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

Full text of DOI:10.1016/j.chempr.2019.07.003

Article
Molecular Design of Stable Sulfamide- and
Sulfonamide-Based Electrolytes for Aprotic Li-
O2 Batteries
Shuting Feng, Mingjun Huang,
Jessica R. Lamb, ..., Collin F.
Perkinson, Jeremiah A. Johnson,
Yang Shao-Horn
fengs@mit.edu (S.F.)
jaj2109@mit.edu (J.A.J.)
HIGHLIGHTS
Designed electrolytes are highly
resistant to chemical degradation
by (su)peroxide
Sulfonamide-based electrolytes
showed electrochemical oxidative
stability at >4.2VLi
Sulfonamides showed higher
cycling stability than tetraglyme
and DMSO
Aprotic lithium-oxygen (Li-O2) batteries show great promise in energy storage and
transportation applications because of their high gravimetric energies, which
potentially represent a several-fold increase over Li-ion batteries. The stable and
reversible operation of Li-O2 batteries, however, is currently hindered by the
severe degradation of common electrolytes. Here, we show that sulfonamide-
based electrolytes, designed on the basis of physical organic chemistry principles,
can exhibit higher (electro)chemical stability than common electrolytes, such as
tetraglyme and DMSO.
Feng et al., Chem 5, 1–12
October 10, 2019 ª 2019 Elsevier Inc.
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Article
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
*
3,4
Yun Guang Zhu,3 Collin F. Perkinson,2 Jeremiah A. Johnson,2,6, and Yang Shao-Horn
*
SUMMARY
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 singlet oxygen,
generated during Li-O2 battery
operation. To date, the vast
majority of studies in this field
have been based on electrolytes
derived from a small set of well-
studied, commercially available
components (e.g., solvents such
as tetraglyme and DMSO and salts
such as lithium
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., lithium bis(trifluoromethane)sulfonimide [LiTFSI]), and exceptionally sta-
ble 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 cy-
cles in a Li-O2 cell when cycled at <4.2 VLi. This study provides guiding princi-
ples for the development of next-generation electrolyte components based on
sulfamides and sulfonamides.
INTRODUCTION
Aprotic lithium-oxygen (Li-O2) batteries show great promise in energy storage and
transportation applications because of their high gravimetric energies, which
potentially represent a 3- to 5-fold increase over Li-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 electro-
lyte components of well-established battery chemistries (e.g., Li ion), such as car-
bonates,5–8 glymes,9–11 dimethyl sulfoxide (DMSO),12–14 and N,N-dimethylforma-
mide (DMF),15 are not stable in the radical-rich, basic, nucleophilic, and oxidizing
environment of the oxygen electrode of Li-O2 batteries (Figure 1). Although refor-
mulation of classical electrolyte components, e.g., with high salt concentrations, has
led to significant stability improvements in some systems, the path toward practical
Li-O2 batteries will most likely require the rational molecular design of novel elec-
trolyte components.16 In an early example of such an approach, Nazar and co-
workers17 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, ketone-based18 and pivalate-based19 electro-
lyte 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.
bis(trifluoromethane)sulfonimide
[LiTFSI]). Although 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.
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Figure 1. Dominant Degradation Mechanisms of Carbonate-, Ether-, and Sulfoxide-Based
Electrolytes and the Molecular Design of Stable Sulfamide- and Sulfonamide-Based Solvents
BTMSA (Top), DMCF3SA (Middle), and BMCF3SA (Bottom) for Aprotic Li-O2 Batteries
RESULTS AND DISCUSSION
Herein, we report three compounds—N-butyl-N,N0,N0-trimethylsulfamide (BTMSA),
N,N-dimethyl-trifluoromethanesulfonamide (DMCF3SA), and N-butyl-N-methyl-tri-
fluoromethanesulfonamide (BMCF3SA)—that are promising for use in aprotic
Li-O2 batteries (Figure 1). These compounds are polar aprotic liquids at room tem-
perature and are capable of dissolving reasonably high amounts of common Li
salts, such as lithium bis(trifluoromethane)sulfonimide (LiTFSI). BTMSA has been
considered as an electrolyte component for Li 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 electro-
chemically 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-sta-
bility relationships obtained from our previous works,16,22 we designed the struc-
tures of these electrolyte components to be devoid of (1) vulnerable C–H bonds
for hydrogen and proton removal, (2) electrophilic centers susceptible to nucleo-
philic substitution, and (3) highly electron-donating functional groups vulnerable
1Department of Chemical Engineering,
Massachusetts Institute of Technology,
to electrochemical oxidation. In this study, these three compounds were shown
Cambridge, MA 02139, USA
both experimentally and computationally (Figure S1) to exhibit exceptional stability
2Department of Chemistry, Massachusetts
Institute of Technology, Cambridge, MA 02139,
USA
toward a variety of harsh conditions encountered in the Li-O2 cathode environment.
