Ajit Kumar*a,
Frederick Ntia,
Jenny Sun
a,
Mahin Malekia,
Steve Rowlandsb,
Paul M. Bayleyb,
Maria Forsyth
*a and
Patrick C. Howlett*a
aInstitute for Frontier Materials(IFM), Deakin University, Burwood, VIC 3125, Australia. E-mail: ajit.kumar@deakin.edu.au; maria.forsyth@deakin.edu.au; patrick.howlett@deakin.edu.au
bLi-S Energy, Deakin University, 75 Pigdons Rd, Waurn Ponds, Victoria 3216, Australia
First published on 7th April 2025
Despite the potential for a greater energy density than lithium-ion batteries, polysulphide dissolution, the polysulphide shuttle effect, and lithium metal instability impede the commercialization of lithium–sulfur (Li–S) batteries. To overcome these obstacles, this study investigates ionic liquids (ILs) as electrolytes, with a particular emphasis on mixed-anion ILs and high concentrations of lithium salt. As demonstrated by undetectable levels in Raman and UV spectroscopy, our results demonstrate that trimethyl-isobutyl phosphonium (P111i4FSI) with 30 mol% lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) efficiently inhibits polysulphide dissolution. With a specific capacity of 625 mA h g−1 (based on sulphur) and a 60% capacity retention after 200 cycles, this electrolyte dramatically enhances Li–S battery performance. These findings show how high-concentration IL electrolytes may stabilise lithium interfaces and reduce polysulfide-related problems, bringing Li–S battery technology closer to real-world uses.
Broader contextAdvanced energy storage systems that are scalable, sustainable, and efficient are required due to the pressing worldwide transition towards carbon-neutral technology and renewable energy. Because of the availability of sulphur and its high theoretical energy density, lithium–sulfur (Li–S) batteries have become a viable option for next-generation energy storage. However, problems including the shuttle effect, polysulphide dissolving, and lithium metal instability have prevented them from being widely used since they result in decreased cycle life and capacity fading.This work advances the field by demonstrating how ionic liquid (IL)-based electrolytes may be utilised to solve these important challenges. These IL electrolytes greatly enhance the electrochemical performance of Li–S batteries by stabilising the lithium metal contact and inhibiting polysulphide dissolution to undetectable levels. The study also identifies a new electrolyte design strategy that strikes a compromise between strong ionic conductivity and chemical stability, providing a workable way to increase battery efficiency and longevity. By bridging the gap between laboratory-scale developments and the real-world implementation of Li–S batteries, these discoveries enhance energy and environmental science. The enhanced performance made possible by IL electrolytes advances Li–S technology towards commercialisation and aids in the worldwide shift to sustainable energy systems for grid storage, electric car integration, and renewable energy integration. |
In Li–S batteries, the discharge process involves the reduction of sulfur (S8) to lithium sulfide (Li2S). This happens through a two-voltage plateau mechanism where intermediate polysulfides (Li2Sn, n ≥ 3) are formed and dissolved in the electrolyte before being reduced to insoluble lithium-sulfides (Li2S2 and Li2S).3,4 When the battery is charged, the process is reversed: lithium sulfides are oxidized back into sulfur, again forming polysulfide intermediates during the reaction. This reversible reaction gives Li–S batteries a high energy density capability.
Despite the promise of Li–S batteries, they have not yet been commercialized due to the major challenge of the “polysulfide shuttle” (PS) effect. During cycling, the soluble polysulfides formed at the cathode can dissolve in the electrolyte and travel to the anode, where they can react with the lithium metal anode. This so-called shuttle effect of polysulfides leads to unwanted side reactions, causing the rapid loss of active material, capacity fading, and poor coulombic efficiency. The deposition of insoluble lithium sulfide (Li2S) at the anode-electrolyte interface further contributes to these issues by forming a barrier that hinders lithium-ion transport.5,6
In addition to the polysulfide shuttle effect, lithium deposition at the metal anode remains a significant challenge in lithium–sulfur (Li–S) batteries. In electrolytes containing polysulfides, lithium growth often forms dendritic/mossy structures.7,8 This results in the anode experiencing volume expansion, which compromises mechanical stability and shortens the battery's cycle life.
Although Li–S batteries hold the promise of achieving high theoretical energy densities, practical issues like unstable anode structures hinder their commercialization. Addressing the challenges of lithium deposition and polysulfide interactions is crucial for overcoming these limitations and fully realizing the commercial viability of Li–S technology.
