.22−24 One of the limitations of ILs is their high viscosity and, hence, low diffusivity of dissolved species. Recent papers report the use of dimethyl sulfoxide (DMSO).25−29 Sharon et al. showed reversible oxygen reduction reaction and oxygen evolution reaction (ORR/OER) behavior in DMSO at a planar gold electrode by electrochemical quartz crystal microbalance (EQCM). However, by using low volumes of electrolyte © 2014 American Chemical Society

together with high-surface-activated carbon-fiber electrodes to magnify any instability, the same authors identified some side products such as LiOH, dimethylsulfone, Li2SO3, and Li2SO4.28 Younesi et al. demonstrated by X-ray photoelectron spectroscopy analysis that after 2 days, DMSO decomposed on the Li2O2 surface whereas no indication of degradation was observed for acetonitrile.29 On the other hand, a screening of the stability of solvents against superoxide radicals demonstrated that acetonitrile, PP13TFSI, and DMSO were the most stable solvents.30 Interestingly, using a combination of DMSO and a nanoporous gold electrode, Peng et al. achieved sustainable and highly reversible Li−O2 formation/decomposition on cycling.25 These different results for DMSO likely highlight the reality that no solvent is completely stable. Side reactions can always be detected depending on the experimental conditions. The key questions are not whether there is decomposition but how much, what is the nature of the decomposition products, and are they electrochemically active? The issue for the use of a given solvent in a practical Li−O2 cell is whether the degree of decomposition over the number of cycles and lifetime of the cell is small compared with the electrolyte volume. Again, the nature of the side reaction products also matters. Finally, a recent paper suggests that N,Ndimethylacetamide may also be a suitable basis for an electrolyte in a Li−O2 cell.31 As the hunt for a good Li−O2 battery electrolyte is still ongoing, we decided to investigate sulfones as a possible Li−O2 Received: May 16, 2014 Revised: July 18, 2014 Published: July 23, 2014 18892

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electrolytes containing 0.1 and 1 M LiTFSI based on DMSO (CH3SOCH3) (Aldrich) were prepared in the same way. In solutions without Li+, 1 M tetrabutylammonium perchlorate (TBAP) (Aldrich) was dissolved. The TBAP was dried at 140 °C under vacuum overnight prior to use. 2.2. Cyclic Voltammetry Tests. CVs were performed using a potentiostat/galvanostat (EG&G 273) controlled by CorrWare software and a VMP3 potentiostat (Biologic). Wellpolished glassy carbon (GC), fritted Ag wire, and platinum wire (Pt) were used as working, reference, and counter electrodes, respectively. The true potential was established versus the ferrocenium/ferrocene (Fc/Fc+) (Aldrich) redox couple. The cyclic voltammetry was performed at 100 mV s−1 at room temperature. Electrolytes were saturated with O2 by bubbling it through the solution for 10 min prior to the start of the experiments. 2.3. Li−O2 Cell Assembly and Electrochemical Measurements. Porous O2 electrodes (Super P carbon: poly(tetrafluoroethylene), PTFE) supported on a stainless steel mesh were prepared and dried following the previously published procedure.37 Nanoporous gold (NPG) electrode foils were prepared by dealloying white gold leaf by floating it on a bath of concentrated nitric acid for 5 min, as previously reported.37 This process resulted in a free-standing film of NPG. The NPG was dried by heating under vacuum at 150 °C overnight. As Li metal stability and cycling efficiency in sulfones are not yet clearly demonstrated and very few studies are dedicated to this topic,38 we use lithium iron phosphate (LiFePO4) composite electrode as both counter and reference electrode. LiFePO4 and Super P carbon (Timcal) were mixed with ethanol before addition of a PTFE suspension (Aldrich, 60 wt %). The slurry (LiFePO4:SuperP:PTFE with 9:0.3:0.7 weight ratio) was stirred for 10 min and was cast onto stainless steel mesh disks (Advent-RM, 100 mesh per inch). Prior to being dried at 200 °C under vacuum for 12 h, the porous electrodes were washed with ethanol: water (1:1) mixture to remove the surfactant from the PTFE suspension. Swagelok cells were assembled in an argon-filled glovebox, were placed in a glass tube, and were purged with pure dried O2 for 30 min prior to measurement. Cycling was carried out on a Maccor instrument, between 2.4 and 4.2 V (vs Li/Li+) at a rate of 70 mA.g−1carbon. 2.4. Cell Disassembly and Analysis of (Dis)charge Products. The procedures of disassembly of Li−O2 cells and the extended examination of O2 cathodes by XRD, FTIR, 1 HNMR, and in situ DEMS were described in detail in a previous publication.37