For example, all three compounds were compatible with reactive species such as
peroxide, superoxide, and singlet O2, whereas the trifluoromethylsulfonamides
3Research Laboratory of Electronics,
BMCF3SA and DMCF3SA displayed enhanced electrochemical oxidative stability
due to the electron-withdrawing CF3 group. The latter could be cycled >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. All together, these results
highlight the power of rational molecular design for electrolyte discovery for Li-O2
batteries.
Massachusetts Institute of Technology,
Cambridge, MA 02139, USA
4Department of Mechanical Engineering,
Massachusetts Institute of Technology,
Cambridge, MA 02139, USA
5These authors contributed equally
6Lead Contact
*Correspondence: fengs@mit.edu (S.F.),
jaj2109@mit.edu (J.A.J.)
We began by using
groups that may be stable in the Li-O2 cathode environment. Sulfamides and
a
computational framework16 to identify functional
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Figure 2. Donor Number and Conductivity of BTMSA, DMCF3SA, and BMCF3SA
(A) The measured 23Na NMR chemical shifts of 20 mM NaTFSI in BTMSA, DMCF3SA, and BMCF3SA are compared with those of DMSO, DMF, DME, and
PC (400 MHz). The 23Na signal from the internal reference, 0.5 M NaClO4 in H2O, was set to 0 ppm.
(B) A trend (dashed) line correlating the donor numbers (DNs) of DMSO (29.8 kcal/mol31), DMF (26.6 kcal/mol31), DME (20.2 kcal/mol30), and PC
(15.1 kcal/mol32) and their measured relative 23Na NMR shifts was used for estimating the DNs of BTMSA, DMCF3SA, and BMCF3SA, which were
determined to be 16.9, 16.4, and 13.3 kcal/mol, respectively.
(C) Conductivities of solutions containing 0.1 M LiTFSI in BTMSA, DMCF3SA, and BMCF3SA are compared with those of G4 as a commercial reference at
various temperatures.
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 chlo-
ride and N-butylmethylamine in 90% yield, whereas 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 ( 3 ) are provided in Table S1.
The donor number (DN) ofan electrolytecomponent,23 which isa quantitative descriptor
of Lewis basicity, can influence the solubility24–27 and lifetime24 of intermediates such as
LiO2,24–26 as well as the discharge product morphology24,27,28 and capacity of Li-O2 bat-
teries.24–27 The DNs of BTMSA, BMCF3SA, and DMCF3SA were estimated by 23Na
NMR.29 Solutions of 20 mM sodium bis(trifluoromethane)sulfonimide (NaTFSI) were
prepared in BTMSA, BMCF3SA, and 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 sulfon-
amide-based electrolytes indicate weaker interactions between Na+ and the solvation
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shell in these electrolytes.29,30 Using the known DNs of DMSO (29.8 kcal/mol31), DMF
(26.6 kcal/mol31), DME (20.2 kcal/mol30), and PC (15.1 kcal/mol32) and their 23Na NMR
shifts,33 we estimated the DNs of BTMSA, DMCF3SA, and BMCF3SA 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 lower superoxide solubility24,26 and thus higher chemical stability than
high-DN solvents such as DMSO and DMF.