The choice of electrolyte can play a significant role in the stability of Li–S batteries. The liquid electrolytes commonly investigated in the extensive research of these devices are not realistically practical for large-scale manufacture and deployment due to safety and toxicity-related issues. Furthermore, the electrolytes used in conventional LIBs, such as carbonate-based solvents, are not compatible with Li–S chemistry.9 Polysulfides react with carbonate solvents, resulting in byproducts that degrade both the electrolyte and the active materials over time.10
To avoid these reactions, ether-based solvents such as dioxolane (DOL) and dimethoxyethane (DME) are commonly used in Li–S cells.11 These solvents are better at stabilizing polysulfides without causing unwanted side reactions. One popular electrolyte formulation is 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in a 1:
1 mixture of DOL and DME, with a 1% weight/weight addition of lithium nitrate (LiNO3). The LiNO3 additive helps passivate the lithium metal anode, preventing the polysulfides from reacting with it.12 However, this protection is lost once the LiNO3 is consumed, and the polysulfide shuttle effect resumes.13,14 Furthermore, the volatile and potentially explosive nature of LiNO3.15,16 presents additional safety concerns, making it challenging to manufacture large-scale Li–S cells with this electrolyte.
Controlling the dissolution of polysulfide intermediates is essential for extending the cycle life of Li–S batteries while maintaining high energy density. Electrolytes with a lower ability to dissolve polysulfides are preferred because they reduce the shuttle effect.6,17 However, achieving low polysulfide solubility is difficult due to the complex nature of sulfur and lithium polysulfides. Sulfur (S8) is hydrophobic and only dissolves in non-polar solvents like benzene, while the fully reduced product, Li2S, is hydrophilic and only dissolves in highly polar solvents like water.18 The intermediate polysulfides (Li2Sn) exhibit varying degrees of polarity depending on their chain length, making it hard to design an electrolyte that works well with all forms of polysulfides.
In addition to controlling polysulfide solubility, the electrolyte must be a good solvent for the lithium salt and must be able to form a stable passivation layer on both the anode and cathode; i.e. a stable, homogeneous SEI/CEI on the electrodes.
Previous research has explored the use of ionic liquids (ILs) as electrolytes in Li–S batteries. Watanabe's group studied the solubility of sulfur and lithium polysulfides (Li2Sn, 4 ≤ n ≤ 8) in ionic liquids with different anions and found that ILs with bis(trifluoromethanesulfonyl)imide (TFSI) anions were effective in reducing polysulfide dissolution.19 These TFSI-based ILs showed low sulfur solubility (about 10 mM atomic sulfur), making them suitable candidates for Li–S batteries, with the best results obtained for an electrolyte composition of 0.5 mol kg−1 LiTFSI in N-methyl-N-propylpyrrolidinium bis(trifluoromethane sulfonyl)imide provide discharge capacity of 600 mA h g−1 at the 50th cycle. The other IL-based electrolyte systems performed poorly due to high polysulfide solubility, side reactions, and slow mass transport.20
In other work, ILs with the smaller fluor sulfonyl imide (FSI) anion did not perform well in Li–S cells;19,21 despite effectively suppressing polysulfide dissolution, they caused significant capacity loss and high overpotential within the first 10 cycles. This was likely due to excessive decomposition of the FSI anion, resulting in an insulating layer on the sulfur cathode with concomitant charge transfer resistance. As a result, TFSI-based ILs have been considered the most promising for Li–S batteries due to their ability to balance low polysulfide solubility and stable cycling performance.22
Several individual challenges need to be addressed to enhance the performance of Li–S batteries and make them commercially viable; these include the control of polysulfide dissolution, the ability to cycle high-capacity lithium anodes without dendrite formation, and discovering an electrolyte that can withstand the highly reactive nature of both the metal Li anode and reactive polysulfide intermediates and form robust, low ionic-resistance passivating layers on both electrodes. The electrolyte is a key component in solving these challenges. Our previous work has shown that high, near-saturation concentrations of LiFSI in various ionic liquids can support high-rate cycling of Li metal anodes as well as high-voltage layered oxide cathodes.23 We have also shown that certain IL chemistries and compositions exhibit extremely low solubility for intermediate polysulfides (Li2Sn, 4 ≤ n ≤ 8).24
In this work, we further show that IL electrolytes based on trimethyl, isobutyl phosphonium [P111i4] [FSI] with the addition of LiTFSI dissolve polysulfides at such low levels that their concentration in the solution is typically below 1 mM sulfur, or even undetectable using UV-vis and Raman spectroscopy. This ultra-low solubility makes them ideal for suppressing the polysulfide shuttle effect in Li–S cells. Additionally, these electrolytes form a low-impedance passivating layer on both the cathode and anode, further enhancing cell performance by limiting polysulfide dissolution.
