Stefan A. Freunberger

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

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, extended-range 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 (O(2)) 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.

Challenges facing lithium batteries and electrical double-layer capacitors.

Energy-storage technologies, including electrical double-layer capacitors and rechargeable batteries, have attracted significant attention for applications in portable electronic devices, electric vehicles, bulk electricity storage at power stations, and load leveling of renewable sources, such as solar energy and wind power. Transforming lithium batteries and electric double-layer capacitors requires a step change in the science underpinning these devices, including the discovery of new materials, new electrochemistry, and an increased understanding of the processes on which the devices depend. The Review will consider some of the current scientific issues underpinning lithium batteries and electric double-layer capacitors.

A reversible and higher-rate Li-O2 battery.

Improving Lithium Batteries Lithium-oxygen batteries have similar volumetric energy densities to lithium-ion batteries, but, because the oxygen part of the battery can be extracted from the air, they have a significant advantage in their gravimetric energy densities. One of the fundamental problems plaguing the nonaqueous Li-O2 system is that the Li2O2 that forms on discharge must be completely reversed on charging, but for most systems, a range of side products form instead of Li2O2. Peng et al. (p. 563, published online 19 July) show that by using dimethyl sulfoxide as the electrolyte, and a porous gold cathode, they can get reversible production and removal of Li2O2 during discharge and charge cycles. Furthermore, the electrolyte-electrode system operates with much faster kinetics than carbon electrodes. A viable lithium-oxygen battery is demonstrated using dimethylsulfoxide electrolyte and a porous gold cathode. The rechargeable nonaqueous lithium-air (Li-O2) battery is receiving a great deal of interest because, theoretically, its specific energy far exceeds the best that can be achieved with lithium-ion cells. Operation of the rechargeable Li-O2 battery depends critically on repeated and highly reversible formation/decomposition of lithium peroxide (Li2O2) at the cathode upon cycling. Here, we show that this process is possible with the use of a dimethyl sulfoxide electrolyte and a porous gold electrode (95% capacity retention from cycles 1 to 100), whereas previously only partial Li2O2 formation/decomposition and limited cycling could occur. Furthermore, we present data indicating that the kinetics of Li2O2 oxidation on charge is approximately 10 times faster than on carbon electrodes.

Reactions in the Rechargeable Lithium–O2 Battery with Alkyl Carbonate Electrolytes

The nonaqueous rechargeable lithium-O(2) battery containing an alkyl carbonate electrolyte discharges by formation of C(3)H(6)(OCO(2)Li)(2), Li(2)CO(3), HCO(2)Li, CH(3)CO(2)Li, CO(2), and H(2)O at the cathode, due to electrolyte decomposition. Charging involves oxidation of C(3)H(6)(OCO(2)Li)(2), Li(2)CO(3), HCO(2)Li, CH(3)CO(2)Li accompanied by CO(2) and H(2)O evolution. Mechanisms are proposed for the reactions on discharge and charge. The different pathways for discharge and charge are consistent with the widely observed voltage gap in Li-O(2) cells. Oxidation of C(3)H(6)(OCO(2)Li)(2) involves terminal carbonate groups leaving behind the OC(3)H(6)O moiety that reacts to form a thick gel on the Li anode. Li(2)CO(3), HCO(2)Li, CH(3)CO(2)Li, and C(3)H(6)(OCO(2)Li)(2) accumulate in the cathode on cycling correlating with capacity fading and cell failure. The latter is compounded by continuous consumption of the electrolyte on each discharge.

The lithium-oxygen battery with ether-based electrolytes.