battery solvent. Sulfones have been studied previously for application in high-voltage Li-ion batteries,32−34 and sulfolane (TMS) has recently been examined as a possible electrolyte for the Li−O2 battery.35 The highest ionic conductivity for a sulfone-based electrolyte at room temperature is obtained with 1 M lithium bis(trifluoromethylsulfonyl)imide (LiTFSI) in ethyl vinyl sulfone (EVS), tetramethylene sulfone (TMS), and ethyl methyl sulfone (EMS). In addition, these three sulfones have a relatively low melting point (Table 1) compared to other Table 1. Formulae and Physico-Chemical Properties of Sulfones36

sulfones such as butyl sulfone (44 °C) or 1-fluoro-2-methylsulfonyl benzene (50 °C).34 For these reasons, we focused our attention on EVS, EMS, and TMS. In addition, we selected both cyclic and linear sulfones in order to compare their stability, expecting the cyclic sulfolane to be less stable than the linear sulfones.19 We first investigated the oxygen reduction reaction/oxygen evolution reaction (ORR/OER) by cyclic voltammetry (CV) for various Li+ concentrations to obtain a preliminary idea of the stability of the solvents. DMSO was included here for comparison purposes, as it has already been shown to be relatively stable25−27,30 compared to other solvents investigated to date.6−10,17−24,30 Then, the electrochemical performance of Li−O2 batteries employing sulfones as the basis of the electrolyte was evaluated. Finally, the reaction mechanism at the O2 cathode while cycling was studied by conducting analyses of (dis)charge products by several complementary techniques: powder X-ray diffraction (PXRD), Fourier transform infrared spectroscopy (FTIR), 1 H-nuclear magnetic resonance (1HNMR), and in situ differential electrochemical mass spectrometry (DEMS). The results established the level of stability of these sulfone electrolytes in the presence of O2 species and their potential use as Li−O2 battery electrolytes.

3. RESULTS AND DISCUSSION 3.1. Voltammetric Analysis of ORR/OER. The electrochemical windows of TMS, EMS, and EVS sulfones and DMSO (reported here for comparison purposes) were determined using TBAP as the supporting electrolyte at room temperature or at 50 °C if the electrolyte was solid at room temperature (Figure 1a−d). The anodic stability deceased in the order TMS (5.2 V vs Li/Li+) > EMS (5 V vs Li/Li+) > EVS (4.3 V vs Li/ Li+), while DMSO had the lowest oxidation stability (4.2 V vs Li/Li+). The cyclic voltammograms obtained in DMSO are shown in Figure 2a. In the presence of O2, but without Li+, cathodic and anodic peaks were observed, confirming that O2 reduction to O2− is chemically reversible in DMSO, as demonstrated by Laoire et al.39 When both O2 and Li+ were present, the oxidation peak for O2− − e− → O2 disappeared,

2. EXPERIMENTAL SECTION 2.1. Preparation of Electrolytes. EVS (C2H5SO2C2H3), TMS (C4H8O2S), and EMS (C2H5SO2CH3) (Aldrich) were dried with activated 4 Å molecular sieves for 2 days before use (Table 1). Battery-grade LiTFSI (LiN(SO2CF3)2) (Aldrich) was dried at 140 °C under vacuum for 12 h. The 0.1 and 1 M LiTFSI−sulfone solutions were prepared in a high-integrity argon-filled glovebox (MBRAUN) with less than 1 ppm of oxygen and water. Because EMS and TMS are solid at room temperature, the samples were heated to obtain a liquid. Adding 1 M LiTFSI salt to the EMS and TMS solvents decreases their melting point so that they become viscous liquids at room temperature. For comparison purposes, 18893

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of TBA+. A general reaction scheme for reduction and oxidation has been proposed previously:39 O2 + Li+ + e− → LiO2

(1)

2LiO2 → Li 2O2 + O2

(2)

LiO2 + Li+ + e− → Li 2O2

(3)

Li 2O2 + 2Li+ + 2e− → 2Li 2O

(4)

in which eq 2 is not electrochemical, but a purely chemical reaction. Possible oxidation reactions in the presence of Li+ are Li 2O2 → LiO2 + Li+ + e− +

LiO2 → O2 + Li + e

−

Li 2O2 → O2 + 2Li+ + 2e−

Figure 1. Electrochemical windows of (a) DMSO, (b) TMS, (c) EMS, and (d) EVS at room temperature (except for EMS which was measured at 50 °C) measured using 1 M TBAP as supporting electrolyte and at a GC electrode.

