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. Author manuscript; available in PMC: 2020 Oct 10.
Published in final edited form as: Chem. 2019 Jul 25;5(10):2630–2641. doi: 10.1016/j.chempr.2019.07.003

Molecular Design of Stable Sulfamide- and Sulfonamide-based Electrolytes for Aprotic Li-O2 Batteries

Shuting Feng 1,5,*, Mingjun Huang 2,5, Jessica R Lamb 2, Wenxu Zhang 2, Ryoichi Tatara 3, Yirui Zhang 4, Yun Guang Zhu 3, Collin F Perkinson 2, Jeremiah A Johnson 2,6,*, Yang Shao-Horn 3,4
PMCID: PMC7442112  NIHMSID: NIHMS1535873  PMID: 32832724

SUMMARY

Electrolyte instability is one of the most challenging impediments to enabling Lithium-Oxygen (Li-O2) batteries for practical use. The use of physical organic chemistry principles to rationally design new molecular components may enable the discovery of electrolytes with stability profiles that cannot be achieved with existing formulations. Here, we report on the development of sulfamide- and sulfonamide-based small molecules that are liquids at room temperature, capable of dissolving reasonably high concentration of Li salts (e.g., LiTFSI), and are exceptionally stable under the harsh chemical and electrochemical conditions of aprotic Li-O2 batteries. In particular, N,N-dimethyl-trifluoromethanesulfonamide was found to be highly resistant to chemical degradation by peroxide and superoxide, stable against electrochemical oxidation up to 4.5 VLi, and stable for > 90 cycles in a Li-O2 cell when cycled at < 4.2 VLi. This study provides guiding principles for the development of next-generation electrolyte components based on sulfamides and sulfonamides.

Keywords: Electrochemistry, Li-O2 batteries, electrolyte design, (electro)chemical stability, sulfonamide, sulfamide, peroxide, superoxide, singlet oxygen

Graphical Abstract

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eTOC blurb

Aprotic lithium-oxygen (Li-O2) batteries show great promise in energy storage and transportation applications owing to their high gravimetric energies that potentially represent a several-fold increase over lithium-ion batteries. The stable and reversible operation of Li-O2 batteries, however, is currently hindered by the severe degradation of common electrolytes. Here, we showed that sulfonamide-based electrolytes, designed based on physical organic chemistry principles, can exhibit superior (electro)chemical stability than common electrolytes such as tetraglyme and DMSO.

INTRODUCTION

Aprotic lithium-oxygen (Li-O2) batteries show great promise in energy storage and transportation applications owing to their high gravimetric energies that potentially represent an increase of 3 to 5 times over lithium-ion batteries.1–4 The stable and reversible operation of Li-O2 batteries is currently hindered by the severe degradation of common electrolytes. Indeed, many of the commonly used electrolyte components of well-established battery chemistries (e.g., Li-ion), such as carbonates,5–8 glymes,9–11 dimethyl sulfoxide (DMSO),12–14 and N,N-dimethylformamide (DMF),15 are not stable in the radical-rich, basic, nucleophilic, and oxidizing environment of the oxygen electrode of Li-O2 batteries (Figure 1). Though reformulation of classical electrolyte components, e.g., by using high salt concentrations, has led to significant stability improvements in some systems, the path toward practical Li-O2 batteries will likely require the rationale molecular design of novel electrolyte components.16 In an early example of such an approach, Nazar and coworkers17 substituted the secondary hydrogens of 1,2-dimethoxyethane (DME) with methyl groups (–CH3) to produce a new solvent with improved stability against hydrogen abstraction. More recently, ketone18- and pivalate19-based electrolyte solvents were reported to be reasonably stable in Li-O2 cells, though cycling studies were limited. Despite these examples, the rational design of electrolyte components remains an underutilized strategy for the discovery of next-generation electrolytes.

Figure 1. Dominant degradation mechanisms of carbonate-, ether-, and sulfoxide-based electrolytes and the molecular design of stable sulfamide- and sulfonamide-based solvents, BTMSA, BMCF3SA, and DMCF3SA, for aprotic Li-O2 batteries.