This discovery marks a significant advancement in Li–S battery technology. By combining ultra-low polysulfide solubility with stable interface formation, these electrolytes enable a quasi-solid-state sulfur redox process, leading to improved cycle life, and overall better performance for Li–S batteries.
After 3 hours of stirring, the color intensity of all the solutions, except for the pure ionic liquid, decreased. The time it took for the color to disappear was inversely proportional to the LiTFSI concentration. After 48 hours, all solutions became colorless, except for the pure ionic liquid. This color disappearance could be due to the reverse reaction, where polysulfides convert back into elemental sulfur and Li2S. Even a small amount of polysulfides can degrade battery performance if present in the electrolyte for more than an hour, thus ideally, we want to decrease the concentration to negligible values.
To further ensure that polysulfide formation was fully suppressed in our electrolyte, we further tested solutions with LiTFSI concentrations above 20 mol%. Fig. 2b shows the color changes in solutions with 0, 30, 40, 45, and 50 mol% LiTFSI, both before and after adding Li2S and elemental sulfur powder at 50 °C. Interestingly, none of these solutions, except for the pure ionic liquid, showed any color change, even after 90 days, suggesting no significant polysulfide dissolution occurred in electrolytes with 30 mol% or more LiTFSI. Raman spectroscopy, known to be a highly sensitive technique for detecting polysulfides, was used to confirm that even trace amounts of polysulfides were not forming in the electrolyte. Polysulfides can be identified by their characteristic Raman peaks in the 400–500 cm−1 range, which correspond to the S–S stretch vibration of the polysulfide molecules.25–27 Fig. 2c illustrates the sample preparation process for this test. We examined two samples: one with pure ionic liquid (neat IL) saturated with lithium polysulfides and another with 30 mol% LiTFSI in P111i4FSI, also saturated with lithium polysulfides. To ensure polysulfide formation, we stirred the mixtures of elemental sulfur and Li2S powder for one week at 50 °C, giving ample time for the polysulfides to form and dissolve. The solutions were then allowed to settle for two days. Heating the mixtures to 50 °C provided the necessary thermal energy to promote polysulfide formation. Afterward, the mixtures were centrifuged to separate any unreacted sulfur or Li2S from the liquid part of the solution. The supernatant liquid was transferred into transparent capillary tubes, which were sealed to prevent contamination from moisture and air during the Raman spectroscopy measurements. The Raman spectra for these samples are shown in Fig. 2d–i. For the sample with pure ionic liquid (0 mol% LiTFSI), the Raman spectra showed clear, sharp peaks at 438 and 474 cm−1, corresponding to dissolved lithium polysulfides. These peaks were detected even with a low laser power of 0.85 mW, which indicates a large quantity of dissolved polysulfides in the solution. In contrast, the Raman spectra of the colorless solution with 30 mol% LiTFSI did not show any detectable polysulfide peaks, even when illuminated with a much higher laser power of 17 mW (20 times higher than the power used for the pure ionic liquid sample). This suggests that the amount of dissolved polysulfides in the 30 mol% LiTFSI solution was extremely low or absent. For comparison, we also examined a blank electrolyte containing 30 mol% LiTFSI in P111i4FSI without any added sulfur or Li2S. The results showed no significant Raman peaks in this blank sample either. These findings indicate that high concentrations of LiTFSI in the P111i4FSI ionic liquid (30 mol% and above) effectively suppress the dissolution of lithium polysulfides. We can thus conclude that electrolytes based on LiTFSI in P111i4FSI with 30 mol% or more LiTFSI significantly reduce polysulfide formation and dissolution which should lead to the improvement of the overall stability and performance of Li–S batteries.
