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Li–O2 and Li–S batteries with high energy storage


REVIEW ARTICLE
PUBLISHED ONLINE: 15 DECEMBER 2012?|?DOI: 10.1038/NMAT3191

Li–O2 and Li–S batteries with high energy storage
Peter G. Bruce1*, Stefan A. Freunberger1, Laurence J. Hardwick1? and Jean-Marie Tarascon2
Li-ion batteries have transformed portable electronics and will play a key role in the electrification of transport. However, the highest energy storage possible for Li-ion batteries is insufficient for the long-term needs of society, for example, extendedrange electric vehicles. To go beyond the horizon of Li-ion batteries is a formidable challenge; there are few options. Here we consider two: Li–air (O2) and Li–S. The energy that can be stored in Li–air (based on aqueous or non-aqueous electrolytes) and Li–S cells is compared with Li-ion; the operation of the cells is discussed, as are the significant hurdles that will have to be overcome if such batteries are to succeed. Fundamental scientific advances in understanding the reactions occurring in the cells as well as new materials are key to overcoming these obstacles. The potential benefits of Li–air and Li–S justify the continued research effort that will be needed.

E

nergy storage will be more important in the future than at any time in the past. Among the myriad energy-storage technologies, lithium batteries will play an increasingly important role because of their high specific energy (energy per unit weight) and energy density (energy per unit volume). Since their introduction in 1991, Li-ion batteries (Fig.?1) have transformed portable electronic devices1–4. New generations of such batteries will electrify transport and find use in stationary electricity storage. However, even when fully developed, the highest energy storage that Li-ion batteries can deliver is too low to meet the demands of key markets, such as transport, in the long term. Reaching beyond the horizon of Li-ion batteries is a formidable challenge; it requires the exploration of new chemistry, especially electrochemistry, and new materials2,5,6. There are few options. Two, based on lithium, are receiving intense interest at the present time and will be discussed in this Review: rechargeable Li–air (hereafter referred to as Li–O2 as O2 is the fuel) and Li–S batteries7. Other options, especially Zn–air, have been reviewed in detail recently elsewhere8–14. Although Li–O2 and Li–S share the same anode, and have active cathode components (O2 and S) that are nearest neighbours in group?16 of the periodic table, there are important differences related to the different chemistry of O and S and the different states of matter of their cathodes. Li–S has been investigated since the 1940s; the problems are formidable and extensive efforts have been made to address them over the intervening 70?years. Important advances have made recently, but significant challenges remain7,15–24. In comparison, Li–O2, especially with a non-aqueous electrolyte, has received much less attention until recently 7,8,25–30. As in the case of Li–S, major challenges will have to be solved if Li–O2 batteries are to succeed. The renaissance of interest in Li–S and the upsurge of interest in Li–O2, based on aqueous and non-aqueous electrolytes, reflects the need for electrochemical energy-storage devices that can offer a leap forward; for example, delivering electric vehicles with a driving range approaching the goal of ~500?km between charging. In the limited space available, we cannot hope to review all the excellent work that has taken place on these two battery technologies. Instead, we shall begin by considering the energy that can be stored in Li–O2 and Li–S cells, and then examine each system, how it operates, and the challenges facing research that attempts to advance Li–O2 and Li–S batteries. Paramount among the present challenges is a fundamental understanding of the chemistry taking place in the cells and the discovery of new materials.
1

Energy storage from theory to practice

The theoretical specific energies (gravimetric energy densities) and energy densities (volumetric energy densities) for Li–S and Li–O2 are given in Table?1, where they are compared with those for Li-ion and Zn–air. The values are based on the cell reactions in column 1, that is, the energy obtained per unit mass or per unit volume of the active components of the anode and cathode. Often a value of 11,586?Wh?kg?1 is quoted for the Li–O2 cell; however this is based on the mass of Li alone. All metal–air cells gain mass (O2) as they discharge, so the mass of O2 should be included, as it is in Table?1. The leap forward in theoretical specific energy on migrating from Li-ion to Li–S and then Li–O2 is clear. It arises because Li2S, Li2O2 and LiOH in the cathode store more Li, and hence charge, than LiCoO2 per unit mass, and Li metal stores more charge per unit mass than a graphite (C6Li) anode. The theoretical energy density is also greater for Li–O2 and Li–S than Li-ion but the gain is not as great as for specific energy. Of course there is always a significant reduction in the energy stored in a battery on moving from theory to practice. A comparison of practical specific energies for several rechargeable batteries is presented in Fig.? 2. The values for established technologies are well attested but for Li–O2 are at best very rough estimates, because so far there are few realistic prototypes on which to base such figures. However, the values quoted are in line with those reported by others16,26,28. Several factors conspire to lower the energy storage of practical Li–S and Li–O2 batteries. The cathode in each case consists of a porous conducting matrix (usually carbon) in which the discharge products form, thus adding mass and volume to the cell. More Li metal than is required for the stoichiometric reaction has to be included, to compensate for the inefficiency of Li-metal cycling. The effect of these factors is illustrated in Box?1 for the case of nonaqueous Li–O2. Furthermore, packaging, current collectors and, in the case of Li–O2 cells, gas diffusion channels, will also reduce the practical energy storage. There is a general rule of thumb that moving from theory to practice involves a reduction by a factor of 3 in energy storage. Prototype Li–S cells have been developed by, for example, Sion Power and, based on their data, the specific energy of a Li–S cell is reduced from 2,567?Wh?kg?1 to 350?Wh?kg?1, that is, a factor of 7. However they state that values in the region of 600?Wh?kg?1 are expected in the near future (factor of 4) and similar values have been suggested by others13,16,31–33. If we apply a factor

School of Chemistry, University of St Andrews, North Haugh, St Andrews, Fife KY16 9ST, Scotland, UK, 2Laboratoire de Réactivité et Chimie des Solides — UMR CNRS 6007, 33?rue Saint-Leu, 80039 Amiens Cedex, France. ?Present address: Stephenson Institute for Renewable Energy, Department of Chemistry, University of Liverpool, Crown Street, Liverpool L69 7ZD, UK. *e-mail: pgb1@st-andrews.ac.uk
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REVIEW ARTICLE
Li-ion Discharge e– e