Figure 1.

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RESULTS AND DISCUSSION

Herein, we report three compounds: N-Butyl-N,N’,N’-trimethylsulfamide (BTMSA), N- butyl-N-methyl-trifluoromethanesulfonamide (BMCF3SA), and N,N-dimethyl-trifluoromethanesulfonamide (DMCF3SA) that are promising for use in aprotic Li-O2 batteries (Figure 1). These compounds are polar aprotic liquids at room temperature, capable of dissolving reasonably high amounts of common Li salts such as lithium bis(trifluoromethane)sulfonimide (LiTFSI). BTMSA has been considered as an electrolyte candidate for lithium batteries,20 and DMCF3SA was recently employed in lithium-sulfur (Li-S) batteries to suppress polysulfide solubility and shuttling.21 However, the utility of these compounds as chemically and electrochemically stable electrolyte components in aprotic Li-O2 batteries has not, to our knowledge, been examined. Guided by a comprehensive stability framework for organic molecules in the Li-O2 oxygen electrode environment and the structure-stability relationships obtained from our previous works,16,22 we designed the structures of these electrolyte components to be lack of (1) vulnerable C-H sites for hydrogen/proton removal, (2) electrophilic centers susceptible against nucleophilic substitution, and (3) highly electron-donating chemical moieties that are vulnerable against electrochemical oxidation. In this study, these three compounds are shown both experimentally and computationally (Figure S1) to exhibit exceptional stability toward a variety of harsh conditions encountered in the Li-O2 cathode environment. For example, all three compounds are compatible with reactive species such as peroxide, superoxide, and singlet O2, while the trifluoromethylsulfonamides BMCF3SA and DMCF3SA display enhanced electrochemical oxidative stability due to the electron withdrawing CF3 group. The latter could be cycled for >90 times in a Li-O2 cell without capacity decay, suggesting that it could represent a significant new addition to the toolbox of electrolyte components. Altogether, these results highlight the power of rational molecular design for electrolyte discovery for Li-O2 batteries.

We began by using a computational framework16 to identify functional groups that may be stable in the Li-O2 cathode environment. Sulfamides and trifluoromethylsulfonamides lacking acidic protons and weak C-H bonds stood out as promising starting points; notably, the exceptional stability of the TFSI anionic component of the commonly used LiTFSI salt provided further motivation for exploring these scaffolds. BTMSA, BMCF3SA, and DMCF3SA were selected for further investigation: BTMSA was synthesized from N,N- dimethylsulfonamoyl chloride and N-butylmethylamine in 90% yield while BMCF3SA and DMCF3SA were prepared via condensation of trifluoromethanesulfonyl chloride and the corresponding secondary amines in 80% – 90% yields. These compounds are clear, colorless liquids at room temperature; their boiling temperatures, viscosities, and dielectric constants (ɛ) are provided in Table S1.

The donor number (DN) of an electrolyte component,23 which is a quantitative descriptor Lewis basicity, can influence the solubility24–27 and lifetime24 of intermediates such as LiO2,24–26 as well as the discharge product morphology24,28,29 and capacity of Li-O2 batteries.24–27 The DNs of BTMSA, BMCF3SA, and DMCF3SA were estimated by 23Na NMR.30 Solutions of 20 mM NaTFSI were prepared in BTMSA, BMCF3SA, DMCF3SA as well as in commonly used electrolyte solvents DMSO, DMF, DME, and propylene carbonate (PC) with 0.5 M sodium perchlorate (NaClO4) in deionized water (H2O) as the internal reference. The 23Na NMR shifts of NaTFSI in these seven electrolytes are shown in Figure 2A. The more upfield (more negative) 23Na shifts recorded in the sulfamide- and sulfonamide-based electrolytes indicate weaker interactions between Na+ and the solvation shell in these electrolytes.30,31 Using the known DNs of DMSO (29.8 kcal/mol32), DMF (26.6 kcal/mol32), DME (20.2 kcal/mol31), and PC (15.1 kcal/mol33) and their 23Na NMR shifts,34 the DNs of BTMSA, DMCF3SA, and BMCF3SA were estimated to be 16.9, 16.4, and 13.3 kcal/mol, respectively (Figure 2B). The low DNs of these three solvents suggest that they will have reduced superoxide solubility24,26 and thus improved chemical stability compared to high-DN solvents such as DMSO and DMF.

