Electrochemical Energy Reviews

, Volume 1, Issue 3, pp 388–402 | Cite as

Toward Better Lithium–Sulfur Batteries: Functional Non-aqueous Liquid Electrolytes

  • Shizhao Xiong
  • Michael Regula
  • Donghai WangEmail author
  • Jiangxuan SongEmail author
Review article


Having been extensively studied in the past five decades, lithium–sulfur (Li–S) batteries possess a high theoretical energy density (~ 2600 Wh kg−1), offering the potential to power advanced twenty-first-century technologies such as electric vehicles and drones. However, a surprisingly complex engineering challenge remains in the application of these batteries: the identification of appropriate electrolytes that are compatible with both sulfur cathodes and lithium metal anodes. Non-aqueous liquid electrolytes, typically consisting of a lithium salt dissolved in an organic solvent, cannot themselves demonstrate effective electrochemical performances. Researchers have found that functional electrolytes offer unique possibilities to engineer the surface chemistries of sulfur cathodes and lithium anodes to enable long-term cycling. In this article, recent progresses in the development of functional non-aqueous liquid electrolytes in Li–S batteries are reviewed, including novel co-solvent solutions, lithium salts, additives, redox mediators, and ionic liquids. Characterization techniques and interpretations are cited to elucidate the effects of these components on the kinetics of sulfur redox reactions, lithium passivation, and cell performance. The information presented and the studies highlighted in this review will provide guidance for future optimized electrolyte designs.

Graphical Abstract


Lithium–sulfur batteries Additive Electrolyte Lithium metal anode Redox mediator 

1 Introduction

Lithium-ion batteries (LIBs) have propelled many different sectors of the global economy in the twenty-first century, particularly portable electronic devices. However, fundamental shifts in key lithium-based battery components are required to continue this electrification revolution. This is because the layered metal oxide cathodes and graphite anodes of current state-of-the-art batteries, dependent on intercalation chemistry, are insufficient to meet growing energy demands, particularly for electric vehicles and grid-level storage applications, due to their low specific capacities and energy densities [1, 2, 3].

One proposed fundamental shift in battery chemistry is toward lithium–sulfur (Li–S) batteries that utilize lithium metal anodes and sulfur cathodes. The theoretical gravimetric energy density of Li–S batteries is nearly six times higher than that of intercalation-dependent LIBs, being attributed to both sulfur and lithium possessing significantly higher specific capacities compared with their counterparts [4, 5, 6]. Additionally, sulfur is an abundant, cost-effective, and environmentally benign cathode material and is ideal for the development of sustainable energy systems.

However, despite more than five decades of investigation, the commercialization of Li–S batteries remains elusive because of the chemical properties of both sulfur and metallic lithium.

Focusing first on the sulfur cathode, several high-level challenges are present that require further research. The first challenge is that sulfur (S8) is an electronic insulator. This results in the need for the addition of materials not as electrochemically active as sulfur to produce electronically conductive environments, preventing the achievement of the theoretical energy density of Li–S batteries [7, 8, 9, 10].

The second challenge is that the dissolution of polysulfides, intermediates of sulfur reacting with lithium, can result in a parasitic “shuttle effect” [11, 12]. Particularly notable in organic electrolytes, investigations have revealed major consequences of polysulfide dissolution on the charge–discharge characteristics of Li–S systems. The sulfur discharge process is often represented by a two-step mechanism (Fig. 1), in which sulfur reacts with lithium to form long-chain polysulfides (Li2Sx, 4 ⩽ x ⩽ 8). These long-chain polysulfides then further react to form short-chain polysulfides/sulfides (x  ⩽ 3) [10, 11, 12]. All of these discharge products have the ability to disproportionate, by which sulfide anions can recombine and dissociate with each other [13, 14]. The shuttle effect is particularly notable between highly soluble long-chain polysulfides and the electrode surfaces, which results in low sulfur utilization during discharge, low columbic efficiency during charging, high self-discharge rates, and lithium metal anode corrosion [12].
Fig. 1

Schematic diagram of the redox reaction in a Li–S cell showing the progression of elemental sulfur (S8) to lithium sulfide (Li2S) in an ether electrolyte. Because sulfur anions (particularly long-chain polysulfides) can disproportionate, these species can react with both the lithium metal anode and the sulfur cathode. This is known as the “shuttle effect”

As for sulfur reduction and oxidation, although the two-step mechanism is helpful in the understanding of the fundamentals of Li–S batteries, it is an oversimplification of the true behavior of the chemical and electrochemical reactions occurring in the system. In practice, researchers have observed two different sulfur redox reaction pathways: a direct reduction of sulfur to S42− and S22− anions and a more complex pathway in which sulfur progresses through both anionic and radical intermediates (Fig. 2) [13, 21]. The properties of the polysulfide solvent, specifically the Lewis donor number, appear to govern which redox process occurs. In addition, the type and concentration of the supporting lithium salt also influences the solubility, the equilibrium, the mass transfer, and the kinetics of the redox reaction [8].
Fig. 2