NATURE MATERIALS DOI: 10.1038/NMAT3191
Li–O2 (non-aq) Discharge e– Li–O2(aq) Discharge e– e O2 Li
+ –

Li–S Discharge e– e O2




+

e





+



+



+

Li+ LixC6 Organic Li1-x CoO2 electrolyte

Li metal

O2 Porous carbon + catalyst Li2O2
+

Li metal

Li

+

O2 Porous carbon + catalyst LiOH

Li metal

Li+ Organic electrolyte Li2S Porous carbon +S

Organic electrolyte

Li -conducting membrane

Aqueous electrolyte

O2-evolution electrode 2LiOH 2Li+ + 2e– + S

? C6Li + Li0.5CoO2

3C + LiCoO2

2Li+ + 2e– + O2

Li2O2

2Li+ + 2e– + ? O2 + H2O

Li2S

Figure 1 | Schematic representations of Li-ion, non-aqueous and aqueous Li–O2 and Li–S cells.

of 4–7 to the Li–O2 battery then the estimated practical specific energy is in the range of ~500–900?Wh?kg?1, that is, the range given in Fig.?2, which is at least 2–3 times greater than Li-ion and would be sufficient to deliver a driving range of more than 550?km, when scaled to the driving range of the Nissan Leaf (see caption to Fig.?2). Practical energy densities for Li–O2 are even more difficult to estimate with any reasonable accuracy in the absence of realistic prototypes and may not exceed Li-ion, that is, the main gain over Li-ion is in specific energy. The practical specific energies for Li–S and Li–O2 given in Fig.?2 represent the performances expected if the problems described in the subsequent sections of this Review can be addressed successfully. The main factor limiting the practical energy storage of Li–O2 and Li–S cells is the need for excess Li in the anode; this especially compromises volumetric energy density owing to the low density of Li metal (0.534?g?cm?3). Thus, improving the cycling efficiency of the Li electrode or replacing it with an alternative (see later) is one of the important challenges of maximizing the energy storage of Li–O2 and Li–S batteries. The price of Li–S and Li–O2 cells will be as important as performance in determining their adoption. The prices given in Fig.?2 are the targets at the pack level outlined by the US Advanced Battery Consortium. Cells must be assembled into packs then into systems that include the battery management. The US Advanced Battery Consortium states that the target long-term minimal selling price of a mass-produced (25,000 units) 40?KWh battery pack is US$150?kWh?1. Clearly, whichever technology is used, a significant reduction in cost is required to enable mass commercialization. Today the Li-ion pack-level price is close to US$500–600?kW?h?1 (ref. 34).

LiOH being oxidized on charge. Because H2O and O2 are involved, this is sometimes referred to as a Li–water battery. Note that we consider only Li–O2 cells with alkaline electrolytes here, as these have been more extensively studied than acidic electrolytes39–41. Although both cells involve O2 reduction on discharge, there are important differences, especially relating to the reactions at the cathode and the role of the electrolyte. As a result, aqueous and nonaqueous Li–O2 will first be considered separately; the challenges that exhibit some commonality with Li–S, especially the Li anode and porous carbon cathode, will then be discussed.

The Li–O2 battery

Schematic representations of Li–O2 cells based on aqueous and non-aqueous electrolytes are shown in Fig.?1. In both cases, on discharge, the Li-metal anode is oxidized, releasing Li+ into the electrolyte, and the process is reversed on charge. At the positive electrode, O2 from the atmosphere enters the porous cathode, dissolves in the electrolyte within the pores and is reduced at the electrode surface on discharge. When a suitable non-aqueous electrolyte is employed, O22? is formed, which, along with Li+ from the electrolyte, forms Li2O2 as the final discharge product. The peroxide is then decomposed on charging: 2Li+?+?O2?+?2e????Li2O2. Note that some authors have reported that discharge down to Li2O is possible, which would increase the energy stored (twice the Li per O, see Box?1), but may be difficult to reverse on charging 35–38. Aqueous electrolytes involve the formation of OH? and then LiOH at the cathode on discharge, according to the equation: 2Li+?+??O2?+?H2O?+?2e????2LiOH, with
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The non-aqueous Li–O2 cell. The rechargeable non-aqueous Li–O2 cell was first reported by Abraham and co-workers25 and interest has expanded rapidly in recent years36,37,42–59. Although important progress has been made by many authors, significant challenges remain; these are summarized in Fig.?3. The cathode, if exposed to ambient air, would be affected by CO2 and H2O, resulting in the formation of Li2CO3 and LiOH instead of Li2O2. Thus it is necessary to remove these gases by, for example, an O2 diffusion membrane covering the outer surface of the cathode that blocks CO2 and H2O. Such a membrane would have to support fast O2 diffusion so as not to limit the rate. Zhang and co-workers have examined many aspects of the non-aqueous Li–O2 cell60–63. These authors have investigated membranes to see if they are suitable for protecting the non-aqueous Li–O2 cell from ingress of CO2 and H2O when operating in ambient air. They investigated hydrophobic polymethylsiloxane and silicalite on a porous metal substrate56,57, and Melinex (DuPont), a polyester–polethyleneglycol copolymer, or high-density polyethylene films62. Using such membranes it was possible to discharge the cell continuously for one month in ambient air with 20% relative humidity, compared with only a few hours in the same environment in the absence of the membrane, demonstrating progress in using the cells in air. The electrolyte is a key component and one of the main challenges at present. It must be stable both to O2 and its reduced species, as well as the LiOx compounds that form on discharge; it must exhibit sufficient Li+ conductivity, O2 solubility and diffusion to ensure satisfactory rate capability, as well as wet the electrode surface and possess low volatility to avoid evaporation at the cathode. Electrolytes based on organic carbonates (for example, LiPF6 in propylene carbonate) have so far been widely used in non-aqueous Li–O2 cells, because of their low volatility, compatibility with Limetal and high oxidation stability (>4.5? V compared with Li+ or Li). However, studies of O2 as a contaminant in Li-ion battery electrolytes showed that organic carbonates are susceptible to nucleophilic attack by reduced O2 species64. Recent studies of Li–O2 cells
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NATURE MATERIALS DOI: 10.1038/NMAT3191
Table?1 | Data for several electrochemical reactions that form the basis of energy-storage devices.
Battery Today’s Li-ion ?C6Li?+?Li0.5CoO2???3C?+?LiCoO2 Zn–air Zn?+??O2???ZnO Li–S 2Li?+?S???Li2S Li–O2 (non-aqueous) 2Li?+?O2???Li2O2 Li–O2 (aqueous) 2Li?+??O2?+?H2O???2LiOH§ Cell voltage (V) 3.8 1,086 2,567 3,505 3,582 Theoretical specific energy (Wh kg?1) 387 1,086 2,567 3,505 3,582