Figure 2. Donor number (DN) and conductivity of BTMSA, DMCF3SA, and BMCF3SA.

Figure 2.

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(A) The measured 23Na NMR chemical shifts of 20 mM NaTFSI in BTMSA, DMCF3SA, and BMCF3SA compared to that of DMSO, DMF, DME, and PC (400 MHz). The 23Na signal from the internal reference, 0.5 M NaClO4 in H2O, is set to 0 ppm.

(B) A trend line, indicated by the dashed line, correlating the DNs of DMSO (29.8 kcal/mol32), DMF (26.6 kcal/mol32), DME (20.2 kcal/mol31), and PC (15.1 kcal/mol33) and their measured relative 23Na NMR shifts was used to estimate the DNs of BTMSA, DMCF3SA, and BMCF3SA, which were determined to be 16.9, 16.4, and 13.3 kcal/mol, respectively.

(C) Conductivity of solutions containing 0.1 M LiTFSI in BTMSA, DMCF3SA, and BMCF3SA compared to that of G4 as a commercial reference at various temperatures.

As electrolytes with lower DNs coordinate with Li+ more weakly, potentially leading to lower charge carrier concentrations and conductivities (nonetheless, we note that besides the DN, solvent dielectric constant and viscosity also influence the ion conductivity), we then employed electrochemical impedance spectroscopy (EIS) to investigate the ionic conductivities of solutions containing 0.1 M LiTFSI in BTMSA, BMCF3SA, and DMCF3SA as a function of temperature, which was compared to that of tetraglyme (G4, DN = 16.6 kcal/mol31) as a reference due to its relative stability against oxygen and its reduction products (Figure 2C).The BTMSA-LiTFSI solution exhibited conductivity approximately two times greater than that of G4 whereas DMCF3SA- and BMCF3SA-LiTFSI solutions had lower conductivities by ~2 and 5 times than that of BTMSA, respectively. These values can be rationalized by considering the dielectric constants (ɛ), DNs and viscosities of these compounds. Although BTMSA and G4 have similar DNs and viscosities, BTMSA has a considerably higher dielectric constant than G4 (>29 vs. 7.79,35 Table S1) and can better screen charges, leading to overall higher charge carrier concentration and conductivity. Additionally, LiTFSI is less dissociated in DMCF3SA than BTMSA, supported by higher Raman shifts of the S-N symmetric stretching of the TFSI anion36,37 (Figure S2), resulting in lower charge carrier concentration and conductivity. Furthermore, we note that BMCF3SA not only solvates Li+ more weakly than DMCF3SA, leading to lower charge carrier concentration, but also has higher viscosity (Table S1), both of which contributed to its lower conductivity than DMCF3SA.