Two proposed sulfur redox reaction pathways. Here, the properties of the solvent appear to govern which mechanism is preferred. In DOL: DME, sulfur redox proceeds solely through anionic intermediates, whereas in DMSO, sulfur redox proceeds through a combination of anionic and radical intermediates. Reproduced with permission [20]. Copyright 2016, American Chemical Society

As for the final challenge, the end products of the Li–S reaction chain, lithium sulfides (Li2S and Li2S2), do not dissolve in organic electrolytes and are also electronic insulators that are detrimental to the passivation layer of electrode surfaces [13, 14]. While disproportionation can help regain some of these species, lithium sulfides remain electrochemically inactive on both the sulfur cathode and the lithium anode surfaces. As this passivation film builds up, the polarization of the cell leads to further capacity fading as active materials face increased interfacial resistance at the electrodes [22].

In addition to the resolution of conquering high-level challenges associated with sulfur and polysulfides, cathode optimization engineering decisions must also take into consideration the effects of electrolyte solutions on lithium metal anodes. Lithium possesses the lowest electrochemical potential of any material and readily reacts with many liquid electrolytes to form a passivation film known as the solid electrolyte interface (SEI) [23, 24]. Left uncontrolled, lithium can agglomerate into long, branched structures called “dendrites” during cycling, which will grow and penetrate separators. This can result in short-circuited batteries, leading to unsafe operating conditions.

Researchers have proven that SEIs formed on both sulfur cathodes and lithium anodes using conventional, organic electrolytes are insufficient for long-term cycling. Recent developments with functional electrolytes offer novel methods to tune surface chemistries and promote reversible reaction sites for active materials. This review will focus on functional non-aqueous liquid electrolytes for Li–S batteries and will aim to provide a comprehensive examination of the unique effects imparted by novel co-solvent solutions, lithium salts, additives, redox mediators, ionic liquids, and other components in an already complex reaction environment. Novel characterization techniques and interpretations are also highlighted to guide future electrolyte design.

2 Fluorinated Co-solvents

The redox reactions of polysulfide species have been widely studied in different solvents, especially in ether-based solvents such as tetrahydrofuran (THF), dimethyl sulfoxide (DMSO), dimethylformamide (DMF), and 1,3-dioxolane (DOL) [18, 25, 26, 27, 28, 29]. In particular, DOL possesses high ionic conductivity, low freezing point (− 95 °C), low viscosity (0.6 mPa s−1), and the ability to form a protective SEI layer on lithium metals [27, 30, 31, 32]. One of the initial motivating factors for introducing DOL into lithium batteries was the observation that DOL can allow for the formation of a passivation layer on anode surfaces [27] in which the polymerization of DOL can be triggered either by trace HTFSI or radical-catalyzed isomerization [33], forming ROLi, HCO2Li, and poly-DOL oligomers with -OLi edge groups (Fig. 3). Although these groups can offer increased SEI flexibility and work to suppress dendrite formation, they do not limit the shuttle effect, contain the self-discharge rate, or promote the high coulombic efficiency necessary for long-term cycling [31, 32, 34, 35].
Fig. 3

The electrochemical and chemical mechanisms for the decomposition of DOL on lithium metal

To properly engineer new electrolyte solvents for Li–S batteries, the effects of electrolytes on the equilibrium between different polysulfide species, sulfur redox reaction pathways, and chemical states of decomposition products must be understood. These factors are heavily dependent on the physiochemical properties of the solvent, including viscosity, solubility of polysulfides, and donor numbers [27, 28, 29, 30, 36]. A practical route toward optimizing many of these properties involves the addition of co-solvents with functional groups that can enhance SEIs. Fluorinated ethers have been extensively studied as co-solvents to provide positive contributions to cell performance.

Fluorinated solvents (Fig. 4) have a lower solubility to polysulfides than conventional ethers. Zhang et al. [37] reported that a 1,1,2,2-tetrafluoroethyl-2,2,3,3-tetrafluoropropylether (TTE) additive demonstrated anti-solvent effects that limited self-discharge rates, while Zu et al. [38] reported that the addition of this fluorinated ether (1:1 v with DOL) improved capacity retention, coulombic efficiency, and Li–S cell shelf life. In the study by Zu et al. [38], the researchers attributed these improved performances to the reaction between TTE and lithium, producing lithium fluoride (LiF) and compounds with conjugated C=C bonds. Together, these components form a LiF-rich protective surface film on both the sulfur cathode and the lithium metal, providing a physical barrier to limit polysulfide dissolution. Additionally, Azimi et al. [39] reported that TTE can better reduce the solubility of long-chain polysulfides as compared with DOL/1,2-dimethoxyethane (DME) (Fig. 5), leading to better retention of active materials in the cathode structure and keeping lithium sulfides electrochemically active upon charging. Azimi et al. [39] also reported that the SEI on the cathode can slow the diffusion of polysulfide intermediates, further improving the utilization of active materials. The TTE electrolyte was paired with lithium nitrate (LiNO3) and demonstrated a 0.7% and 0% self-discharge rate after resting for 10 h and subsequently cycling at room temperature and 55 °C, respectively [39].
Fig. 4