REVIEW ARTICLE
Theoretical energy density (Wh l?1) 1,015 6,091* (ZnO) 2,199? (Li?+?Li2S) 3,436? (Li?+?Li2O2) 2,234|| (Li?+?H2O?+?LiOH)

*Based on volume of ZnO at the end of discharge; ?based on the sum of the volumes of Li at the beginning and Li2S at the end of discharge; ?based on the sum of the volumes of Li at the beginning and Li2O2 at the end of discharge; §assuming the product is anhydrous LiOH and alkaline conditions; and ||based on the sum of the volumes of Li?+?H2O consumed and the LiOH at the end of discharge.

with organic carbonate electrolytes have demonstrated electrolyte degradation does indeed occur on discharge65–69. Significantly, studies have also shown that there is little or no evidence of Li2O2 formation occurring in parallel with the electrolyte degradation69. This is important as it rules out the possibility that the remarkable ability of Li–O2 cells with organic carbonate electrolytes to sustain cycling (up to 100 cycles65) could be explained by reversible Li2O2 formation, whereas the ubiquitous capacity fading (Fig.? 3) is explained by simultaneous electrolyte degradation. Instead, such cells cycle by degradation of the electrolyte on discharge to form lithium propyl dicarbonate (C3H6(OCO2Li)2), Li2CO3, HCO2Li, CH3CO2Li, CO2 and H2O, with the decomposition products being oxidized on charge69. The implication of these results is that all studies of the Li–O2 cell with organic carbonate electrolytes are likely to be significantly affected by electrolyte degradation on discharge. This highlights the importance of presenting data from analytical methods (for example, Fourier transform infrared, Raman and mass spectrometry) to demonstrate that the product actually formed on discharge is Li2O2, when reporting results on Li–O2 cells. Recently, attention has turned to other electrolytes, especially ethers, including tetraglyme (CH3O(CH2CH2O)4CH3), dimethoxyethane (CH3OCH2CH2OCH3; DME) and polyethyleneoxide (PEO). Such ethers are certainly more stable than organic carbonates towards reduced O2 and do exhibit Li2O2 formation at ~2.7?V, at least on the first discharge42,43,64,70–72. Recently, cells using DME and LiClO4 electrolyte and all-carbon-nanofibre electrodes have shown enhanced cycling stability compared with typical cells with carbonate-based electrolyte, though fading is still present 59. However, investigation of linear (diglyme, triglyme and tetraglyme) and cyclic (1,3-dioxolane, 2-methyl tetrahydro furan) ethers in Li–O2 cells demonstrated electrolyte degradation to form Li2CO3 and HCO2Li, and CH3CO2Li, which becomes more severe on cycling 70,71. Studies of DME in Li–O2 cells by mass spectrometry showed that only ~60% of the O2 consumed on discharge is released on charge60. If these results are confirmed by further studies then the implication is that ethers, although more stable than carbonate electrolytes, are not the final solution to the challenge of identifying a suitable non-aqueous electrolyte for Li–O2 batteries. Provided Li2O2 can be formed on discharge then it has been shown that it can be oxidized on charge49,73. Such charging has been investigated by constructing electrodes in the discharged state, that is, by incorporating Li2O2 in the electrode49,73. The dominant gas evolved on charging is O2, consistent with Li2O2 oxidation without electrolyte decomposition71,73. The gap between the charge and discharge voltage has been linked to the nature of the catalyst in the positive electrode, with values ranging from 0.6? to 1.5? V37,46,49–51,74–76. However, many of these studies were carried out in electrolytes containing organic
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carbonates, which we now know decompose on discharge, as discussed above65,66,69. Therefore it is difficult to know the extent to which the catalyst is catalysing the electrolyte decomposition on discharge and the oxidation of the decomposition products on charge, rather than Li2O2 formation and decomposition. The effect of the catalyst on charging Li2O2 has been investigated by charging electrodes constructed with Li2O2 incorporated in the as-prepared cathode (Fig.?4.)37,77. Even when organic carbonate electrolytes are used, it has been shown by differential electrochemical mass spectrometry that only oxidation of Li2O2 occurs73. These studies were further facilitated by the development of a fast screening technique based on the catalytic degradation of H2O2 (H2O2(aq)?→?H2O?+??O2↑), which exhibits similar trends to the electrochemical oxidation of Li2O2 (Li2O2(s)?→?2Li?+?O2↑). Investigation included a variety of transition metal oxides, with nanowires of α-MnO2 giving the lowest charging voltage in Li–O2 cells and the most facile H2O2 decomposition of the materials investigated in the study 77. Overall, it has been demonstrated that catalysts lower the charging voltage of Li2O237,77. Studies of O2 reduction in non-aqueous electrolytes have been carried out for several decades in solvents more stable than organic carbonates64,78. However, it is important to understand the

1,000 Speci?c energy (Wh kg–1) 800 600 400 200 0 50 km 80 km 100 km Pb–acid Ni–Cd Ni–MH Li–ion Future Zn–air Li–S Li–ion 600 900 600 < 150 Today > 200 km

> 550 km

> 400 km > 225 km

160 km

Li–air < 150 R&D

Price (US$ kW h–1) 200

< 150 < 150

Available

Under development

Figure 2 | Practical specific energies for some rechargeable batteries, along with estimated driving distances and pack prices. For future technologies, a range of anticipated specific energies are given as shown by the lighter shaded region on the bars in the chart for rechargeable batteries under development and R&D. The values for driving ranges are based on the minimum specific energy for each technology and scaled on the specific energy of the Li-ion cells (140?Wh?kg?1) and driving range (160?km) of the Nissan Leaf136. The prices for technologies under development represent targets set by the US Advanced Battery Consortium137.
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REVIEW ARTICLE
Box 1 | Factors influencing the energy storage of a practical Li–O2 battery.