The chemical stabilities of BTMSA, BMCF3SA, and DMCF3SA were evaluated under conditions mimicking the oxygen electrode of aprotic Li-O2 batteries using a previously established protocol.16,22 These electrolyte components were combined with 0.5 equivalents of lithium peroxide (Li2O2) and KO2 powders, and the resulting mixtures were stirred at 80°C for three days. The lack of appreciable change in the 1H NMR spectra collected before and after the exposure to Li2O2 and KO2 (Figure 3A–C) indicate that these compounds are highly resistant to chemical degradation by peroxide and superoxide, in good agreement with our computational analyses (Figure S1). To account for possible solid and gas products formed during the chemical stability tests, we have additionally performed quantitative NMR analyses (Figure S3; see the Supplemental Information for more details) and determined that the fractions of BTMSA, BMCF3SA, and DMCF3SA remaining intact in the presence of 10 equiv. Li2O2 and KO2 at 80°C for three days are 97.8, 99.3, and 102.1%, respectively. Overall, the quantitative NMR studies showed that our synthesized solvents were highly stable in the chemical stability tests. We note that G4 also didn’t yield detectable degradation products under the same test conditions. In contrast, 14.1% of the DMSO sample decomposed to form dimethyl sulfone, as indicated by the new peak at ~3 ppm38 (Figure 3D). Additionally, recent reports39,40 have proposed electrolyte degradation due to the formation of singlet O2 (1O2) during Li-O2 battery operation. To investigate the reactivity of BTMSA, BMCF3SA, DMCF SA and DMSO with 1O2, 1O2 was generated by irradiating solutions containing these electrolyte components as well as the photosensitizer zinc tetraphenylporphyrin (ZnTPP) in a custom-made photoreactor (Figure S4A–G; component list in Table S2); the generation of 1O2 in these solutions was verified by detecting its emission at 1,270 nm (Figure S4H; see the Supporting Information for experimental details). 1H NMR spectra (Figure 3E) collected before and after irradiation reveal a new peak at ~2.94 ppm for the solution containing DMSO while no observable change for our synthesized compounds and G4 after 8 hours of irradiation, highlighting their stability toward 1O2.

Figure 3. Chemical stability of BTMSA, DMCF3SA, and BMCF3SA.

Figure 3.

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1H NMR analyses of the chemical stability of (A) BTMSA, (B) DMCF3SA, (C) BMCF3SA, and (D) DMSO. Teal and red spectra were obtained before and after the chemical stability test, in which the samples were mixed with 0.5 equiv. commercial Li2O2 and KO2 powders. The mixtures were stirred and maintained at 80°C for three days.

(E) 1H NMR analyses of the solutions containing 50 μL DMSO, G4, BTMSA, BMCF3SA, and DMCF3SA and 150 μM ZnTPP in 0.5 mL d-ACN before and after exposure to 1O2. Length of irradiation is indicated by “_#hr” followed by sample name. Note: the changing peaks at ~2.25 ppm are attributed to adventitious water.

We further confirmed that these electrolytes are chemically and electrochemically stable upon discharge in a real Li-O2 battery environment. Li-O2 cells with electrolytes containing 0.2 M LiTFSI in BTMSA, BMCF3SA, and DMCF3SA sandwiched by carbon paper with a gas diffusion layer (CP-GDL) cathode and a Li metal anode were fully discharged at 0.03 mA/cm2 with a voltage cutoff of 2.0 VLi. Cells containing electrolyte components with higher DNs – BTMSA (16.9 kcal/mol) and DMCF3SA (16.4 kcal/mol) – exhibited higher full discharge capacities (1.04 and 0.95 mAh/cm2, respectively, comparable to the full discharge capacities of Li-O2 cells employing similar CP-GDL electrodes and a G4-based electrolyte reported previously41) than the lower-DN compound, BMCF3SA (DN = 13.3 kcal/mol, full discharge capacity = 0.79 mAh/cm2). This observation agrees with the previously reported trend between higher-DN electrolytes and higher discharge capacities in Li-O2 batteries.24,26 XRD characterization of CP-GDL cathodes after full discharge showed Li2O2 as the discharge product (Figure S5A). After full discharge, the electrolytes were collected and analyzed by FTIR (Figure S5B), 1H (Figure S5C) NMR, and 19F NMR (Figure S5D), and compared to the pristine electrolytes. No perceivable change was observed in the FTIR or NMR spectra for all three electrolytes, indicating that these electrolytes are resistant to chemical degradation under full discharge conditions. Additionally, we performed pressure-tracking experiment to show that the ratios of electron and O2 consumption during galvanostatic discharge in the G4- and DMCF3SA-based cells were both highly close to the 2 e− per O2 ideality (Figure S6A–B). Furthermore, we quantified the yield of Li2O2 in the G4- and DMCF3SA-based cells using titration and UV-Vis spectroscopy (Figure S6C–D; see the Supplemental Information for more details). These experiments showed that the Li2O2 yield in the DMCF3SA cell was significantly higher than that of G4, 85.5 ± 2.7% vs. 77.0 ± 1.7% (Figure S6E; the yield was normalized by the total O2 consumption determined in the pressure-tracking experiments; the error bars represent one standard deviation based on three replicate trials for each electrolyte), indicating that our synthesized electrolyte DMCF3SA exhibited superior discharge stability over G4.