We then analyzed the Raman spectra of the mixed anion electrolyte, which contained 30 mol% LiTFSI in P111i4FSI and was saturated with lithium polysulfides. These results, along with the spectra for the blank electrolyte (without any lithium polysulfides), are presented in Fig. 3c and d, with zoomed-in views in Fig. 3g and h. After comparing the spectra from Fig. 3f, g, and h, we can conclude that the solubility of lithium polysulfides in the 30 mol% LiTFSI in P111i4FSI electrolyte is extremely low such that even the highly sensitive Raman spectroscopy could not detect any species when operating at full laser power (100%). In addition to Raman spectroscopy, we also used UV spectroscopy (UV-2600, SHIMADZU) to further analyze the solubility of polysulfides in both the neat P111i4FSI and the 30 mol% LiTFSI/P111i4FSI mixture. Fig. 3i shows the UV absorption spectra for these solutions. While UV-VIS spectra are generally challenging for determining the exact composition of polysulfide species, previous research has identified the absorption bands for various polysulfides: 490–500 nm for S82−, 450–470 nm for S62−, and around 420 nm for S42−.20
In Fig. 3i the black-colored graph represents the UV spectrum for the 30 mol% LiTFSI in P111i4FSI solution, saturated with Li2S8 and other lithium polysulfides (Li2Sn, n ≥ 4). This graph shows nearly zero absorbance above 400 nm, indicating that none of these polysulfide species (S82−, S62−, S42−) were detectable in the solution at levels measurable by UV spectroscopy. In contrast, the green-colored graph represents the lowest concentration sample (1.25 mM sulfur in P111i4FSI without any Li salt), which shows significant UV absorption, confirming the presence of detectable amounts of sulfur. This comparison suggests that the concentration of sulfur in the 30 mol% LiTFSI in P111i4FSI solution is far lower than 1.25 mM. It is likely below the quantification or detection limits of the UV and Raman spectroscopy methods we used. This finding indicates that the solubility of polysulfides in the 30 mol% LiTFSI electrolyte is extremely low-essentially negligible or close to zero.
Overall, both the Raman and UV spectroscopy results verify the above conclusion that high concentrations of LiTFSI in P111i4FSI greatly suppress the dissolution of lithium polysulfides. Having satisfied ourselves that these mixed anion, high salt concentration LiTFSI in P111i4FSI electrolytes can both support Li metal cycling and suppress the presence of polysulfide, we further investigated these electrolytes in full cell configurations as discussed below.
Fig. 4a shows how the discharge capacity changes over cycles for Li–S cells using different electrolyte compositions: 10 mol% LiFSI (black), 30 mol% LiFSI (red), 10 mol% LiTFSI (blue), 30 mol% LiTFSI (green), 50 mol% LiFSI (purple), and 50 mol% LiTFSI (yellow) in the P111i4FSI ionic liquid electrolyte at 50 °C. Fig. 4b presents the charge capacity over the cycles, while Fig. 4c illustrates the coulombic efficiency (how well charge is retained) for each cycle. Fig. 4d shows the discharge capacity for the 10th and 20th cycles of these cells, and Fig. 4e highlights the coulombic efficiency for the 1st and 20th cycles. Long-term cycling performance, shown in Fig. 4f, compares discharge–charge capacity over time for cells with 30 mol% LiTFSI, 50 mol% LiFSI, and 50 mol% LiTFSI in P111i4FSI electrolytes. Here the high concentration LiFSI in P111i4FSI is used as a comparison since this electrolyte composition has previously been shown to support high-rate, high-capacity cycling.28 The charge–discharge cycling was conducted within a voltage range of 1.5–2.8 V vs. Li/Li+. During discharge, the cells displayed a single voltage plateau at 2.0 V, which likely corresponds to a quasi-solid–solid transition from S8 to Li2Sn (Fig. 4g). During the charge cycle, a single plateau at 2.4 V indicates the oxidation of lower-order polysulfides back to elemental sulfur (S8).13 Fig. 4h shows the capacity retention after the 20th and 200th cycles for cells with 30 mol% LiTFSI, 50 mol% LiFSI, and 50 mol% LiTFSI in P111i4FSI electrolytes.
Fig. S2† demonstrates that the P111i4FSI-based electrolyte with LiFSI salt also significantly enhances the utilization of the cathode active material, resulting in considerably higher charge and discharge capacities compared to traditional solvent-based electrolytes. As illustrated in Fig. S3,† increasing the LiFSI salt concentration to 30 mol% within the P111i4FSI electrolyte improves cathode utilization more effectively than a lower concentration of 10 mol%. Similarly, Fig. S4† shows that the P111i4FSI-based electrolyte with LiTFSI salt markedly boosts cathode active material utilization, leading to higher charge and discharge capacities compared to conventional electrolytes. As depicted in Fig. S5,† increasing the LiTFSI salt content to 30 mol% in the P111i4FSI electrolyte enhances cathode utilization more than a 10 mol% concentration of the same salt. This is consistent with the hypothesis that highly concentrated IL electrolytes can reduce polysulfide dissolution in the electrolyte.