NATURE MATERIALS DOI: 10.1038/NMAT3191

Excess Li

Beginning with the simplest model for a non-aqueous Li–O2 cell, assuming the cathode is composed only of Li2O2, (that is, no porosity, or carbon, or binder) and assuming a stoichiometric quantity of Li in the anode, see Fig.? B1a; the specific energy and energy density are 3,505?Wh?kg?1 and 3,436?Wh?l?1 respectively, that is, the values given in Table?1. The effect of introducing excess Li metal in the anode to compensate for the Li loss on cycling is shown. For a threefold excess (n?=?4) the specific energy and energy density become 1,800?Wh?kg?1 and 1,290?Wh?l?1. Values for discharge to Li2O are also shown.
a
Energy density (Wh l–1)

Porous cathode

Taking account of porosity in the cathode, let us assume 20% of the cathode volume is occupied by C, 20% by electrolyte and 60% by Li2O2 at the end of discharge. The resulting specific energies and energy densities are shown in Fig.? B1b, where the porous cathode is coupled with a Li anode; a stoichiometric amount of Li is assumed to provide a benchmark for the highest possible energy storage that is, in principle, possible. Comparison with a Li-ion cell is shown. Bars framed green represent values for Li2O2, those framed orange are for discharge to Li2O. The electrolyte thickness is 10 μm.

Li2O2 Li consumed

Li + ?O2

? Li2O2

3,500 3,000 2,500 2,000 1,500 1,000

1

or 5,000 Energy density (Wh l–1) 4,000 3,000 2,000 1,000 1

2 3 Equivalents of Li

4

3,500 3,000 2,500 2,000 1,500 1,000

Speci?c energy (Wh kg–1) Speci?c energy (Wh kg–1)

5,000 4,000 3,000 2,000 2 3 Equivalents of Li 4 1,000

Li2O

Li + ?O2

? Li2O

b
4,000 Maximum Li–O2 Li2O Volumetric energy density (Wh l–1) Gravimetric speci?c energy (Wh kg–1) 3,000 Li2O2 2,000 3,000 4,000 Maximum Li–O2 Li2O Li2O2 2,000 LiCoO2/C

1,000 LiCoO2/C 0

1,000

0

Figure B1 | Specific energy and energy densities of the non-aqueous Li–O2 battery. a, Excess Li in the anode. b, A porous cathode. The values for today’s Li-ion battery (LiCoO2/C) shown as comparison.

fundamental mechanism of O2 reduction in the presence of Li+ and the formation of Li2O2, as well as the mechanism of Li2O2 oxidation on charging, if progress is to be made on Li–O2 cells. Detailed electrochemical investigations have probed the influence of salt and solvent on O2 reduction36,42,72. In?situ spectroelectrochemical measurements involving Raman and mass spectrometry have also been applied, which offer the important benefit of directly identifying the species involved in the reactions (intermediates and products) on discharge and charge73. (Note that in ref.?73 the peak maximum for O2 reduction in the cyclic voltammogram is 2.2?V, however the peak
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onset (the start of the reduction) is at 2.7?V, corresponding to the voltage of the discharge plateau.) The consensus from the various studies is that the mechanism of O2 reduction on discharge is: O2?+?e??→?O2? O2??+?Li+?→?LiO2 2LiO2?→?Li2O2?+?O2 (1a) (1b) (1c)

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NATURE MATERIALS DOI: 10.1038/NMAT3191
Anode Discharge e– Problems of Li metal ? Dendrite formation ? Cycling e ciency ? Requires stable solidelectrolyte interphase ? Safety issues e
Li metal


REVIEW ARTICLE
Cathode Cathode needs a membrane to block CO2 and H2O, while allowing O2 to pass. O2 Li
+



+

Li metal

O2 H2O CO2

O2 Porous carbon + catalyst Li2O2

Organic electrolyte Capacity fading Discharge 2,000 Charge

4.5 Potential (V versus Li/Li+) Electrolyte ? Stability ? Conductivity ? Volatility ? O2 solubility, di usivity 4.0 3.5 3.0 2.5 2.0 0

Voltage gap Charge

Capacity (mAh g–1 C)

1,000

Discharge 1,000 Capacity (mAh g–1 C) Porous cathode design ? Pore size, distribution ? Catalyst — type, distribution, loading 2,000

0

0

5

10

15

20

25

30

Cycle number

Figure 3 | Challenges facing the non-aqueous Li–O2 battery. Data illustrating capacity fading and the voltage gap were collected from a cell with an organic carbonate electrolyte, Li–1?M?LiPF6 in propylene carbonate–(superP:Kynar:α-MnO2 nanowires)69. Use of more stable ether electrolytes and carbon fibre nanotube electrodes — for example, Li–LiClO4 in dimethoxyethane–carbon fibre nanotube — reduces, although does not eliminate, capacity fading59.