A selection of fluorinated ether co-solvents that have been tested for Li–S battery electrolytes

Fig. 5

Schematics of the polysulfide dissolution during the charge–discharge process in a Li–S battery with a DOL/DME and b DOL/TTE electrolytes, illustrating how polysulfides are confined to the cathode in TTE-containing electrolytes. c The transition from a dark red to light yellow polysulfide solution demonstrates the effect of increasing the concentration of TTE on polysulfide dissolution. Reproduced with permission [39]. Copyright 2015, American Chemical Society

Anti-solvent effects have also been observed in studies using bis(2,2,2-trifluoroethyl) ether (BTFE) and hexafluoroisopropyl methyl ether (HFME) [40, 41]. For example, Wang et al. [40, 41] reported that the use of BTFE decreased self-discharge rates to 4% and 25% in low-loading (< 1 mg cm−2) and high-loading (~ 5 mg cm−2) sulfur cathodes, respectively, and attributed this to the formation of a LiF-rich interface. The researchers suggested that with low viscosity and good wettability, BTFE can promote fast electrochemical reaction kinetics of sulfur species, particularly in the second plateau of the discharge process, improving sulfur utilization and enabling high-sulfur loading cathodes to be paired with both pre-lithiated graphite [40] and lithium metal anodes (Fig. 6) [41]. Zhang et al. [36] employed HFME (60 wt%) as an anti-solvent agent in a “strengthened pseudo-concentrated electrolyte” and was able to obtain stable long-term cycling performances.
Fig. 6

The charge–discharge profiles of Li–S batteries cycled with a DOL/DME and b DOL/BTFE electrolytes. c The cycling performance at a current rate of C/10. d (A) The dissolution of 0.25 M Li2S8 in DOL/DME and (B) The insolubility of the same solution in DOL/BTFE. Reproduced with permission [40]. Copyright 2017, American Chemical Society

1,2-(1,1,2,2-tetrafluoroethoxy)ethane (TFEE) has also be examined as a co-solvent. Dominko et al. [42], using a combination of the operando sulfur K-edge XANES (X-ray absorption near edge structures) analysis, the UV/Vis spectroscopy, and the COSMO-RS computation, revealed that TFEE can allow Li+ ions to be two times more likely to coordinate with polysulfides and LiTFSI as compared with reference ether samples, resulting in lower polysulfide solubility in electrolyte solutions. In addition, the researchers reported that this coordination ensures that Li+ ions do not interact with oxygen groups in both TEGDME and DOL, limiting the rate of electrolyte decomposition. Despite these promising enhancements, however, the conductivity of the electrolyte solution was found to be one order of magnitude lower than DOL/DME even at the lowest concentration of TFEE tested (8/2 v/v DOL/TFEE).

Finally, Wen et al. [43] reported a sulfur reduction mechanism with three plateaus rather than the normally observed two in Li–S batteries in their study of 1,3-(1,1,2,2-tetrafluoroethoxy)propane (FDE, 80 v%). The proposed three steps include the following: (1) Li2S8/Li2S6 → Li2S4, (2) Li2S4 → Li2S3/Li2S2/Li2S, and (3) Li2S2 → Li2S, in which the third step is the notable difference between FDE and conventional ether electrolytes. Gu et al. [43] attributed this plateau in the voltage profile to the novel surface chemistry of the passivated lithium metal and noted that the length of the second and third step plateaus was highly dependent on the current rate, a general characteristic of polysulfide reduction.

In contrast to ether electrolytes, carbonate electrolytes are generally incompatible with polysulfides. This is because polysulfides react irreversibly with carbonates to form numerous undesired by-products instead of progressing through the standard polysulfide reaction chain [44]. However, studies have reported that the decomposition products of fluoroethylene carbonate (FEC) can exhibit properties beneficial to Li–S batteries with notable beneficial species being formed including LiF, LixPFy, and LixPOFy, which are analogous to species formed during LiNO3 decomposition (highlighted later in this review) [35].

3 Lithium Salts

Li–S batteries typically contain two different types of lithium salts: (1) supporting lithium salts, which act as charge carriers, and (2) lithium polysulfides, which are the intermediates of the sulfur reduction reaction. Increasingly, a third lithium salt is employed by researchers to provide additional SEI components to limit polysulfide reactivity with fresh lithium metal.