A wholistic view of the effect of electrolyte composition on Li–S cell performance is presented in Fig. 4. The data shows that a concentration of LiTFSI salt (30 mol%) in the mixed FSI/TFSI anion electrolyte and 50 mol% of LiFSI in single anion leads to better cathode utilization. In contrast, lower salt concentrations (10 mol% in both LiTFSI and LiFSI) and 30 mol% LiFSI result in a lower capacity. The 30 mol% mixed-anion electrolyte cell delivered a discharge capacity of ∼1050 mA h g−1 in the first cycle and ∼740 mA h g−1 in the 20th cycle, while the 50 mol% single-FSI-anion electrolyte cell showed ∼1030 mA h g−1 initially and ∼730 mA h g−1 after 20 cycles. This demonstrates that similar electrochemical performance can be attained at a lower total salt concentration through the use of a mixed-anion approach. These results highlight how mixed-anion ionic liquid electrolytes can minimize polysulfide dissolution, promote stable and conductive electrode interfaces, and support sustained battery performance with reduced salt content. This indicates that not only a higher lithium salt concentration but also using the mixed TFSI/FSI anions at lower overall Li salt concentration in these IL electrolytes further reduces polysulfide dissolution and appears to support the formation of more stable and conductive electrode interfaces (as will be discussed later). This leads to lower cell polarization, better utilization of the cathode, and more sustained capacity over time.
Fig. 5a shows the typical discharge curve of Li–S cells in standard electrolytes, with two plateaus at 2.3 V and 2.1 V. The first plateau, at 2.3 V, is due to the formation of higher-order polysulfide intermediates (Li2Sn, n ≥ 4), which dissolve in the electrolyte.3 The second plateau, at 2.1 V, corresponds to the formation of lower-order sulfides (Li2Sn, 1 ≤ n ≤ 2), which are insoluble.19 A schematic of possible polysulfide intermediates formed during the lithiation and discharge process is shown in Fig. 5b. This schematic helps compare voltage profiles in different electrolyte systems. Fig. 5c shows the discharge curve of Li–S cells using the new electrolyte studied in this work (30 mol% LiTFSI in P111i4FSI), which has a single voltage plateau at 2.0 V. This system follows the quasi-solid-state sulfur reduction, without the formation of higher-order electrolyte soluble polysulfides. The initial discharge capacity of the sulfur electrode is 1050 mA h g−1, around 63% of the sulfur's theoretical capacity.
Fig. 5d shows the discharge–charge curve of Li–S cells using this new electrolyte. The first charge capacity of the sulfur electrode is 890 mA h g−1, which is 60 mA h g−1 less than the first discharge capacity. This loss in capacity could be due to irreversible electrolyte decomposition that passivates the sulfur electrode. There is a noticeable difference between the voltage profiles of the first and second discharge cycles. In the first cycle, the discharge voltage plateau at 50% depth of discharge (DOD) is 1.89 V, but this increases to 1.97 V in the second cycle. Similarly, the charge voltage plateau at 50% depth of charge (DOC) decreases from 2.42 V in the first cycle to 2.37 V in the second cycle. These changes in the discharge and charge voltage plateaus may be due to the sulfur electrode–electrolyte interface becoming less resistant.
During the first lithiation, the voltage drops from the open-circuit voltage (OCV) to 1.85 V and then gradually rises to 1.88 V. To separate the effects of electrode passivation from sulfur lithiation, a control cell was made using an electrode of the same composition (HSAC, C65, and CMC) but without sulfur. This control cell, with a lithium metal anode, was cycled at a constant current within the same potential window (1.5 to 2.8 V) using the same 30 mol% LiTFSI in P111i4FSI electrolyte. Fig. 5e shows the discharge–charge curve of the Li-HSAC cells. The initial charge capacity is 50 mA h g−1, mostly capacitive, which is 180 mA h g−1 less than its first discharge capacity. This loss in capacity is likely due to irreversible electrolyte breakdown that passivates the HSAC electrode. During the first lithiation, a voltage plateau at 1.76 V indicates electrolyte breakdown. The voltage drops from OCV to 1.73 V and then gradually rises to 1.76 V, possibly due to a decrease in resistance at the electrode–electrolyte interface. Finally, Fig. 5f shows the discharge–charge capacity of Li-HSAC cells over several cycles. After the first few cycles, passivation ends, and only capacitive storage remains.