(5a) (5b) (6)

1 O2 + 2Li+ + 2e− (7) 2 The presence of Li2O at the end of discharge 4 and especially its oxidation on charge 7 remains a point of debate.39 Nevertheless, recent in situ X-ray photoelectron spectroscopy (XPS) studies on Li-oxygen electrochemistry clearly show evidence of the formation and oxidation of Li2O when solidstate electrolytes are used.40,41 The broadness of the oxidation peak in DMSO and the presence of Li+ indicate that the oxidation is complex. To investigate the reactions in more detail, CVs with different cutoff potentials on reduction were recorded, Figure 3a and b. Provided reduction is confined to Li 2O →

Figure 3. Cyclic voltammograms applying different cutoff potential with O2 in (a) 0.1 M Li+ DMSO, (b) 1 M Li+ DMSO, (c) 0.1 M Li+ TMS, and (d) 1 M Li+ TMS. Rate 100 mV.s−1.

the higher potential peak, as the reduction limit is then made more negative, the magnitudes of the oxidation peaks at around 2.6 and 3 V increase. On extending the cutoff to lower potentials, such that the second reduction peak is included, the oxidation peak at ∼2.6 V disappears and the broad oxidation peaks at the higher voltages, >3.3 V, increase, Figure 3a and b. The oxidation processes above 3 V are most likely associated with oxidation of Li2O2. However, DMSO is not the subject of this paper, which instead focuses on TMS, EVS, and EMS.

Figure 2. Cyclic voltammograms in 1 M TBAP (black and blue curves), 0.1 M LiTFSI (red curves), and 1 M LiTFSI (green curves) in (a) DMSO, (b) TMS, (c) EMS, and (d) EVS without and with O2. Rate 100 mV.s−1.

reflecting the fact that the O2 reduction mechanism in Li+containing solutions is different from that seen in the presence 18894

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Interestingly, the first discharge plateaus in both systems are marginally lower (50 mV) than the plateaus of the following cycles, and the first charge plateaus are higher (100 mV) than the plateaus of the following cycles. This may indicate some degree of electrode conditioning, a phenomenon often observed in intercalation electrodes in lithium-ion batteries.42−45 The fifth discharge capacity is over 1000 mAh· g−1carbon in both EMS and TMS. Because capacities fade quickly after 5−6 cycles and because almost no capacity is left after 10 cycles, Figure 4b and c, the analysis of the electrodes on the first discharge and first charge as well as on the fifth discharge and fifth charge was carried out. The PXRD data at the end of the first discharge shows clear evidence of the presence of Li2O2 for both EMS and TMS, Figure 5a and b.