To examine the stability of these electrolytes against charging in Li-O2 batteries, we first examined the electrochemical oxidative stability of electrolytes containing 0.1 M LiTFSI in BTMSA, BMCF3SA, and DMCF3SA using potentiostatic measurements, cyclic voltammetry (CV), and linear sweep voltammetry (LSV). The potentiostatic measurements were performed under an oxygenated environment in a two-electrode electrochemical cell held at various potentials from 3.4 to 5.0 VLi for 3 hours each (Figure 4A). The electrochemical cell consists of a glass fiber separator impregnated with the electrolyte and sandwiched between stainless steel mesh (316), current collector, and Li foil. The same measurement was performed on DMSO- and G4-based electrolytes for comparison. The sulfamide- and sulfonamide-based electrolytes exhibited high stability against electrochemical oxidation (oxidative current < 5 μA, zoomed-in view in Figure 4B) at potentials ≤ 4.5 VLi, similar to the G4-based electrolyte (Figure S6F). In contrast, the cell containing DMSO showed 1 ~ 2 orders of magnitude greater oxidative current. At higher potentials (≥ 4.8 VLi, Figure 4C), sulfonamides with the electron-withdrawing −CF3 moiety, BMCF3SA, and DMCF3SA, exhibited considerably greater electrochemical oxidative stability (oxidative current < 20 μA) than the sulfamide BTMSA (oxidative current 50 ~ 220 μA), in excellent agreement with our computational prediction (Figure S1). The electrochemical oxidation stability of these three electrolytes were further tested against high-surface-area carbon electrodes as carbon-based electrodes commonly-used in aprotic Li-O2 batteries can participate in parasitic reactions, especially at high charging potential.42–45 We performed cyclic voltammetry (CV, 1 mV/s, 2.0 – 5.0 VLi, Figure S6G) tests in Ar and linear sweep voltammetry (LSV, 0.1 mV/s, from open circuit voltage to 5.0 V, Figure S6G inset) tests in O2 using CP-GDL as the working electrode. Under both Ar and O2, the BTMSA- and DMSO-based electrolytes exhibited increasing oxidative current at > 4.2 VLi on carbon electrodes. In contrast, the −CF3 containing compounds, BMCF3SA, and DMCF3SA, showed significantly improved oxidative stability (> 4.5 VLi). Notably, the oxidative current recorded in O2 for the DMSO-based electrolyte was 1 ~ 2 orders of magnitude higher than all other electrolytes at 5 VLi, as shown in Figure S6G inset.

Figure 4. Electrochemical oxidative stability of BTMSA, DMCF3SA, and BMCF3SA.

Figure 4.

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Potentiostatic electrochemical stability tests of electrolytes containing 0.1 M LiTFSI in BTMSA, DMCF3SA, and BMCF3SA compared to DMSO in an oxygenated environment on stainless steel (316) electrodes at (A) potential ≤ 4.5 VLi with enlarged view in (B) and (C) potential ≥ 4.8 VLi.