The rechargeable Li–air (O2) battery is receiving a great deal of interest because theoretically it can store significantly more energy than lithium ion batteries, thus potentially transforming energy storage. Since it was first described, a number of aspects of the Li–O2 battery with a non-aqueous electrolyte have been investigated. The electrolyte is recognized as one of the greatest challenges. To date, organic carbonate-based electrolytes (e.g. LiPF6 in propylene carbonate) have been widely used. However, recently, it has been shown that instead of O2 being reduced in the porous cathode to form Li2O2, as desired, discharge in organic carbonate electrolytes is associated with severe electrolyte decomposition. As a result it is very important to investigate other solvents in the search for a suitable electrolyte. In this regard much attention is now focused on electrolytes based on ethers (e.g. tetraglyme (tetraethylene glycol dimethyl ether)). Ethers are attractive for the Li–O2 battery because they are one of the few solvents that combine the following attributes: capable of operating with a lithium metal anode, stable to oxidation potentials in excess of 4.5 V versus Li/Li, safe, of low cost and, in the case of higher molecular weights, such as tetraglyme, they are of low volatility. Crucially, they are also anticipated to show greater stability towards reduced O2 species compared with organic carbonates. Herein we show that although the ethers are more stable than the organic carbonates, the Li2O2 that forms on the first discharge is accompanied by electrolyte decomposition, to give a mixture of Li2CO3, HCO2Li, CH3CO2Li, polyethers/ esters, CO2, and H2O. The extent of electrolyte degradation compared with Li2O2 formation on discharge appears to increase rapidly with cycling (that is, charging and discharging), such that after only 5 cycles there is little or no evidence of Li2O2 from powder X-ray diffraction. We show that the same decomposition products occur for linear chain lengths other than tetraglyme. In the case of cyclic ethers, such as 1,3dioxolane and 2-methyltetrahydrofuran (2-Me-THF), decomposition also occurs. For 1,3-dioxolane, decomposition forms polyethers/esters, Li2CO3, HCO2Li, and C2H4(OCO2Li)2, and for 2-Me-THF the main products are HCO2Li, CH3CO2Li; in both cases CO2 and H2O evolve. The results presented herein demonstrate that ether-based electrolytes are not suitable for rechargeable Li–O2 cells. A Li–O2 cell consisting of a lithium metal anode, an electrolyte, comprising 1m LiPF6 in tetraglyme, and a porous cathode (Super P/Kynar) was constructed as described in the Experimental Section. The cell was discharged in 1 atm O2 to 2 V. The porous cathode was then removed, washed with CH3CN, and examined by powder X-ray diffraction (PXRD) and FTIR. The results are presented in Figure 1 and Figure 2. The PXRD data demonstrate the presence of Li2O2, consistent with previous PXRD data for a Li–O2 cell with a tetraglyme electrolyte at the end of the first discharge. However, examination of the FTIR spectra, Figure 2, reveals that, in addition to Li2O2, other products form. Although the FTIR spectra provide clear evidence of electrolyte decom-

The carbon electrode in nonaqueous Li-O2 cells.

Carbon has been used widely as the basis of porous cathodes for nonaqueous Li-O(2) cells. However, the stability of carbon and the effect of carbon on electrolyte decomposition in such cells are complex and depend on the hydrophobicity/hydrophilicity of the carbon surface. Analyzing carbon cathodes, cycled in Li-O(2) cells between 2 and 4 V, using acid treatment and Fentons reagent, and combined with differential electrochemical mass spectrometry and FTIR, demonstrates the following: Carbon is relatively stable below 3.5 V (vs Li/Li(+)) on discharge or charge, especially so for hydrophobic carbon, but is unstable on charging above 3.5 V (in the presence of Li(2)O(2)), oxidatively decomposing to form Li(2)CO(3). Direct chemical reaction with Li(2)O(2) accounts for only a small proportion of the total carbon decomposition on cycling. Carbon promotes electrolyte decomposition during discharge and charge in a Li-O(2) cell, giving rise to Li(2)CO(3) and Li carboxylates (DMSO and tetraglyme electrolytes). The Li(2)CO(3) and Li carboxylates present at the end of discharge and those that form on charge result in polarization on the subsequent charge. Li(2)CO(3) (derived from carbon and from the electrolyte) as well as the Li carboxylates (derived from the electrolyte) decompose and form on charging. Oxidation of Li(2)CO(3) on charging to ∼4 V is incomplete; Li(2)CO(3) accumulates on cycling resulting in electrode passivation and capacity fading. Hydrophilic carbon is less stable and more catalytically active toward electrolyte decomposition than carbon with a hydrophobic surface. If the Li-O(2) cell could be charged at or below 3.5 V, then carbon may be relatively stable, however, its ability to promote electrolyte decomposition, presenting problems for its use in a practical Li-O(2) battery. The results emphasize that stable cycling of Li(2)O(2) at the cathode in a Li-O(2) cell depends on the synergy between electrolyte and electrode; the stability of the electrode and the electrolyte cannot be considered in isolation.