Therefore, let us now consider the CVs obtained in these three sulfones, Figure 2b−d, commencing with TMS. In the absence of Li+, O2 reduction is chemically reversible, although not completely, as demonstrated by the lower charge on oxidation of O2− compared with reduction of O2, perhaps indicating that O2− reacts with TMS to some extent. Addition of Li+ eliminates the oxidation of O2−, and a broad oxidation peak at higher potentials (3.4−3.8 V) appears which may be associated with oxidation of Li2O2. The absence of the oxidation peak at ∼2.6 V, associated with O2− oxidation, implies that O2− reacts quickly with Li+. As the cutoff potential on reduction is made more negative, the oxidation peaks between 3.4 and 3.8 V increase, Figure 3c and d, similar to the DMSO case. Turning to EMS, in absence of Li+, O2 reduction is chemically reversible although not completely as for TMS. Addition of Li+ eliminates the oxidation of O2−, but no broad peak appears at higher potentials as was the case for TMS. Finally EVS, Figure 2 d, shows no evidence of O2− oxidation in the absence of Li, indicating that O2− reacts with the solvent. In the presence of Li+, there is only minor evidence for oxidation processes above 3 V. In some cases (DMSO and EVS), the reduction is composed of two peaks, and in the others, it is a broad peak; both are consistent with a complex, multistep reduction, in accord with the mechanisms shown in eqs 1−3. It is clear that the electrochemistry of sulfones in O2 is complex on reduction and oxidation. It cannot be understood fully on the basis of the CVs presented here alone. However, as the purpose of the present paper is to investigate the efficacy of sulfone-based electrolytes for use in Li−O2 batteries and not the detailed electrochemistry of sulfones in O2 per se, we have used the CVs only as a preliminary indicator of stability, investigating below the electrolytes in Li−O2 batteries. From this preliminary CV screening, it is clear that EVS is unstable in the presence of superoxide radicals which we expect to attack the vinyl group known to be very reactive. 1H NMR studies (not reported here) on products extracted from an air cathode after the first discharge in EVS confirmed the presence of decomposition products from residual EVS; this result confirms that the vinyl group in EVS is easily attacked by the superoxide radicals. We therefore decide to focus our deeper investigation on TMS and EMS as electrolytes for Li−O2 batteries. 3.2. Cycling and Analysis of (Dis)charge Products from Li−O2 Batteries Using EMS and TMS-Based Electrolytes. Li−O2 batteries with 1 M LiTFSI in EMS or with 1 M LiTFSI in TMS were investigated. First discharge capacities were over 1800 mAh.g−1carbon, and the discharge plateaus were the same in both solvents, Figure 4. Charge plateaus at 3.9 and 3.8 V were observed for EMS and TMS, Figure 4a.

Figure 5. PXRD patterns (Kα1Cu) collected from cathodes discharged and charged in 1 M LiTFSI in (a) EMS and (b) TMS. *, Li2O2; ▲, Li2CO3.

FTIR spectra support this view, Figure 6a and b; however, the broad peaks at 1400−1600 cm−1 suggest the formation of

Figure 6. FTIR spectra of O2 cathodes discharged in 1 M LiTFSI in (a) EMS and (b) TMS.

Li2CO3, although this is not evident in the PXRD patterns at the end of the first discharge. We suppose that Li2CO3 is amorphous and cannot be detected by PXRD. Lithium formate (1610 cm−1 in FTIR) is also identified on the first discharge, as are other peaks in the range of 900−1300 cm−1, the origin of which could not be identified. 1HNMR results confirm the presence of lithium formate and indicate the coexistence of lithium acetate impurities, Figure 7a and b.

Figure 4. (a) Comparison of cycle 1 for Li−O2 batteries using 1 M LiTFSI in EMS or TMS as the electrolytes; load curves for 10 cycles of the Li−O2 battery (b) in EMS and (c) in TMS. 18895

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Figure 8. In situ DEMS measured during (a, c) 5th discharge and (b, d) 5th charge of a Li−O2 battery with 1 M LiTFSI in (a, b) EMS and (c, d) TMS electrolyte; rate of CV: 0.05 mV.s−1; cell current, blue curve; flux of O2, green curve; and CO2, red curve, 100 times amplified. Flow of carrier gas: 0.2 mL min−1.

Figure 7. 1HNMR spectrum of D2O solution after washing the cathodes discharged and charged in 1 M LiTFSI (a) EMS and (b) TMS.

decomposition of side-reaction products on the fifth charge observed by FTIR previously (Figure 6b). It has been shown recently that carbon-based electrodes not only undergo oxidation above 3.5 V in the presence of Li2O2 oxidation but also promote electrolyte decomposition on discharge and charge.46 To verify that the side reactions observed above are due primarily to solvent decomposition rather than to the electrode, the composite Super P-based cathodes were replaced with nanoporous gold (NPG), which has been shown previously to be stable.25 The FTIR spectra of cathodes discharged in EMS or TMS give clear evidence of Li2O2 formation on the first discharge, Figure 9. We also noted