Differential electrochemical mass spectrometry (DEMS) has been used to assess the rechargeability of Li-O2 cells in various electrolytes employing carbonates, DMSO, and glymes, where deviations from the ideal 2 e− per O2 chemistry on discharge/charge are used to quantify the (electro)chemical instability of electrolytes. Generally, the rechargeability of Li-O2 cells in common electrolytes follows the order: glymes > DMSO > carbonates,10,46 where glyme- and carbonate-based electrolytes enable predominately O2 and CO2 evolution on charge,46 respectively (O2 evolution has also been observed in DMSO-based cells, although the rate of O2 evolution is lower than that of glyme10). Here we employed DEMS to detect gases evolved on charge under galvanostatic conditions following galvanostatic discharge of Li-O2 cells using the more electrochemically stable electrolytes, DMCF3SA, and BMCF3SA (Figure 5); the recorded O2 evolution rate is compared to the 2 e− per O2 ideality. While the voltage profiles of the DMSO-based cell increased gradually to ~4.2 VLi (Figure 5A), only O2 was detected, where the corresponding O2 evolution rate peaked at ~0.4 μmol/h then gradually decayed, yielding an overall evolution of 2.7 μmol O2 (5.2 e−/O2). In addition to O2, significant CO2 evolution was observed for the last 30% charging capacity (beginning at ~4.2 VLi and ~0.2 mAh/cm2). In contrast, the G4- and DMCF3SA-based cells showed a long plateau at ~4.2 VLi, which was accompanied by only O2 evolution with a steady O2 evolution rate of ~0.4 μmol/h (overall O2 evolution = 4.0 μmol for G4 and 4.1 μmol for DMCF3SA, both corresponding to 3.5 e−/O2), as shown in Figure 5B and 5C. However, as the charging potential increased above 4.3 VLi, O2 evolution was surpassed by CO2 evolution in both G4-and DMCF3SA-based cells (Figure 5B and 5C). While the charging potential of the BMCF3SA-based cell increased slowly, only O2 was detected upon charging of the cell at voltages below 4.2 V, where the first 70% charging capacity (~0.2 mAh/cm2) remained below 3.9 VLi, after which the voltage increased steadily to 4.3 VLi (Figure 5D). The O2 evolution rate of the BMCF3SA-based cell during early stage of charge was approximately twice as high as the DMSO-based cell (~0.6 vs. 0.3 μmol/h). As the potential of the BMCF3SA-based cell increased from 3.9 VLi to 4.3 VLi, O2 production increased again then eventually diminished, evolving 4.4 μmol O2 overall (3.2 e−/O2); CO2 evolution became dominant at > 4.2 VLi. We note that the overall e−/O2 values for the DMSO- and G4-based cells in this work are higher than those reported previously by McCloskey et al. (5.2 vs. 4.110 for DMSO, and 3.5 for G4 vs. 3.246 and 2.610 for DME), which is likely due to systematic instrumentation errors. Nevertheless, these DEMS results suggest that our new electrolytes exhibit superior O2 evolution efficiency and (electro)chemical stability compared to DMSO.

Figure 5. Differential electrochemical mass spectrometry (DEMS) analysis of Li-O2 cells employing DMSO-, G4-, DMCF3SA-, and BMCF3SA-based electrolytes.

Figure 5.

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Galvanostatic charging (0.03 mA/cm2) curves and gas evolution rates on charge of Li-O2 cells containing 0.2 M LiTFSI in (A) DMSO, (B) G4, (C) DMCF3SA, and (D) BMCF3SA.