Several studies suggest also the direct reductions36,72: LiO2?+?Li+?+?e??→?Li2O2 Li2O2?+?Li+?+?e??→?2Li2O (1d) (1e)

Equation (1d) occurs at lower voltages than equation (1a)73 and so far there has been limited evidence for the formation of Li2O. Oxidation of Li2O2 follows: Li2O2?→?2Li+?+?2e??+?O2 (2)

In other words, the process on charging is not the reverse of discharge; the latter involves O2? as an intermediate, whereas the former does not. The different pathways followed on reduction and oxidation do not violate the principle of microscopic reversibility, but arise because the kinetics of oxidizing Li2O2 directly are faster than reversing the three steps on reduction, especially disproportionation73. The result of these different pathways is the observed separation of the charge and discharge voltages. Another factor that may contribute to the voltage gap is if singlet O2 is formed on oxidation of Li2O2, whereas reduction involves the more stable triplet state72. The singlet-to-triplet O2 transition is spin forbidden and hence kinetically hindered. If the transition kinetics are significantly slower than O2 evolution on oxidation of Li2O2, then the difference in energy between the singlet and triplet states (~0.9?V) could influence the voltage gap. Because voltage gaps smaller than 0.9?V have been observed (typically 0.7?V), at most any difference in the spin states of O2 is likely to make only a partial contribution to the gap. Other electrolytes have been explored, but in much less detail than the organic electrolytes. Investigation of hydrophobic ionic liquids demonstrated that they can maintain less than 1% H2O
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content after 100?hours of operation, in the case of 1-ethyl-3-methyl imidizolium bis(trifluoromethane sulphonyl)imide. On discharge, a capacity of 5,360? mAh? g?1 (based on carbon alone) has been reported for operation of up to 56?days52. Solid electrolytes have also been investigated, in particular, cells incorporating the Li+ conductor 18.5Li2O:6.07Al2O3:37.05GeO2:37.05P2O5 (LAGP) can sustain some 40 cycles at elevated (40–100?°C) temperatures79. Cells with a solid PEO-based polymer electrolyte can be charged at relatively low voltages, as noted above72. However, in view of what is now understood about the reactions in liquid electrolytes, it will be important in the future to investigate the nature of the discharge products and the electrode reaction mechanisms in the ionic liquid and solid
4.5 No catalyst Potential (V versus Li/Li+) NiO 4.0 Co3O4 CuO Electrolytic MnO2 α-MnO2 bulk α-MnO2 nanowire
e 2O 3 α-F

3.5

3.0

0

250

500

750

1,000

1,250

Capacity (mAh g–1 C)

Figure 4 | First galvanostatic charge, i?=?70?mA?g?1 C (that is, Li2O2 oxidation) for various catalyst-containing Li–O2 cells in this study77. Figure adapted with permission from ref.?77, ??2010 ECS.
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REVIEW ARTICLE
Anode Problems of Li metal ? Cycling e ciency ? Li/ceramic interface ? Safety issues Requires Li+–conducting membrane (such as LISICON) to protect Li. Precipitation of LiOH on LISICON membrane and membrane instability in strong acidic and basic electrolyte. Discharge e– e– – +

NATURE MATERIALS DOI: 10.1038/NMAT3191
Cathode Cathode needs membrane blocking CO2 , while allowing O2 to pass O2 Li+ Aqueous electrolyte Porous carbon + catalyst O2 LiOH O2 CO2 Optimize gas-di usion electrode for Li–O2 cell. ? O2 + H2O 2OH–

Li metal

Li+-conducting membrane

O2-evolution electrode Requires cheap, e cient O2–evolution catalyst. Electrolyte Limited LiOH solubility in H2O: precipitation of LiOH, blocking electrode pores.

Requires cheap, e cient oxygen-reduction catalyst.

Figure 5 | Challenges facing the aqueous Li–O2 battery. Image of LISICON membrane reproduced with permission from ref.?39, ??2010 ECS.

electrolytes before any conclusions can be formed concerning their potential use in practical Li–O2 cells.

The aqueous Li–O2 cell. Pioneering work by Visco and colleagues in 200726, involving the development of protected Li anodes, opened the way to the rechargeable aqueous Li–O2 battery 26,40,41. Additional challenges facing rechargeable aqueous Li–O2 are summarized in Fig.?5. Unsurprisingly, significant attention has focused on protecting the Li-metal anode from the aqueous electrolyte, using a Li+-conducting but electronically insulating membrane, such as LISICON-type glass ceramic (lithium superionic conductor, Li(1+x+y)AlxTi2?xSiyP(3?y)O12), without which the rechargeable aqueous Li–O2 cell could not function. This is discussed later in the section comparing Li–O2 and Li–S as it also has applications in non-aqueous Li–O2 and Li–S. The Li anode is far from the only problem facing aqueous Li–O2 cells. Although one of the advantages of such a cell is that there is, by definition, no need to avoid ingress of H2O from the atmosphere, CO2 remains a problem as it leads to the formation of Li2CO3. As a result, it is necessary to apply a protective membrane to the outer
5.0 4.5 4.0 3.5 Cell voltage (V) 3.0 2.5 2.0 1.5 1.0 0.5 0.0 0 20 40 60 80 100 120 140 160 Discharge Charge

Capacity (mAh g-1)

Figure 6 | Load curve of an aqueous Li–O2 cell. Cell cycled over a limited range to avoid precipitation of LiOH. Figure reproduced with permission from ref.?41, ??2010 ECS.
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surface of the cathode that permits O2 diffusion while blocking CO2, or to remove the latter from the atmosphere around the battery by alternative means. The aqueous electrolyte is not inert, it takes part in the cathode reaction, see Table? 1. As discharge proceeds, the solvent is consumed and LiOH is generated, resulting in a rapid increase in the concentration of the latter. Saturation is reached at 5.3?M?l?1, corresponding to a specific capacity of only ~170?mAh?g?1 (based on the electrolyte alone) that is, barely more than a Li-ion cell. To avoid limiting the capacity, a flow-cell design, in which the electrolyte is constantly replenished with a fresh solution, has been proposed; the same authors have also proposed recharging by collecting the LiOH, extracting the Li metal then mechanically replacing the anode80,81. Otherwise, to achieve a specific capacity that offers a meaningful advantage over Li-ion cells, LiOH must be allowed to precipitate as a solid. Such precipitation has important consequences; not only will solid LiOH block the ceramic protecting the anode, but it can also clog the porous cathode. One solution to the latter is the use of an anion-exchange membrane at the cathode/electrolyte interface; this permits OH? generated in the porous cathode to be transported out, while blocking Li+ entering the electrode, hence LiOH precipitates outside the cathode. This membrane also blocks any carbonates that might be formed by CO2 ingress, preventing the formation of Li2CO3. An interesting solution to the problem of recharging the aqueous Li–O2 cell containing solid LiOH is the incorporation of a third electrode, which is used for O2 evolution on recharging (Fig.?5). The third electrode, which remains immersed in the liquid aqueous electrolyte, was found to more effectively oxidize LiOH than if solid LiOH is allowed to precipitate in the pores of the carbon cathode. Such a cell has been demonstrated to operate for several hundreds of cycles39. As in the case of the non-aqueous Li–O2 cell, a voltage gap between charge and discharge is evident. At a current density of 0.1?mA?cm?2 this can be as low as 0.3?V (ref.?39) or ~0.9?V at a higher current density of 0.5?mA?cm?2 (Fig.?6)41. The reduction and evolution of O2 in aqueous media have been studied extensively for many years. In contrast to the reduction of O2 to Li2O2 in non-aqueous electrolytes, reduction to LiOH necessitates cleavage of the O–O bond, a reaction that requires a catalyst, as does O2 evolution. The most widely used catalyst is Pt (refs 39,41), although other materials, including manganese oxides82, metallic copper 83 and various perovskite oxides84 have been investigated. One advantage of separating out the electrodes for O2 reduction
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NATURE MATERIALS DOI: 10.1038/NMAT3191
Anode Problems of Li metal ? Dendrite formation ? Cycling e ciency ? Requires stable solidelectrolyte interphase ? Safety issues ? Formation of insulating Li2S layer on Li anode Discharge e–
Li metal