Because 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 were
compared with those of tetraglyme (G4, DN = 16.6 kcal/mol30) as a reference
because of its relative stability against oxygen and its reduction products (Figure 2C).
The BTMSA-LiTFSI solution exhibited conductivity approximately 2-fold greater
than that of G4, whereas DMCF3SA- and BMCF3SA-LiTFSI solutions had conductiv-
ities ꢀ2- and 5-fold lower than that of BTMSA, respectively. These values can be
rationalized by considering the dielectric constants ( 3 ), 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 versus 7.79;34 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
anion35,36 (Figure S2), resulting in lower charge carrier concentration and conductiv-
ity. Furthermore, we note that BMCF3SA not only solvates Li+ more weakly than
DMCF3SA, leading to lower charge carrier concentration, but also has higher viscos-
ity (Table S1), both of which contributed to its lower conductivity than DMCF3SA.
The chemical stabilities of BTMSA, DMCF3SA, and BMCF3SA were evaluated under
conditions mimicking the oxygen electrode of aprotic Li-O2 batteries according to a
previously established protocol.16,22 These electrolyte components were combined
with 0.5 equiv lithium peroxide (Li2O2) and KO2 powders, and the resulting mixtures
were stirred at 80ꢁC for 3 days. The lack of appreciable change in the 1H NMR
spectra collected before and after the exposure to Li2O2 and KO2 (Figures 3A–3C)
indicate that these compounds are highly resistant to chemical degradation by
peroxide and superoxide, in good agreement with our computational analyses (Fig-
ure S1). To account for possible solid and gas products formed during the chemical
stability tests, we additionally performed quantitative NMR analyses (Figure S3; see
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 80ꢁC for 3 days were 97.8%, 99.3%, and 102.1%, respectively.
Overall, the quantitative NMR studies showed that these solvents were highly stable
in the chemical stability tests. We note that G4 also did not yield detectable degra-
dation products under the same test conditions. In contrast, 14.1% of the DMSO
sample decomposed to form dimethyl sulfone (DMSO2), as indicated by the new
resonance at ꢀ3 ppm37 (Figure 3D). Additionally, recent reports38,39 have proposed
electrolyte degradation due to the formation of singlet O2 (1O2) during Li-O2 battery
operation. To investigate the reactivity of BTMSA, BMCF3SA, DMCF3SA, and DMSO
with 1O2, we generated 1O2 by irradiating solutions containing these electrolyte
components as well as the photosensitizer zinc tetraphenylporphyrin (ZnTPP) in a
custom-made photoreactor (Figures S4A–S4G; component list in Table S2); we veri-
1
fied the generation of O2 in these solutions by detecting its emission at 1,270 nm
4
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Figure 3. Chemical Stability of BTMSA, DMCF3SA, and BMCF3SA
(A–D) 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, respectively, in which the samples were mixed with 0.5 equiv commercial Li2O2 and KO2 powders. The mixtures were
stirred and maintained at 80ꢁC for 3 days.
(E) 1H NMR analyses of the solutions containing 50 mL DMSO, G4, BTMSA, BMCF3SA, and DMCF3SA and 150 mM ZnTPP in 0.5 mL d-ACN before and
after exposure to 1O2. Length of irradiation is indicated by ‘‘_#h’’ following the sample name.
(Figure S4H; see Supplemental Information for experimental details). 1H NMR
spectra (Figure 3E) collected before and after irradiation revealed a new resonance
at ꢀ2.94 ppm for the solution containing DMSO, whereas no observable change was
observed for our synthesized compounds or G4 after 8 h of irradiation, highlighting
their stability toward 1O2.