Li-O2 cells employing 0.2 M LiTFSI in DMCF3SA as the electrolyte were subject to prolonged galvanostatic cycling tests (discharge at 0.03 mA/cm2, charge at 0.02 mA/cm2, and capacity cutoff at 0.1 mAh/cm2 unless otherwise noted). The discharge-charge profiles of select cycles are presented in Figure 6A. The electrolytes and positive electrodes of DMCF3SA-based cells were collected after select cycles and analyzed using 1H NMR (Figure 6B, Figure S7A, and Figure S7D, compared to spectra of pristine electrolytes in Figure S7F), 19F NMR (Figure S7B and Figure S7E), and FTIR (Figure S7C); the results are compared to DMSO- and G4-based cells cycled under the same galvanostatic conditions (cycling profiles in Figure S7). The 1H NMR analyses (Figure 7B) revealed clear new peaks for the DMSO-based electrolyte after the first cycle (capacity cutoff = 0.3 mAh/cm2); the signal attributable to dimethyl sulfone (DMSO2)38 significantly intensified after the 10th cycle (~2.9 ppm in Figure 6B and Figure S7D). The spectrum for G4 exhibited numerous new peaks in the range of 3.6~4.7 ppm (Figure 6B) as well as a clear peak attributable to formate47,48 (~8.1 ppm, Figure S7A) after 92 cycles. Our FTIR analyses also confirmed the presence of formate in G4-based electrolyte at ~1700 cm−1 (Figure S7C). In contrast, the DMCF3SA-based electrolyte collected after the 1st (capacity cutoff = 0.3 mAh/cm2), 5th, 25th, and 92nd cycles did not display new peaks in the 1H NMR (Figure 6B) and FTIR (Figure S7C) spectra, highlighting the superior stability of this electrolyte under prolonged cycling conditions. Additionally, 19F NMR analysis (Figure S7B and Figure S7E) of the DMCF3SA-based electrolyte showed a negligible change for the first 25 cycles; nonetheless, the spectrum collected after 92 cycles revealed a small amount of degradation products, likely resulting from parasitic reactions with the Li metal electrode and/or the oxygen electrode. Taken together, these data suggest that our electrolytes exhibit (electro)chemical stability superior to that of G4 and DMSO under prolonged cycling conditions and are promising for use in aprotic Li-O2 batteries.

Figure 6. Cycling stability of Li-O2 cells employing DMCF3SA-based electrolytes.

Figure 6.

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(A) Galvanostatic discharge (0.03 mA/cm2) and charge (0.02 mA/cm2) profiles of select cycles (1st, 5th, 10th, 25th, 50th, and 80th cycles) of a Li-O2 cell employing 0.2 M LiTFSI in DMCF3SA as the electrolyte.

(B) 1H NMR analyses on DMCF3SA-, G4-, and DMSO-based electrolytes collected after select cycles (denoted by “_cycle#”). Asterisk (*) indicates the signal of water from d-ACN solvent.

In summary, we present three electrolytes based on BTMSA, BMCF3SA, and DMCF3SA with enhanced chemical and electrochemical stability in aprotic Li-O2 batteries. These compounds were shown to be stable in the presence of commercial Li2O2 and KO2 powders as well as under galvanostatic, full discharge conditions, likely due to the suppressed solubility of discharge reaction intermediates (e.g., Li+O2–) resulting from low electrolyte donor numbers (DNs). In contrast, DMSO decomposes significantly under the same testing conditions. Additionally, BMCF3SA and DMCF3SA are considerably more stable against electrochemical oxidation (Vox > 4.5 VLi) than DMSO and BTMSA, which can be attributed to the electron-withdrawing effect of the −CF3 moiety. Differential electrochemical mass spectrometry (DEMS) measurements show O2 as the vastly predominant gas evolved on charge in Li-O2 cells employing sulfonamide-based electrolytes, which, notably, exhibit ~50% higher overall O2 evolution than the DMSO cell. Li-O2 cells employing the DMCF3SA-based electrolyte have been cycled for 90 times without capacity decay. The results presented in this study demonstrate that sulfamide- and sulfonamide-based electrolytes are promising for aprotic Li-O2 battery electrolytes. In addition, this work highlights the power of molecular design in the context of Li-O2 battery chemistry.

EXPERIMENTAL PROCEDURES

Chemical stability tests

A 10 mL microwave vial was charged with 0.5 mL sulfamide- or sulfonamide-based solvents a stir bar. After three cycles of freeze-pump-thaw to remove the air, the vial was transferred into the glove box. Then 0.5 equiv. Li2O2 and 0.5 equiv. KO2 were added into the vial. After the vial was sealed, it was moved out of the glove box and heated in an oil bath at 80°C for 3 days. The reaction mixture was cooled down and treated with d6-DMSO. The mixture was further centrifuged. The liquid layer was analyzed with 1H and 19F-NMR.