Additional peaks are also evident, in accordance with the presence of other decomposition products, as noted in the FTIR. On the first charge, the PXRD data show that the Li2O2 decomposes, which is supported by FTIR. The FTIR also shows that Li2CO3 persists in the case of EMS. At the end of the fifth discharge, Li2O2 is still evident in both PXRD and FTIR results. Li2CO3 is now observed in the PXRD patterns of electrodes that had been cycled in TMS based electrolyte, Figure 5b. In the FTIR, additional peaks between 900 and 1300 cm−1 are again observed. The FTIR spectra show that the ratio of Li2CO3 to Li2O2 increases, in both EMS and TMS, because of the accumulation of Li2CO3 during cycling, which is likely to cause the capacity fading, as noted for other electrolytes previously.19,37 On the fifth charge, there is some residual Li2O2 together with Li2CO3 and Li formate in the FTIR, indicating that all three compounds accumulate. In situ DEMS for the EMS based electrolyte shows O2 consumption on fifth discharge and the corresponding O2 evolution on charge, Figure 8a and b. O2 reduction starts at 2.8 V, and Li2O2 decomposition starts at 3.1 V. The amount of O2 consumed and released follows the current measured during the CV. The molar ratio of electrons to O2 on discharge is 2.15, which supports the formation of Li2O2 observed by PXRD and FTIR along with some side reactions (2e−/O2 would correspond to Li2O2 alone). However, on the fifth charge, CO2 evolution was observed, especially above 3.6 V. The CO2 evolution is expected to arise from decomposition of the sidereaction products. The ratio of electrons to O2 is around 2.3 on charge, in accordance with the presence of side reactions. DEMS in the TMS-based electrolyte shows similar results. O2 reduction starts from 2.8 V, Figure 8c, and the O2 consumption corresponds to 2.12 mol of electrons per mole of O2. Very little CO2 evolution was observed during the discharge. On charge, the decomposition of products formed on the previous discharge starts from 3 V with the evolution of oxygen from the decomposition of Li2O2. Above 3.6 V, more than 2.33 e−/ O2 is observed and CO2 is evolved (Figure 8d), consistent with

Figure 9. FTIR spectra of nanoporous gold (NPG) electrode discharged in 1 M LiTFSI EMS and TMS.

evidence of Li2CO3; the amount was somewhat smaller compared with the data in Figure 6, indicating that when carbon is replaced by NPG there is less Li2CO3 formation. These results support the view that both electrolyte decomposition and the carbon electrode are a source of Li2CO3 formation.14,15,46 However, of significance for the present paper, the data for nanoporous gold after five cycles shows an increase in the 18896