We further confirmed that these electrolytes are chemically and electrochemically
stable upon discharge in a real Li-O2 battery environment. Li-O2 cells with electro-
lytes containing 0.2 M LiTFSI in BTMSA, DMCF3SA, and BMCF3SA 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
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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
previously40) 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 X-ray diffraction (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 Fourier-transform
infrared spectroscopy (FTIR) (Figure S5B), 1H NMR (Figure S5C), and 19F NMR (Fig-
ure 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 a pressure-tracking experiment to show that the ratios
of electron (eꢂ) and O2 consumption during galvanostatic discharge in the G4-
and DMCF3SA-based cells were both highly close to the 2 eꢂ/O2 ideality (Figures
S6A and S6B). Furthermore, we quantified the yield of Li2O2 in the G4- and
DMCF3SA-based cells by using titration and UV-visible spectroscopy (Figures S6C
and S6D; see Supplemental Information for more details). These experiments
showed that the Li2O2 yield in the DMCF3SA cell was significantly higher
(85.5% G 2.7%) than that of G4 (77.0% G 1.7%) (Figure S6E; the yield was normal-
ized 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 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, DMCF3SA, and BMCF3SA by using potentiostatic mea-
surements, cyclic voltammetry (CV), and linear sweep voltammetry (LSV). The poten-
tiostatic measurements were performed under an oxygenated environment in a
2-electrode electrochemical cell held at various potentials from 3.4 to 5.0 VLi for
3 h each (Figure 4A). The electrochemical cell consisted of a glass fiber separator
impregnated with the electrolyte and sandwiched between a 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 sulfon-
amide-based electrolytes exhibited high stability against electrochemical oxidation
(oxidative current < 5 mA; 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 oxidative current that was 1ꢀ2 orders of magnitude greater. At higher
potentials (R4.8 VLi; Figure 4C), sulfonamides with the electron-withdrawing –CF3
moiety, BMCF3SA and DMCF3SA, exhibited considerably greater electrochemical
oxidative stability (oxidative current < 20 mA) than the sulfamide BTMSA (oxidative
current 50–220 mA), in excellent agreement with our computational prediction (Fig-
ure S1). The electrochemical oxidation stability of these three electrolytes was
further tested against high-surface-area carbon electrodes given that carbon-based
electrodes commonly used in aprotic Li-O2 batteries can participate in parasitic
reactions, especially at high charging potential.41–44 We performed CV (1 mV/s,
2.0–5.0 VLi; Figure S6G) tests in Ar and LSV (0.1 mV/s, from open-circuit voltage to
5.0 V; Figure S6G inset) tests in O2 by 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
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Figure 4. Electrochemical Oxidative Stability of BTMSA, DMCF3SA, and BMCF3SA
Potentiostatic electrochemical stability tests of electrolytes containing 0.1 M LiTFSI in BTMSA,
DMCF3SA, and BMCF3SA are compared with DMSO in an oxygenated environment on stainless-
steel (316) electrodes at potential % 4.5 VLi (A); enlarged views at potential R 4.8 VLi are shown in (B)
and (C).
stability (>4.5 VLi). Notably, the oxidative current recorded in O2 for the DMSO-
based electrolyte was 1ꢀ2 orders of magnitude higher than that of all other electro-
lytes at 5 VLi, as shown in the Figure S6G inset.
Differential electrochemical mass spectrometry (DEMS) has been used to assess the
rechargability of Li-O2 cells in various electrolytes employing carbonates, DMSO,
and glymes, where deviations from the ideal 2 eꢂ per O2 stoichiometry on discharge
and charge are used to quantify the (electro)chemical instability of electrolytes.10,45
Generally, the rechargeability of Li-O2 cells in common electrolytes follows the order
of glymes > DMSO > carbonates,10,45 where glyme- and carbonate-based
electrolytes enable predominately O2 and CO2 evolution on charge,45 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 after galvanostatic discharge of
Li-O2 cells by using the more electrochemically stable electrolytes, DMCF3SA
and BMCF3SA (Figure 5); the recorded O2 evolution rate was compared to the
2 eꢂ/O2 ideality. Whereas the voltage profiles of the DMSO-based cell increased
gradually to ꢀ4.2 VLi (Figure 5A), only O2 was detected, and the corresponding
O2 evolution rate peaked at ꢀ0.4 mmol/h and then gradually decayed, yielding an
overall evolution of 2.7 mmol O2 (5.2 eꢂ/O2). In addition to O2, significant CO2 evo-
lution 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 mmol/h (overall O2 evolution = 4.0 mmol for G4 and
4.1 mmol for DMCF3SA, both corresponding to 3.5 eꢂ/O2), as shown in Figures 5B
and 5C. As the charging potential increased above 4.3 VLi, however, O2 evolution
was surpassed by CO2 evolution in both G4- and DMCF3SA-based cells (Figures
5B and 5C). Whereas 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
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Figure 5. Differential Electrochemical Mass Spectrometry Analysis of Li-O2 Cells Employing
DMSO-, G4-, DMCF3SA-, and BMCF3SA-Based Electrolytes
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.