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Cathode S8 Potential (V versus Li/Li+) 3.0 Address polysulphide solubility Charge process

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+

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Li+ Organic electrolyte Porous carbon +S

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Figure 7 | Challenges facing the Li–S battery. Load curve and schematic showing PEG?200-coated CMK-3–S composites that impede diffusion of the polysulphides into the electrolyte; reproduced from ref.?17, ??2009 NPG. Capacity fading for Li–S cell using graphene-nanosheets cathode; adapted with permission from ref.?21, ??2011 Elsevier.

and evolution is that the former is not degraded by O2 evolution and, in principle, different catalysts can be employed at the different electrodes, avoiding the need for both catalysts to be stable in the same voltage range. Although the basic mechanisms of O2 reduction and evolution in aqueous electrolytes are well known, the specific processes in those containing lithium salts have received much less attention and merit further study in view of the present interest in aqueous Li–O2 cells84. In this Review we have focused on alkaline electrolytes, as these have been most widely used in Li–O2 cells so far. However, acidic electrolytes may also be used41 and these give rise to higher voltages (up to ~4.25?V versus Li/Li+): 2Li?+??O2?+?2H+?→?2Li+?+?H2O.

The Li–S battery

The rechargeable Li–S cell is shown in Fig.? 1 and operates by reduction of S at the cathode on discharge to form various polysulphides that combine with Li to ultimately produce Li2S. Such cells have many attractive features, including: (i) the natural abundance and low cost of S; and (ii) high theoretical energy storage (Table?1)15–17. Yet the promise of a device with greater energy storage and cycle life than Li-ion has not yet materialized; even after decades of development, the Li–S battery has still not reached mass commercialization. Several problems inherent in the cell chemistry remain and are summarized in Fig.?7. Among such problems, discussed in detail in a previous review 16, are: (i) poor electrode rechargeability and limited rate capability 85,86 owing to the insulating nature of sulphur and the solid reduction products (Li2S and Li2S2); (ii) fast capacity fading owing to the generation of various soluble polysulphide Li2Sn (3?≤?n?≤?6) intermediates87–90, which gives rise to a shuttle mechanism91; and (iii) a poorly controlled Li/electrolyte interface. The shuttle mechanism arises because the soluble polysulphides that are formed at the cathode are transported to the anode where they are reduced to lower polysulphides, which are then transported
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back to the cathode, where they become reoxidized and then return to the anode. If, at the anode, reduction proceeds to form insoluble Li2S2 or Li2S, then this can deposit on the anode and elsewhere. Much of the recent work to improve Li–S cells builds on previous approaches. Considerable effort has been devoted to designing porous composite cathodes that are capable of delivering electrons efficiently to the S as well as trapping the soluble polysulphides. These aspects of Li–S battery research are described in the section comparing Li–O2 and Li–S. A different approach to the problem of minimizing transport of the soluble polysulphides from cathode to anode involves the use of organosulphur-based polymer systems with S–S linkages offering high specific energy and being capable of reversibly cleaving and reforming on reduction and oxidation in the molecular skeleton92–94. The charge/discharge reactions of the systems are based on the redox chemistry of thiolates (RS?), which can be oxidized to give the corresponding radical (RS), which can, in turn, couple to form disulphides (RSSR)95,96. Poly(2,2′-aminophenyl-disulphide) was an early example96. The redox chemistry of the dimercaptothiadiazole polymer has also been extensively studied and was shown to polymerize to form the highly insoluble polydisulphide. A marked advantage of conjugation with the electron-poor thiadiazole ring is a substantial increase in the discharge potential plateau by approximately 0.6–2.8? V. Asides from this, polyvinyl disulphide polymers were also shown to deliver a sustainable reversible capacity of 400? mAh? g?1 for at least 200? cycles97, and other polyvinyl sulphides containing more sulphur atoms in Sn units (2?<?n?<?7) are being studied at present. Last, it should be recalled that earlier, Degott reported92 polythiene-type conjugated polymers with the expected large capacity of 630?mAh?g?1. However, the polymer had poor kinetics owing to the undesirable crosslinks between chains. Incorporating this approach with mesoporous carbon electrodes may offer an interesting way forward.
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REVIEW ARTICLE
As for Li–O2, the Li-electrolyte interface constitutes another important challenge for the Li–S battery that is discussed in the section comparing Li–O2 and Li–S. Finding an electrolyte to combat the irreversible loss of sulphur associated with the formation of soluble Li2S6, Li2S4 and insoluble Li2S2 or Li2S is one of the greatest challenges for Li–S. Yet solutions remain elusive, even after intensive study for many years. Special attention has been given to cyclic or linear ethers such as tetrahydrofuran87,98–101 1,3-dioxolane23,99–101, dimethoxyethane23 and tetra(ethylene glycol)-dimethyl ether 102 because of the absence of chemical attack from S? centres. However, certain general trends have emerged: (i)?solvents with high permittivity and donor numbers solvate long- and short-chain lithium polysulphides; and (ii)?solvents with high viscosity tend to inhibit the dissociation of S82? and disproportionation of lithium polysulphide. Such differences explain the strong dependency of the shape of the voltage–capacity curve (for example, length of the first versus second plateaus) on the nature of the electrolyte. It is unlikely that a single solvent can fulfil all the requirements of an electrolyte for the Li–S battery, based on the long-standing research on mixtures of solvents103,104 or even the exploration of ionic liquids22. The likelihood of a breakthrough in electrolyte design, although not impossible, is slim unless fresh directions, such as reversible organic or inorganic species that chelate/bind sulphide species, are explored. Faced with the problem of liquid electrolytes, researchers have investigated solid-state Li–S batteries, using a wide variety of polymers (PEO-based)105 or gel (polyvinylidene-fluoride-based)20,106 electrolytes or inorganic glass–ceramic electrolytes (Li2S–P2S5) combined with a Cu and S positive electrode composite107. Success was limited, with cells suffering from rapid capacity decay or large polarization from the persistence of lithium dendrite formation and polysulphide dissolution. These problems were not alleviated by going to the solid state. Many strategies have focused on one component of the cell. In contrast, Scrosati and colleagues18 have adopted a holistic approach by modifying the entire cell. Their cell is assembled in the discharged state (Li2S) by using a lithium sulphide–carbon composite (made by ball milling) as the cathode. The Li negative electrode is replaced by a Sn–C–Li alloy with higher chemical stability towards sulphides and not prone to dendrite formation. The electrolyte is a polymer-gel membrane formed by ‘gelling’ with 1?M?LiPF6 in ethylene carbonate and dimethyl carbonate, saturated with Li2S, with a PEO and lithium trifluoromethane sulphonate (LiCF3SO3) polymer matrix incorporating a ZrO2 nanofiller. The cell can sustain a specific energy of ~1,100?Wh?kg?1 for tens of cycles, a value not previously
1,400 Speci?c capacity (mAh g–1) 1,200 1,000 800 600 400 200 0 0 5 10 Cycle number 15 20