3.1 Lithium Nitrate

Among lithium salt additives, lithium nitrate (LiNO3) has been the most comprehensively investigated additive for Li–S batteries. The surface chemistry of films formed on lithium by using LiNO3 was first characterized by Aurbach et al. [35] in 2009. Passivation layers formed by using LiNO3 can limit the contact between fresh lithium and polysulfides, delivering a facile approach to suppress the redox shuttle effect and continuous corrosion of lithium anodes [35, 45, 46]. SEIs with LiNO3 primarily consist of LixNOy species (LiNO2, Li3N, and Li2N2O2) if cycled with lithium polysulfides and DOL, enabling fast SEI formation and long-term cycling of lithium anodes in polysulfide-rich environments [35, 47, 48]. Li et al. [49] demonstrated the synergistic reaction mechanisms between LiNO3 and polysulfides, which can promote dendrite-free lithium anodes that can accommodate a deposition capacity of up to 6 mAh cm−2 under a current density of 2 mA cm−2. The researchers suggested that LiNO3 and Li2S8 can work together to suppress LiTSFI decomposition, as evidenced by depressed –CF3 peaks in the XPS spectra, corresponding to SEI-degrading products of LiCF3 and Li2NSO2CF3. The top layer of the SEI was also found to consist of oxidized sulfur compounds (such as Li2SO4 and Li2S2O3), whereas the bottom layer was found to consist of reduced nitrogen compounds (Li2N2O2, LiNxOy) and lithium sulfides (Fig. 7). Furthermore, Cheng et al. [48] used a lithium metal anode with this advanced SEI to cycle both sulfur and LiNi0.5Co0.2Mn0.3O2 (NCM) cathodes and reported that the resulting Li–S cell demonstrated impressive long-term cycling performances. The resulting Li-NCM cells delivered a specific capacity of up to 150 mAh g−1.
Fig. 7

Schematic of SEIs formed in DOL/DME electrolytes with LiNO3, Li2S6, and LiNO3 + Li2S6. Reproduced with permission [67]. Copyright 2014, Elsevier

However, despite these promising performance enhancements, there are notable shortcomings of relying on LiNO3. For one, the continuous, irreversible consumption of LiNO3 during the SEI formation results in increased SEI resistance [50, 51, 52]. Secondly, LiNO3 is unable to prevent the irreversible loss of sulfur species during the SEI formation, leading to capacity fading [53]. Finally, LiNO3 comes with significant safety concerns because of the strong oxidation of NO3 groups, particularly in the presence of sulfur and carbon black [32, 54].

3.2 Lithium Borates

Lithiumbis(oxalato) borate (LiBOB) is a common LIB additive that produces less corrosive by-products and is more thermally stable compared with LiNO3. Surface films formed by using LiBOB exhibit two types of passivating films, one of which is dependent on the concentration of salts in the electrolyte and the other is dependent on the electrochemical interactions of the salts with Li–S battery components [55]. In turn, these passivating films govern the effectiveness of controlling the polysulfide shuttle effect. Xiong et al. [55] reported that a 4 wt% LIBOB solution can provide the best initial cycling performance at the lower voltage plateau and produce a capacity of 300 mAh g−1 at the upper voltage plateau, providing the best balance for the formation of an insulating layer that can drive Li+ diffusion. The researchers also reported, however, that over the course of cycling, the SEI undergoes severe structural breakdowns despite having a smoother surface morphology. Wu et al. [56] reported a fluorinated-LiBOB derivative, lithium oxalyldifluoroborate (LiODFB), provided a LiF-rich SEI layer that was effective with fluorinated ethers. Here, the researchers suggested that the LiF-rich SEI was formed from reactions between LiODFB, and both DOL and DME. The optimal LiODFB concentration was 2 wt%, in which the passivation film resistance was the lowest. Concentrations above 2 wt% would result in thick SEI layers on the lithium metal surface. In this study, the resulting cells using LiODFB achieved a coulombic efficiency of 94.6% after 100 cycles without LiNO3. However, similar to LiBOB, sizable cracks were present in the SEI and limited the electrochemical impact of the additive on Li–S cell performances.

3.3 Lithium Halides

Lithium halide salts are neither good electronic nor ionic conductors. Despite this, they have been regularly incorporated into SEI layers because their insulating behavior can actually increase Li+ diffusion rates through the SEI to ensure uniform lithium deposition. As presented above, co-solvents and salts have been specifically designed by researchers to generate LiF. The other lithium halide salts, lithium chloride (LiCl), lithium bromide (LiBr), and lithium iodide (LiI), have also been studied.