the BMCF3SA-based cell during the early stage of charge was approximately twice
as high as the DMSO-based cell (ꢀ0.6 versus 0.3 mmol/h). As the potential of the
BMCF3SA-based cell increased from 3.9 to 4.3 VLi, O2 production increased again
and then eventually diminished, evolving 4.4 mmol O2 overall (3.2 eꢂ/O2); CO2 evo-
lution became dominant at >4.2 VLi. We note that the overall eꢂ/O2 values for the
DMSO- and G4-based cells in this work are higher than those reported previously
by McCloskey et al. (5.2 versus 4.110 for DMSO and 3.5 for G4 versus 3.245 and
2.610 for DME), which is most likely due to systematic instrumentation errors. Never-
theless, these DEMS results suggest that our new electrolytes exhibit higher O2 evo-
lution efficiency and (electro)chemical stability than DMSO.
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 electro-
lytes and positive electrodes of DMCF3SA-based cells were collected after select
cycles and analyzed by 1H NMR (Figures 6B, S7A, and S7D; spectra of pristine
electrolytes are shown in Figure S7F), 19F NMR (Figures S7B and S7E), and FTIR
(Figure S7C); the results were compared to those of DMSO- and G4-based cells
cycled under the same galvanostatic conditions (cycling profiles in Figure S7). The
1H NMR analyses (Figure 6B) revealed clear new resonances for the DMSO-based
electrolyte after the first cycle (capacity cutoff = 0.3 mAh/cm2); the signal attributable
to DMSO2 significantly intensified after the 10th cycle (ꢀ2.9 ppm in Figures 6B
37
and S7D). The spectrum for G4 exhibited numerous new resonances in the range
of 3.6–4.7 ppm (Figure 6B) as well as a clear peak attributable to formate46,47 (ꢀ8.1
ppm, Figure S7A) after 92 cycles. FTIR analyses also confirmed the presence of
formate in G4-based electrolyte at ꢀ1,700 cmꢂ1 (Figure S7C). In contrast, the
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Figure 6. Cycling Stability of Li-O2 Cells Employing DMCF3SA-Based Electrolytes
(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#’’). The asterisk (*) indicates the signal of water from d-ACN solvent.
DMCF3SA-based electrolyte collected after the 1st (capacity cutoff = 0.3 mAh/cm2),
1
5th, 25th, and 92nd cycles did not display new peaks in the H NMR (Figure 6B) and
FTIR (Figure S7C) spectra, highlighting the superior stability of this electrolyte under
prolonged cycling conditions. Additionally, 19F NMR analysis (Figures S7B and 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 product, most 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.