NATURE MATERIALS DOI: 10.1038/NMAT3191
achieved for a Li-metal-free battery. The relatively large polarization (~2?V) between charge and discharge, which penalizes the energy efficiency of the system, could be addressed by substituting Sn for the Si electrode108. Although different in a number of ways, there are important similarities between some of the challenges facing aqueous and non-aqueous Li–O2 as well as Li–S, summarized in Figs?3, 5 and 7. All rely on porous positive electrodes. Several strategies have been explored to develop cathodes with engineered porosity and structure capable of efficient electron transport to the S and capture of the polysulphides formed on discharge; two of the most important challenges for Li–S cells. Most of the approaches, since the work of Peled? et? al. in 198890, have involved the fabrication of disordered mesoporous carbon–sulphur composites86,109–112 occasionally loaded with inorganic nanofillers, for example, Mg0.6Ni0.4O (ref.?113), Al2O3 (ref.? 114), or composites with sulphur embedded in electronically conducting polymers to immobilize polysulphides115–117. However, the most significant advance has been made by Nazar and colleagues17 who have demonstrated that cathodes based on ordered nanostructured mesoporous carbon–sulphur composites provide higher and more sustained, reversible capacities (Fig.?8). They also demonstrated the feasibility of further delaying the diffusion of polysulphide out of the cathode structure by functionalizing the pore surfaces of the carbon with polyethylene glycol (PEG) chains of varying molecular weights17. Subsequently, similar levels of performance were observed using porous carbonaceous composites118–120. Sandwich-type graphene-sheet–sulphur nanocomposites, where the expanded graphite or graphene layers act as sulphur microcontainers, exhibit excellent performance121,122. Overall, these improvements were ascribed to the intimate contact on the nanoscale between the insulating sulphur and the conductive carbon framework, while maintaining accessibility for the electrolyte. A further approach involving nanodimension sulphur–polythiophene core–shell or nanosulphur–(polypyrrole-co-aniline) composites123,124 exhibits a similar cycle life to the ordered mesoporous carbons. A recent advance that further stabilized the sulphur cathodes relies on the concept of using polysulphide reservoirs composed of porous silica19 or metal–organic framework additives125 embedded in the carbon– sulphur composite. Such approaches reduce the capacity fading to values of ~80% of the initial capacity after 80 cycles. This is a significant advance, yet is still not sufficient for most applications. For Li–S, the insulating S resides in the pores, which also trap the LinS discharge products. For Li–O2 the porous electrode must support diffusion of O2 gas, so that O2 is transported to the electrolyte/electrode interface as much as possible through the gas phase rather than by slower diffusion in the electrolyte. In the case of the non-aqueous cell, clogging of the pores by the solid discharge product, Li2O2, must be minimized. A variety of carbon substrates has been investigated by several authors43,49,54,62,76,126,127. Pores that are too small lead to clogging whereas those that are too large may inhibit recharging. Mesoporosity is expected to offer the best compromise and be able to accommodate the catalyst particles, which must be distributed across the pore surface. The deposition of an insulating solid such as Li2O2, if it formed a continuous film, would soon block the electrode surface (its growth limited by electron tunnelling, assuming ionic conductivity is lower than electronic in Li2O2). Controlling the morphology of the discharged products in the pores is therefore important 128. Ensuring some solubility for Li2O2 is an advantage as it would promote higher charge and discharge rates. It is also important to ensure that the electrolyte wets the pore surfaces. Most porous electrodes are based on carbon, because of its low cost, high conductivity and processability. Concerning the stability of the carbon, oxidation (corrosion) may occur if high voltages are accessed on charging. A recent study has indicated that
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Some common challenges of Li–O2 and Li–S