3.3.1 Lithium Fluoride

Although mostly studied as a decomposition product in other additives, LiF has received attention recently as an additive itself. Archer et al. [57] optimized an SEI using a 150nm-thick LiF layer on Cu based on density functional theory calculations and demonstrated that LiF can provide a low energy barrier for Li-ion diffusion. The researchers here also reported that in Li/Cu cells, LiF-coated Cu electrodes prepared by using radio-frequency magnetron sputtering can maintain 99% coulombic efficiency through 90 cycles and achieve a low voltage hysteresis of 70 mV. In another example, Zhang et al. [58] studied LiF protection on Cu substrates by depositing LiF through the in situ hydrolysis of lithium hexafluorophosphate (LiPF6) and reported that the uniform deposition of LiF was responsible for the formation of a columnar SEI, ensuring that lithium deposition only occurred in the radial direction. The obtained initial columbic efficiency at a deposition capacity of 1 mAh cm−2 and a current rate of 0.5 mA cm−2 was 94% by using the LiF-rich protection layer, as compared with just 85% on bare Cu. Furthermore, Cui et al. [59] coated LiF onto lithium metal using a gaseous Freon R134a (1,1,1,2-tetrafluoroethane) in which the C–F bonds in the Freon were broken by lithium metal to form LiF. This reaction also forms (CH2F–CF2) Li+, which subsequently undergoes both α and β elimination to produce additional LiF (as well as carbenes and olefins). In this study, the reactant gas was deposited onto the lithium metal at 150 °C at a pressure of 0.5 atm for 20 h, producing a 40-nm-thick LiF layer. In subsequent testing results, Li–S cells with LiF-coated Li-reduced graphene oxide were found to be capable of producing a specific capacity of 1000 mAh g−1 and a coulombic efficiency of ~ 99% at a current rate of 0.5C after 100 cycles. Huang et al. [60] also employed LiF as a component of an artificial protective layer (APL) by blending LiF with poly(vinylidene-co-hexafluoropropylene) (PVDF-HFP) at a 1:2 mass ratio to form a film. This film was found to demonstrate a Young’s modulus of 6.72 GPa, which vastly exceeded SEI films with only the polymer (0.8 GPa). Although this APL was not tested in ether-based electrolytes, it still offers a unique SEI design method that can be incorporated into Li–S batteries.

3.3.2 Lithium Chloride

LiCl has also been studied as an LIB additive by researchers. For example, Nazar et al. [61] synthesized an SEI layer rich in LiCl in their study through the direct reduction of metal chlorides onto lithium metals. Placed in a 0.167 M metal chloride solution, an SEI formed on lithium metal that consisted of both LiCl and LixM alloys (M = In, Zn, Bi, or As), whereby insulating LiCl (14.5 wt% in the SEI) prevented lithium dendrite growth while the alloys provided efficient Li+ transport. The Operando optical microscopy demonstrated dendrite-free lithium deposition in the presence of the LixM/LiCl composite protection film after cycling for 200 cycles at a capacity of 2 mAh cm−2 and a current rate of 4 mA cm−2. Paired with a lithium titanate cathode (Li4Ti5O12, LTO), the alloy/chloride-protected lithium metal anode delivered a stable capacity for 1500 cycles in a DOL/DME electrolyte without LiNO3.

3.3.3 Lithium Bromide and Lithium Iodide

LiBr and LiI [53] can form smooth and stable surfaces on lithium metal anodes and have been explored by Wu et al. [62] as LIB additives. The researchers reported that upon the oxidation at > 3 V versus Li/Li+, Br and I can form radical species that initiate chain reactions which can result in the oligomerization of DME, the main component of a protective surface film (neither LiBr nor LiI leads to the polymerization of DOL). Employed in Li–S batteries, a DOL/DME electrolyte saturated in LiBr (0.15 M) can result in only 16% active material loss after 200 cycles at a rate of C/5. With C-Li2S cathodes, the addition of 0.5 M LiI was found to decrease the overpotential of the cell from 3.3 to 2.75 V. Even without LiNO3, LiI was found to be capable of ensuring no Li2S precipitation on lithium metal surfaces and provide a high capacity of 1400 mAh g−1-sulfur at a rate of C/5.

4 Phosphorus-Containing Compounds

Phosphorus pentasulfide (P2S5) has been used extensively to synthesize solid-state lithium-ion conductors [63, 64, 65]. It has been adopted for Li–S batteries because P2S5 is electrochemically stable in the Li–S reduction window and can react with all orders of polysulfides/sulfides, notably Li2S and Li2S2, to ensure that these insulating materials do not form on electrode surfaces. A product of the P2S5/polysulfide reaction is Li3PS4, which is a superionic conductor that is a major component of lithium metal passivation films [66, 67]. Other phases, such as Li4P2P7 and Li2P2S6, can also form as a result of disproportionation reactions associated with the polysulfide formation. In one study, it was reported that the addition of a 5 wt% Li2S/P2S5 solution into a 1 M LiTFSI in TEGDME electrolyte produced a 1334 mAh g−1-sulfur discharge capacity and a coulombic efficiency of > 98% after the fourth cycle [66].

Other phosphorous-containing species, such as tris(2,2,2-trifluoroethyl) phosphite (TTFP) [68] and dimethyl methylphosphonate (DMMP) [69], can act as flame retardant additives in carbonate electrolytes through the “free radical capture mechanism”. The decomposition of these phosphorus-containing compounds produces both PO and F, and both of which can react with H and OH that are produced in thermal runaway conditions. In addition, studies [69] have also reported that at the optimized values of 10 wt% TTFP and 11 wt% DMMP, the decomposition products of these additives can form conductive interfacial films on the surfaces of either pPAN@S cathodes or lithium metal anodes despite being highly viscous.