In summary, we present three electrolytes based on BTMSA, DMCF3SA, and BMCF3SA
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 pow-
ders as well as under galvanostatic, full discharge conditions most likely because of the
suppressed solubility of discharge reaction intermediates (e.g., Li+O2ꢂ) resulting from
low electrolyte DNs. In contrast, DMSO decomposed significantly under the same
testing conditions. Additionally, BMCF3SA and DMCF3SA were considerably more sta-
ble against electrochemical oxidation (Vox > 4.5 VLi) than DMSO and BTMSA, which can
be attributed to the electron-withdrawing effect of the –CF3 group. DEMS measure-
mentsshowedO2 asthe vastlypredominantgasevolvedoncharge inLi-O2 cellsemploy-
ing sulfonamide-based electrolytes, which notably exhibited ꢀ50% higher overall O2
evolution than the DMSO cell. Li-O2 cells employing the DMCF3SA-based electrolyte
were cycled 90 times without capacity decay. The results presented in this study demon-
strate thatsulfamide-and sulfonamide-basedelectrolytesare promisingfor aprotic Li-O2
battery electrolytes. In addition, this work highlightsthe 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 with a stir bar. After three cycles of freezing, pumping, and thawing to
Chem 5, 1–12, October 10, 2019
9

Please cite this article in press as: Feng et al., Molecular Design of Stable Sulfamide- and Sulfonamide-Based Electrolytes for Aprotic Li-O2 Bat-
teries, Chem (2019), https://doi.org/10.1016/j.chempr.2019.07.003
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 80ꢁC for 3 days. The reaction mixture was
cooled down and treated with d6-DMSO. The mixture was further centrifuged. The
1
liquid layer was analyzed with H and 19F NMR.
Electrochemical Measurements
The potentiostatic oxidative stability tests (Figure 4) were conducted in an electro-
chemical cell consisting of either a piece of stainless-steel (316) mesh (D =
12.7 mm) or a carbon paper with gas diffusion layer electrode (CP-GDL, Freuden-
berg 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 Li foil (D = 15 mm, Chemetall, Ger-
many). The cells were assembled in an Ar glove box (H2O < 0.1 ppm, O2 < 0.1
ppm, MBraun, USA). For tests conducted in an oxygenated environment, the assem-
bled 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 the cell was held at open circuit voltage for 1 h, a series
of potentials were applied for 3 h 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 (Figures 5 and 6) were performed with an elec-
trochemical cell consisting of a CP-GDL, one glass fiber separator (D = 18 MM, What-
man, Grade GF/A) impregnated with the electrolyte (0.2 M LiTFSI in the solvent of
interest), and a Li 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) and 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 with a VMP3 potentiostat (BioLogic Sci-
ence Instruments).
SUPPLEMENTAL INFORMATION
Supplemental Information can be found online at https://doi.org/10.1016/j.chempr.
2019.07.003.
ACKNOWLEDGMENTS
The authors would like to thank the Samsung Advanced Institute of Technology for
funding this research, Dr. Moungi Bawendi for his support in the photolumines-
cence 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 a National Science Foundation (NSF) graduate research fellow-
ship under grant no. 1122374 and by the Center for Excitonics, an Energy Frontier
Research Center funded by the US Department of Energy (DOE) Office of Science
Basic Energy Sciences Program under award no. DE-SC0001088 (Massachusetts
Institute of Technology). This research used resources of the National Energy
Research Scientific Computing Center, a DOE Office of Science User Facility sup-
ported under contract no. DE-AC02-5CH11231, and the Extreme Science and En-
gineering Discovery Environment, which is supported by NSF grant no. ACI-
1548562.
10
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Please cite this article in press as: Feng et al., Molecular Design of Stable Sulfamide- and Sulfonamide-Based Electrolytes for Aprotic Li-O2 Bat-
teries, Chem (2019), https://doi.org/10.1016/j.chempr.2019.07.003
AUTHOR CONTRIBUTIONS
Conceptualization, S.F., M.H., J.A.J., and Y.S.-H.; Methodology, S.F., M.H., W.Z.,
J.R.L., R.T., Y.Z., Y.G.Z., C.F.P., J.A.J., and Y.S.-H.; Investigation, S.F., M.H.,
W.Z., J.R.L., R.T., Y.Z., Y.G.Z., and C.F.P.; Writing – Original Draft, S.F. and
M.H.; Writing – Review & Editing, S.F., M.H., W.Z., J.R.L., R.T., C.F.P., J.A.J., and
Y.S.-H.; Funding Acquisition and Supervision, J.A.J. and Y.S.-H.
DECLARATION OF INTERESTS
The authors have applied for a patent on this work.
Received: March 13, 2019
Revised: June 19, 2019
Accepted: July 3, 2019
Published: July 25, 2019
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