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(10) Xu, W.; Viswanathan, V. V.; Wang, D.; Towne, S. A.; Xiao, J.; Nie, Z.; Hu, D.; Zhang, J.-G. Investigation on the Charging Process of Li2O2-Based Air Electrodes in Li-O2 batteries with Organic Carbonate Electrolytes. J. Power Sources 2011, 196, 3894−3899. (11) Oh, S. H.; Black, R.; Pomerantseva, E.; Lee, J.-H.; Nazar, L. F. Synthesis of a Metallic Mesoporous Pyrochlore as a Catalyst for Lithium−O2 Batteries. Nat. Chem. 2012, 4, 1004−1010. (12) Bryantsev, V. S.; Giordani, V.; Walker, W.; Blanco, M.; Zecevic, S.; Sasaki, K.; Uddin, J.; Addison, D.; Chase, G. V. Predicting Solvent Stability in Aprotic Electrolyte Li-Air Batteries: Nucleophilic Substitution by the Superoxide Anion Radical (O2(•-)). J. Phys. Chem. A 2011, 115, 12399−12409. (13) Lu, Y.-C.; Gallant, B. M.; Kwabi, D. G.; Harding, J. R.; Mitchell, R. R.; Whittingham, M. S.; Shao-Horn, Y. Lithium−Oxygen Batteries: Bridging Mechanistic Understanding and Battery Performance. Energy Environ. Sci. 2013, 6, 750−768. (14) McCloskey, B. D.; Speidel, A.; Scheffler, R.; Miller, D. C.; Viswanathan, V.; Hummelshoj, J. S.; Norskov, J. K.; Luntz, A. C. Twin Problems of Interfacial Carbonate Formation in Nonaqueous Li−O2 Batteries. J. Phys. Chem. Lett. 2012, 3, 997−1001. (15) McCloskey, B. D.; Scheffler, R.; Speidel, A.; Girishkumar, G.; Luntz, A. C. On the Mechanism of Nonaqueous Li-O2 Electrochemistry on C and Its Kinetic Overpotentials: Some Implications for Li-Air Batteries. J. Phys. Chem. C 2012, 116, 23897−23905. (16) Gallant, B. M.; Mitchell, R. R.; Kwabi, D. G.; Zhou, J.; Zuin, L.; Thompson, C. V.; Shao-Horn, Y. Chemical and Morphological Changes of Li-O2 Battery Electrodes upon Cycling. J. Phys. Chem. C 2012, 116, 20800−20805. (17) Amine, K.; Curtiss, L. A.; Lu, J.; Chun Lau, K.; Zhang, Z.; Sun, Y.-K. Li-Air Batteries Having Ether-Based Electrolytes. U.S. Patent 2013/0230783 A1. (18) Jung, H. G.; Hassoun, J.; Park, J.-B.; Sun, Y.-K.; Scrosati, B. An Improved High-Performance Lithium−Air Battery. Nat. Chem. 2012, 4 (7), 579−585. (19) Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwick, L. J.; Bardé, F.; Bruce, P. G. The Lithium−Oxygen Battery with Ether-Based Electrolytes. Angew. Chem., Int. Ed. 2011, 50, 8609−8613. (20) Veith, G. M.; Nanda, J.; Delmau, L. H.; Dudney, N. J. Influence of Lithium Salts on the Discharge Chemistry of Li-Air Cells. J. Phys. Chem. Lett. 2012, 3, 1242−1247. (21) Younesi, R.; Hahlin, M.; Treskow, M.; Scheers, J.; Johansson, P.; Edström, K. Ether Based Electrolyte, LiB(CN)4 Salt and Binder Degradation in the Li−O2 Battery Studied by Hard X-ray Photoelectron Spectroscopy (HAXPES). J. Phys. Chem. C 2012, 116, 18597−18604. (22) Mizuno, F.; Nakanishi, S.; Shirasawa, A.; Takechi, K.; Shiga, T.; Nishikoori, H.; Iba, H. Design of Non-Aqueous Liquid Electrolytes for Rechargeable Li-O2 Batteries. Electrochemistry 2011, 79, 876−881. (23) De Giorgio, F.; Soavi, F.; Mastragostino, M. Effect of Lithium Ions on Oxygen Reduction in Ionic Liquid-Based Electrolytes. Electrochem. Commun. 2011, 13, 1090−1093. (24) Herranz, J.; Garsuch, A.; Gasteiger, H. A. Using Rotating Ring Disc Electrode Voltammetry to Quantify the Superoxide Radical Stability of Aprotic Li−Air Battery Electrolytes. J. Phys. Chem. C 2012, 116, 19084−19094. (25) Peng, Z. Q.; Freunberger, S. A.; Chen, Y. H.; Bruce, P. G. A Reversible and Higher-Rate Li-O2 Battery. Science 2012, 337, 563− 566. (26) Xu, D.; Wang, Z.-l.; Xu, J.-j.; Zhang, L.-l.; Zhang, X.-b. Novel DMSO-Based Electrolyte for High Performance Rechargeable Li−O2 Batteries. Chem. Commun. 2012, 48, 6948−6950. (27) Trahan, M. J.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M. A.; Abraham, K. M. Studies of Li−Air Cells Utilizing DimethylsulfoxideBased Electrolyte. J. Electrochem. Soc. 2013, 160, A259−A267. (28) Sharon, D.; Afi, M.; Noked, M.; Garsuch, A.; Frimer, A. A.; Aurbach, D. Oxidation of Dimethyl Sulfoxide Solutions by Electrochemical Reduction of Oxygen. J. Phys. Chem. Lett. 2013, 4, 3115− 3119.

amount of Li2CO3 confirming that EMS and TMS are unstable and hence not suitable as the basis of electrolytes for a Li−O2 battery.

4. CONCLUSIONS It has been shown that ethyl methyl sulfone and tetramethylene sulfone are more stable than ethyl vinyl sulfone as electrolytes for Li−O2 batteries. Up to five cycles, Li2O2 formation decomposition is the main reaction in EMS and TMS. However, the capacity fades, and this is associated with the accumulation of Li2CO3 at the cathode, clearly identified by FTIR and PXRD. After 10 cycles, capacity fading is severe. Although degradation of the carbon electrode contributes to the formation of Li2CO3, as demonstrated by the experiments using a nanoporous gold electrode, the electrolyte decomposition is still a major problem rendering the sulfones unsuitable as electrolytes for Li−O2 batteries.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +32 2 712 88 03. Fax: +32 2 712 33 99. Author Contributions

F.B. and P.G.B. wrote the manuscript. Y.C., L.J., S.S., and J.F. performed the experiments and analyzed the experimental data. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS F.B. would like to thank the financial support of the Battery Research Division (M6) at the Toyota Higashi-Fuji Technical Centre.

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REFERENCES

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