5 Redox Mediators

Redox mediators help overcome the slow kinetics of multi-step reactions involving oxidation on electrode surfaces (within a certain potential range) and the diffusion into active intermediates. These processes are also reversible; redox mediators react with the active intermediates, diffuse to the electrodes, and regenerate their state of charge [70]. Redox mediators can assist soluble active intermediates in Li–S batteries to promote fast electron transfer. Several redox mediators have been reported to activate Li2S and accelerate electrochemical kinetics for high-performance Li–S batteries (Table 1).
Table 1

Summary of redox mediators that have been tested for Li–S batteries



Molecular formula

Working mechanism



Polysulfide carriers

Dimethyl disulfide (DMDS)

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1000 mAh g−1 at 4 mg cm−2 sulfur loading

[71, 72, 73]


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1250 mAh g−1 and around 600 mAh g−1 after 50 cycles


Amine-capped aniline trimer (ACAT)

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1268 mAh g−1 at 4.5 mg cm−2 sulfur loading


Biphenyl-4,4′-dithiol (BPD)

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BPD + S n 2−  ↔ BPD-Sn (1 ⩽ n ⩽ 4)

900 mAh g−1 and around 650 mAh g−1 after 100 cycles


Carbon disulfide


Open image in new window

Not available

[55, 77]

N-methyl-N-ethyl pyrrolidinium bromide (MEP-Br)

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Around 1200 mAh g−1 at 0.2 C and maintained 930 mAh g−1 after 100 cycles


Electron transfer agents

Indium iodide


Li2Sx (x = 1, 2) + I3 → Li2Sx (x ≥ 3) + I

By charging, I − e  →  I3

983 mAh g−1 at 0.2 C, 647 mAh g−1 after 200 cycles

[79, 80]

Bis-(pentamethyl-cyclopentadienyl) iron (FeCp2*)

Open image in new window

Li2Sx (x = 1, 2) + [FeCp2*]+ → Li2Sx (x ≥ 3) + [FeCp2*]0

By charging, [FeCp2*]0 − e → [FeCp2*]+

Reversible capacity of 500 mAh g−1 (based on the mass of Li2S) after 150 cycles


Bis(cyclopentadienyl) cobalt (CoCp2*)

Open image in new window

Li2Sx (x = 1, 2) + [CoCp2*]+ → Li2Sx (x ≥ 3) + [CoCp2*]0

By charging, [CoCp2*]0 − e → [CoCp2*]+

800 mAh g−1 after 50 cycles


Bis-(pentamethyl-cyclopentadienyl) chromium (CrCp2*) and bis-(pentamethyl-cyclopentadienyl) nickel (NiCp2*)

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Li2Sx (x ≥ 3) + [CrCp2*]0 or [NiCp2*]0 ↔ Li2Sx (x = 1, 2) + [CrCp2*]+ or [NiCp2*]+

Not available


Ethyl viologendiperchlorate (EtV)

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Li2Sx (x ≥ 3) + EtV0 ↔ Li2Sx (x = 1, 2) + EtV+

30 cycles with coulombic of around 100%


Keggin-type polyoxometalates (POMs)

[PMo12O40]3− (PMo), [SiMo12O40]4− (SiMo), [PW12O40]3− (PW), [SiW12O40]4− (PW)

Li2Sx + [XM12O40]n ↔ Li2Sy + [XM12O40]m (x > y, n > m)

1280 mAh g−1 and 622 mAh g−1 after 100 cycles



Perylenebisimide (PBI)-polysulfide (PS)

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Not available

44 Wh L−1 (volumetric energy density) at 4 mg cm−2 loading of sulfur


Benzo[ghi]peryleneimide (BPI)

Open image in new window

Li2Sx + BPI∙− → Li2S + BPI

BPI + e → BPI∙−

691 mAh g−1


5.1 Polysulfide Carriers

Polysulfide carriers are organosulfide compounds that can react with lithium sulfides or polysulfides through a mechanism similar to the disproportionation. These sulfur-rich compounds diffuse to electrodes and release sulfur-containing species through electrochemical reactions. The first proof of concept of this type of redox mediator was dimethyl disulfide (DMDS), in which the electrochemical reduction pathway proceeds through dimethyl polysulfides instead of lithium polysulfides [71]. DMDS suppresses the formation of unfavorable lithium sulfides on cathodes, resulting in high capacities at high-sulfur loading levels and high capacity retention rates (Fig. 8) [72]. Furthermore, organopolysulfides can improve the mechanical flexibility and toughness of SEI layers in which the organic group can behave as a “plasticizer” for the surface film in addition to its role as a redox mediator [73].
Fig. 8