The electrochemical impedance spectroscopy (EIS) analysis of Li–S cells is modelled using an equivalent circuit Fig. 7(a), which includes different resistances and constant phase element (CPE) representing key cell processes. Rel represents the bulk resistance of the ionic liquid electrolyte, RSEI,Li corresponds to the resistance of the solid-electrolyte interphase (SEI) at the lithium anode, and RSEI,S represents SEI resistance at the sulfur cathode. The charge transfer resistance at the sulfur electrode (RCT,S) reflects the kinetics of the S redox reactions. A constant phase element for diffusion (CPEdiff) models lithium-ion diffusion limitations at low frequencies, highlighting mass transport effects in the electrolyte and electrode structure.
Observing the electrochemical behaviour with 30 mol% LiTFSI in P111i4FSI at open circuit voltage (OCV) Fig. 7(b), the charge transfer resistance at the sulfur electrode (RCT,S), prior to lithiation, is 585 Ω, indicating an impediment to the electrochemical reactions which would lead to slow reaction kinetics. The charge transfer resistance at the sulfur electrode (RCT,S) decreases from 585 Ω at OCV to 96 Ω post-lithiation as shown in Fig. 7(c). After lithiation, RSEI,S is 7.7 Ω, the SEI resistance at the lithium electrode (RSEI,Li) decreases from 149.7 Ω at OCV to 20.25 Ω after lithium stripping, signifying the modification of a passivation layer that stabilizes the lithium electrode.
Fig. 7(d) and (e) display the EIS spectra for 50 mol% LiFSI in P111i4FSI electrolyte. A similar trend is observed but with lower resistance values. At OCV, RCT,S is 349 Ω which decreases to 46.3 Ω after the first lithiation step. After lithiation, RSEI,S is 7.4 Ω the SEI resistance at the lithium electrode (RSEI,Li) decreases from 279 Ω at OCV to 38.4 Ω after lithium stripping, signifying the modification of a passivation layer that stabilizes the lithium electrode. Both electrolytes show a significant reduction in charge transfer resistance (RCT,S) after lithiation as a result of favorable SEI formation at the sulfur electrode.
The difference in voltage between the charge and discharge cycles of the cell, referred to as polarization, is a direct indicator of the stability of the sulfur electrode–electrolyte interface. In general, higher polarization reflects greater cell resistance. A high-resistance interface leads to a more polarized cell, resulting in a large overpotential. This is undesirable, as it decreases the efficiency of the battery. In Fig. 8a, the voltage vs. real capacity graph shows the cell performance during cycling. The capacity fade observed could be due to several factors related to the sulfur electrodes. To better understand how polarization evolves during cycling, we normalized the charge and discharge capacity and display the voltage vs. depth of discharge–charge capacity, which is shown in Fig. 8b. This normalized graph provides a clearer picture of the cell polarization behavior. The key observation from this analysis is that there was no significant change in polarization during cycling. This indicates that the SEI, formed by the mixed anion ionic liquid electrolyte at the sulfur electrode–electrolyte interface, remained stable throughout the cycling test. Additionally, the combination of a stable SEI and the electrolyte's ability to suppress polysulfide dissolution contributes to the excellent utilization of sulfur and higher discharge capacity. Since the electrolyte has ultralow or negligible solvating power for polysulfides, the cell can maintain its performance over multiple cycles, with minimal loss in capacity.
In Fig. 9c, the separator from the cycled cell was washed in neat-IL (P111i4FSI) to transfer any polysulfides that may have been trapped in the separator into the ionic liquid solution. This solution was then placed in a sample container for UV spectroscopy analysis to determine the presence of polysulfides. Fig. 9d shows the UV spectra obtained from this analysis. As shown in Fig. 9e, the absorption spectrum displayed nearly zero absorbance above 400 nm. This suggests that none of the higher-order polysulfide species (such as S82−, S62−, and S42−) were present in detectable amounts in the solution. The results from this analysis indicate that the polysulfide concentration in the electrolyte was extremely low and essentially undetectable within the resolution of the UV spectroscopy. This finding suggests that the cycling process in the tested Li–S cell led to minimal polysulfide dissolution. By limiting polysulfide dissolution, the electrolyte can prevent the shuttle effect, which improves battery efficiency and prolongs its lifespan. The clean appearance of the lithium anode and the lack of significant polysulfide absorption in the UV spectra further confirm that the system effectively suppressed the unwanted side reactions typically seen in Li–S batteries, at least under the conditions utilized here (50 °C for 15 cycles).
Footnote |
† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d5eb00009b |
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