Electrochemical measurements

The potentiostatic oxidative stability tests (Figure 4) were conducted in an electrochemical cell consisting either a piece of stainless steel (316) mesh (D=12.7 mm) or a carbon paper with gas diffusion layer electrode (CP-GDL, Freudenberg H23C2, Fuel Cells Etc, D=12.7 mm) as the working electrode, one glass fiber separator (D=18 MM, Whatman®, Grade GF/A) impregnated with the electrolyte (0.1 M LiTFSI in the solvent of interest), and a lithium foil (D=15 mm, Chemetall, Germany). The cells were assembled in an Argon glove box (H2O < 0.1 ppm, O2 < 0.1 ppm, MBraun, USA). For tests conducted in an oxygenated environment, the assembled cell was transferred to a second glove box (H2O < 1 ppm, O2 < 1 %, MBraun, USA) and pressurized with dry O2 (99.994% purity, H2O < 2 ppm, Airgas, USA). In the potentiostatic tests, after holding the cell at open circuit voltage for one hour, a series of potentials were applied for three hours each: 3.4, 3.6, 3.8, 4.0, 4.2, 4.4, 4.5, 4.8, and 5.0 VLi while the current was recorded.

Galvanostatic discharge and charge (Figure 5 and Figure 6) were performed using an electrochemical cell consisting a CP-GDL, one glass fiber separator (D=18 MM, Whatman®, Grade GF/A) impregnated with the electrolyte (0.2 M LiTFSI in the solvent of interest), and a lithium foil (D=15 mm, Chemetall, Germany). All cells were assembled in an Ar glove box (H2O < 0.1 ppm, O2 < 0.1 ppm, MBraun, USA), pressurized with high-purity, dry O2 (99.994% purity, H2O < 2 ppm, Airgas, USA) in a second glove box (H2O < 1 ppm, O2 < 1 %, MBraun, USA).

All electrochemical tests were conducted employing a VMP3 potentiostat (BioLogic Science Instruments).

Supplementary Material

1

The Bigger Picture.

Lithium-oxygen (Li-O2) batteries can potentially transform energy storage and transportation with a several-fold increase in energy density over the state-of-the-art Li-ion batteries. The development of rechargeable Li-O2 batteries faces substantial challenges such as severe electrolyte instability against the highly reactive oxygen species, including superoxide, peroxide, and single oxygen, generated during Li-O2 battery operation. To date, the vast majority of studies in this field are based on electrolytes derived from a small set of well-studied, commercially available components (e.g., solvents such as tetraglyme, DMSO, etc. and salts such as LiTFSI). Though great progress has been made through optimization of such formulations, the use of physical organic chemistry principles to rationally design new molecular components may enable the discovery of electrolytes with stability profiles that cannot be achieved with existing formulations.

Highlights.

  • Designed electrolytes are highly resistant to chemical degradation by (su)peroxide

  • Sulfonamide-based electrolytes showed electrochemical oxidative stability at >4.2VLi

  • Sulfonamides showed superior cycling stability compared to tetraglyme and DMSO

ACKNOWLEDGMENTS

The authors would like to thank Samsung Advanced Institute of Technology (SAIT) for funding this research, Dr. Moungi Bawendi for his support in the photoluminescence spectroscopy experiment, and Graham Leverick for his assistance in the pressure tracking and DEMS experiments. S.F. gratefully acknowledges the Link Foundation for an Energy Fellowship. J.R.L. gratefully acknowledges the National Institutes of Health for a Postdoctoral Fellowship (1F32GM126913–01A1). C.F.P. was supported by the National Science Foundation Graduate Research Fellowship under Grant No. 1122374 and by the Center for Excitonics, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences under Award Number DE-SC0001088 (MIT). This research used resources of the National Energy Research Scientific Computing (NERSC) Center, a DOE office of Science User Facility supported by the Office of Science of the U.S. Department of Energy under Contract No. DE-AC02–5CH11231, and the Extreme Science and Engineering Discovery Environment (XSEDE), which is supported by National Science Foundation grant number ACI-1548562.

Footnotes

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DECLARATION OF INTERESTS

The authors have applied for a patent on this work.

REFERENCES AND NOTES

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