Photographs of electrolyte solutions at corresponding discharge stages of Li–S cells. The top row of photographs illustrates the conventional electrolyte, whereas the bottom row illustrates the electrolyte with 50 vol% DMDS. Reprinted with permission [71]. Copyright 2016, Wiley

5.2 Electron Transfer Agents

Electron transfer agents can oxidize insoluble lithium sulfides into soluble polysulfides and improve active material utilization. These electron transfer agents can be continuously regenerated on electrodes. For example, Li et al. [83] used bis-(pentamethyl-cyclopentadienyl) iron (FeCp2*) as an electron–hole transfer agent to improve the utilization of lithium sulfides at a charging potential of 3.2 V. With ethyl viologendiperchlorate (EtV) [84], the electrochemical reduction reactions include the following: (1) EtV2+ to EtV0, (2) sulfur to S x 2− (4 ⩽ x ⩽ 8), and (3) S42− (Li2S4) to Li2S2. EtV can also participate in the charging process, oxidizing S42− to S82−, which, in turn, can activate L2S2 to reform sulfur. However, many electron transfer agents that are listed in Table 1 irreversibly decompose on lithium metal and corrode anode surfaces. Therefore, additional surface engineering considerations must be made to protect lithium metals not only from polysulfides, but also from redox mediator additives.

5.3 Other Redox Mediators

A supramolecular gel network, formed by secondary interactions between perylenebisimide (PBI) and lithium polysulfides, has also been reported by Frischmann et al. [85] to improve sulfur utilization. Unlike polysulfide carriers or electron transfer agents, PBI employs a large molecular network dependent on π-stacking to facilitate the charge transport of polysulfides, and PBI was chosen by the researchers in their study from 85 computationally determined structures, because the electrochemical potential of its redox couple closely matches that of S82−/S42−. Gerber et al. [86] also selected benzo[ghi]peryleneimide (BPI) using a similar computational platform in their study and found that in the operating window of Li–S batteries, BPI can undergo a one-electron reduction to BPI∙−. The radical anion subsequently reduces polysulfides to Li2S away from electrode surfaces and builds porous, three-dimensional structures of Li2S.

6 Ionic Liquids

Various salts, such as alkali fluorides or chlorides, can be turned into liquids through simply raising the temperature beyond their melting points. Other salts are liquid even below 100 °C, often containing bulkier cations and anions. These latter salts are referred to as “ionic liquids” (ILs) [87].

ILs have been investigated not only for Li–S batteries, but also for many other energy applications [88]. Advantageous properties of ILs include high thermal stability, wide liquidus range (often > 200 °C), wide electrochemical window, and low volatility [89]. In addition, many of the cations and anions that are used to make ILs possess similar characteristics to salts already employed in Li–S batteries. Therefore, ILs have been investigated both as stand-alone electrolytes and components of binary solutions in traditional electrolytes.

ILs can be designed to possess low solubility to polysulfides (though it is not necessarily one of their inherent characteristics) [90, 91]. Park et al. [90] demonstrated that fluorosulfonyl amide-type anions possess weak coordination ability with Li+ ions in polysulfides, reducing their solubility. On the other hand, IL cations were found to have no effects on solubility, but can significantly alter the electrolyte’s ionic conductivity and electrochemical performance [91]. In addition, Eshetu et al. [92] reported that although certain anions can suppress the polysulfide dissolution ([BETA], PF6, BF4 and [FSA]), coexisting side reactions were detrimental to electrochemical performances. In particular, the XRD patterns for cathodes cycled in N,N-propylmethylpyrrolidinium bis(fluorosulfonyl)amide ([P13][FSA]) revealed the existence of insulating Li2SO4 and Li2S, which were not observed by using the fluorosulfonyl amide-type [P13][TFSA].

One key drawback to IL-based systems is their high viscosity, which causes an increase in cell polarization. A trivial solution to this issue is to add Li–S compatible organic electrolyte components to the IL, such as DME [93], DOL [94], or DME/DOL mixtures [95]. Because the IL is still the dominant species in the electrolyte, the safety benefits of using ILs are retained. Table 2 summarizes the electrochemical performance of these studies.
Table 2

Summary of the electrochemical performance of Li–S batteries by using ionic liquid/organic solvent electrolyte solutions


Ionic liquid

Solvent amount

Specific capacity (mAh g−1-sulfur)






33% (volume)

~ 900

0.2 C





10% (weight)

~ 550

50 mA g−1





50% (volume)

~ 700, ~ 600

0.2 C, 0.5 C



Additionally, LiNO3 can have a notable impact on the performance of ILs. For example, Barghamadi et al. [96] reported that a LiNO3-enhanced [P13][TFSA]/TEGDME electrolyte was able to achieve a 99% coulombic efficiency after 100 cycles. In addition, Wang et al. [95] reported that by using DOL/DME electrolytes in conjunction with LiNO3 and [PP13][TFSI], self-discharge rates were limited to 4% after one day of storage. Compared with DOL/DME-based electrolytes, however, IL-containing electrolytes resulted in significantly lower discharge capacities because of the restricted mobility of polysulfides to lithium metal surfaces.