Figure 8 | Cycle-life data of Li–S cell. Comparison of Li–S cells with a cathode composed of S imbibed into the mesoporous carbon CMK-3 incorporating PEG (upper points) versus a similar cell but with a mechanical mixture of CMK-3 and S (lower points) at a rate of 168?mA?g–1 and at room temperature. Figure reproduced from ref.?17, ??2009 NPG.
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NATURE MATERIALS DOI: 10.1038/NMAT3191
carbon is relatively stable on reduction (discharge) in non-aqueous Li–O2 cells71. In the case of aqueous cells, as noted above, the preferred three-electrode cell configuration should minimize deposition of solid LiOH in the porous cathode. The main challenge here is therefore to design a good gas-diffusion electrode incorporating a low-cost reduction catalyst. The Li negative electrode, whether coupled with O2 or S as the positive electrode, suffers from several problems (Figs?3, 5 and 7). The growth of ‘mossy’ Li on charging requires excess Li to mitigate against loss on each cycle, adding a severe penalty to the overall energy density of the system. The limited amount of Li resources leads to concerns about Li shortage in view of the increasing demands of automotive transportation. The parasitic consumption of the metal can occur by oxidation from O2 crossover from the porous cathode or by reaction with (poly)sulphides leading to insoluble Li2S species at the negative electrode. Such issues must be addressed if we are to develop Li-lean and long-lasting Li–O2 or Li–S systems. A twin-track approach aimed at improving the Li electrode and exploring alternative anodes to Li metal should be adopted. Considering first improving the Li metal anode, neither the dynamics of the Li surface roughness on cycling nor the chemistry of the solvent reactions with Li have been fully addressed, despite considerable effort involving: (i) new anode designs; (ii) new surface treatments; (iii) Li+-conducting barriers; and (iv) the use of additives. Encouraging results were obtained by adding LiNO3 to the cathode or the electrolyte in Li–S cells129, which was shown to modify the solid electrolyte interphase on Li, thus increasing the cycling stability 24. Undoubtedly Li–O2 and Li–S cells will benefit greatly from the use of a Li+-conducting ceramic, such as LISICON, protecting the Li metal as it isolates the negative electrode chemistry from that at the positive electrode, hence widening the choice of solvents to be used, as they are no longer exposed to Li metal. This, for instance, enables the use of dimethylformamide or dimethylsulphoxide for Li–O2 cells. Additionally, the ceramic will suppress the shuttle mechanism in Li–S cells. The Li+-conducting ceramic protecting the Li anode is relatively thick (minimum 0.3? mm in the case of LISICON)130 and adds mass and volume to the cell. It is therefore important to develop membranes that are thinner. Furthermore, the LISICON ceramic is unstable in contact with Li and cycling the Li/ceramic interface is difficult. To solve these problems a Li+-conducting layer (typically based on a polymer, gel or non-aqueous liquid electrolyte) is placed between the Li anode and the ceramic26,40,41,131–133. The ceramic/electrolyte interface also has problems. In the case of Li–S cells, LISICON is partially reduced by the polysulphides, hence it is necessary to add a protective membrane or change the ceramic to enhance its stability against sulphur reduction. Similarly the ceramic/aqueous-electrolyte interface suffers from degradation at the extreme ends of the pH scale in the electrolyte in aqueous Li– O2 cells, and can become obstructed by deposition of LiOH39. To address this problem, a cation-exchange polymer has been applied to the surface of the ceramic that is exposed to the aqueous electrolyte39. Alternatively the electrolyte may be modified, for example, by addition of acetic acid41. Overall, these issues call for greater efforts in developing Li+-conducting ceramic electrolytes with enhanced chemical stability. A different approach to protecting the Li metal to improve its cyclability is to use a different negative electrode, for example, Si or Sn, which operate through alloying–dealloying reactions. Such electrodes would have to be prelithiated at some point in the cellassembly process, unless the cells could be assembled in the discharged state with Li2O2 or Li2S in the cathode in the case of Li–O2 and Li–S, respectively. However, for Li–S, this has proved to be problematic owing to the resulting poor reversibility and high polarization134. Such approaches do not prevent possible oxidation of the lithiated alloys by O2 crossover from the cathode in Li–O2 cells. For the cathode, a single discharge of a Li–S–O2 hybrid was
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REVIEW ARTICLE
recently demonstrated by Zhang and colleagues135. It is argued that the presence of O2 in the cell could be beneficial as it would oxidize the soluble polysulphide into insoluble sulphur, thus reducing self-discharge.

Summary

Li–S and Li–O2 cells both offer substantial increases in specific energy compared with Li-ion, but the gain in energy density is, at best, modest. Li–O2 has a higher specific energy than Li–S. O2 is free and S is cheap, and the main constituent of the cathode in both cases is C. All represent advantages in minimizing the cost of the cells compared with Li-ion. Li–S is closer to mass commercialization than Li–O2, but the problems of Li–S have been known for many years and have not yet been completely solved. Li–O2 has received much less attention than Li–S until recently, which means, as yet, no intractable problems have been identified, but of course such problems may be discovered in the future. It is important that future work focuses on understanding the fundamental chemistry taking place in the Li–O2 and Li–S cells if progress is to be made; otherwise any attempt to develop a technology is likely to fail. The recent observation that the operation of Li–O2 cells with organic carbonate electrolytes is dominated by decomposition reactions and not Li2O2 formation highlights the importance of fundamental studies. There is a degree of commonality in some of the problems facing Li–S and Li–O2, especially concerning the Li anode and the porous C cathode. It is not yet clear if any of the advanced batteries that offer to take us beyond Li-ion will become a commercial success. However, despite the difficulties, our society needs energy-storage devices with much higher levels of energy storage than ever before. Li–O2 and Li–S batteries are among the few contenders that can exceed the stored energy of Li-ion. We must devote more intensive research to address the problems of the Li–O2 and Li–S batteries.

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Acknowledgements

P.G.B. is indebted to the EPRSC and Toyota Motor Europe for support. The authors wish to express their thanks to S.?Visco, M.?Armand and R.?Demir-Cakan and the ALISTOREERI members for helpful discussions. P.G.B. and J.M.T. are members of ALISTOREERI?—?European Network of Excellence on Lithium Batteries.

Additional information

The authors declare no competing financial interests.

29

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