7 Other Additives

Additives that have been studied but have not been categorized up to this point include copper acetate, transition metal oxides, toluene, and water.

Zu et al. [97] first studied copper acetate as an additive for protecting lithium metal anodes in polysulfide-rich environments and found that if cycled with LiNO3, the additive can decompose to CuS and Cu2S. These sulfides, in turn, can control available lithium deposition sites and contribute to a smooth lithium metal surface morphology.

Transition metal oxides have also been extensively studied as polysulfide absorbers in cathode structures, but their impact on battery performance as electrolyte additives is still unclear [98]. Conder et al. [99] visualized liquid polysulfides by using the operando X-ray diffraction for the first time and demonstrated that fumed-SiO2 can adsorb long-chain polysulfides to limit the shuttle effect, providing a coulombic efficiency of nearly 90% without LiNO3.

Finally, two common solvents, toluene [100] and water [101], have been shown to offer different methods for improving Li–S batteries. Toluene possesses a lower viscosity than TEGDME, which improves electrolyte mass transfer characteristics. However, because of the inherently low conductivity of toluene, its electrochemical effectiveness is limited. As for water, it can react with lithium metal to form LiOH, which is a better Li-ion transporter than the common ether decomposition product Li2CO3 [35, 102]. The key drawback with water is similar to that of LiNO3; once consumed, the shuttle effect infects the cell.

8 Future Perspectives

A thorough understanding of the evolution of species in Li–S electrolytes is needed to enable a new generation of energy storage systems. Functional electrolytes (summarized in Fig. 9) can offer increased ionic conductivities, restricted polysulfide dissolution, enhanced electrochemical kinetics of active materials, refined passivation of lithium anodes, and improved safety.
Fig. 9

Summary of the functional non-aqueous electrolyte components studied for Li–S batteries

Co-solvents and supporting lithium salts can be used to enhance Li+ transport in bulk electrolytes under wide temperature ranges. In Li–S cells, Li+ transport takes place via polysulfides, which themselves interact with salt anions and solvent molecules. An SEI readily forms once all these components contact with highly reactive lithium metal. Controlling this SEI layer is the key to enable long-term cycling.

Many types of novel materials can offer unique properties to tackle this challenge. Fluorinated solvents can offer improved safety with lowered volatility and nonflammability. Lithium salts, especially LiNO3 and lithium halides, can form passivation films that prevent the shuttle effect, increase the Li-ion diffusion rate, and contribute to a uniform SEI morphology. Phosphorus-containing compounds such as pentasulfide (P2S5), tris(2,2,2-trifluoroethyl) phosphite (TTFP), and dimethyl methylphosphonate (DMMP) can provide flame retardant capabilities that improve battery safety, an issue that is becoming increasingly more important in high energy density environments such as the Li–S. Redox mediators can overcome the slow kinetics of multi-step polysulfide reactions during the charge–discharge process. These mediators—acting either as polysulfide carriers, electron transfer agents, or supramolecular networks—improve the sulfur utilization by offering an additional redox pathway to activate lithium sulfides. Ionic liquids possess high thermal stability and compatibility with lithium. The IL anion controls the polysulfide dissolution and the SEI formation, whereas its cations can govern ionic conductivities. Going forward, researchers could consider combining these functional components into a single electrolyte (for example, LiNO3 in ILs), but as the number of components in the electrolyte increases, the likelihood for undesired side reactions also increases.

In conclusion, the complex chemical environment that occurs in Li–S electrolytes remains understudied compared with the development of sulfur cathode structures. Both sulfur and lithium possess fundamental chemical properties that have challenged researchers and engineers for decades. Nevertheless, the functional components discussed in this review offer unique research directions to overcome the challenges and allow for the commercialization of Li–S batteries. Further engineering of novel electrolytes can enable the use of lithium metal anodes in sulfur-rich environments, offering dramatic increases in the capacity and the energy density that will accelerate the adoption of Li–S batteries into electric vehicles and energy grids.



SX and JS would like to thank the National Natural Science Foundation of China (No. 51602250) and Thousand Youth Talents Plan Project of China for their funding support. MR and DW would like to acknowledge the Office of Vehicle Technologies of the U.S. Department of Energy under Contract No. DE-EE0007795 for its support of this work.


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Copyright information

© Shanghai University and Periodicals Agency of Shanghai University 2018

Authors and Affiliations

  1. 1.State Key Laboratory for Mechanical Behavior of MaterialsXi’an Jiaotong UniversityXi’anChina
  2. 2.Department of Chemical EngineeringThe Pennsylvania State UniversityUniversity ParkUSA
  3. 3.Department of Mechanical and Nuclear EngineeringThe Pennsylvania State UniversityUniversity ParkUSA

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