A stable cathode for the aprotic Li–O2 battery

Rechargeable lithium-air (O2) batteries are receiving intense interest because their high theoretical specific energy exceeds that of lithium-ion batteries. If the Li-O2 battery is ever to succeed, highly reversible formation/decomposition of Li2O2 must take place at the cathode on cycling. However, carbon, used ubiquitously as the basis of the cathode, decomposes during Li2O2 oxidation on charge and actively promotes electrolyte decomposition on cycling. Replacing carbon with a nanoporous gold cathode, when in contact with a dimethyl sulphoxide-based electrolyte, does seem to demonstrate better stability. However, nanoporous gold is not a suitable cathode; its high mass destroys the key advantage of Li-O2 over Li ion (specific energy), it is too expensive and too difficult to fabricate. Identifying a suitable cathode material for the Li-O2 cell is one of the greatest challenges at present. Here we show that a TiC-based cathode reduces greatly side reactions (arising from the electrolyte and electrode degradation) compared with carbon and exhibits better reversible formation/decomposition of Li2O2 even than nanoporous gold (>98% capacity retention after 100 cycles, compared with 95% for nanoporous gold); it is also four times lighter, of lower cost and easier to fabricate. The stability may originate from the presence of TiO2 (along with some TiOC) on the surface of TiC. In contrast to carbon or nanoporous gold, TiC seems to represent a more viable, stable, cathode for aprotic Li-O2 cells.

Charging a Li–O2 battery using a redox mediator

The non-aqueous Li-air (O2) battery is receiving intense interest because its theoretical specific energy exceeds that of Li-ion batteries. Recharging the Li-O2 battery depends on oxidizing solid lithium peroxide (Li2O2), which is formed on discharge within the porous cathode. However, transporting charge between Li2O2 particles and the solid electrode surface is at best very difficult and leads to voltage polarization on charging, even at modest rates. This is a significant problem facing the non-aqueous Li-O2 battery. Here we show that incorporation of a redox mediator, tetrathiafulvalene (TTF), enables recharging at rates that are impossible for the cell in the absence of the mediator. On charging, TTF is oxidized to TTF(+) at the cathode surface; TTF(+) in turn oxidizes the solid Li2O2, which results in the regeneration of TTF. The mediator acts as an electron-hole transfer agent that permits efficient oxidation of solid Li2O2. The cell with the mediator demonstrated 100 charge/discharge cycles.

The role of LiO2 solubility in O2 reduction in aprotic solvents and its consequences for Li-O2 batteries.

When lithium-oxygen batteries discharge, O2 is reduced at the cathode to form solid Li2O2. Understanding the fundamental mechanism of O2 reduction in aprotic solvents is therefore essential to realizing their technological potential. Two different models have been proposed for Li2O2 formation, involving either solution or electrode surface routes. Here, we describe a single unified mechanism, which, unlike previous models, can explain O2 reduction across the whole range of solvents and for which the two previous models are limiting cases. We observe that the solvent influences O2 reduction through its effect on the solubility of LiO2, or, more precisely, the free energy of the reaction LiO2(*)u2005⇌u2005Li(sol)(+)u2009+u2009O2(-)(sol)u2009+u2009ion pairsu2009+u2009higher aggregates (clusters). The unified mechanism shows that low-donor-number solvents are likely to lead to premature cell death, and that the future direction of research for lithium-oxygen batteries should focus on the search for new, stable, high-donor-number electrolytes, because they can support higher capacities and can better sustain discharge.

Oxygen reactions in a non-aqueous Li+ electrolyte.

Oxygen (O2) reduction is one of the most studied reactions in chemistry.1 Widely investigated in aqueous media, O2 reduction in non-aqueous solvents, such as CH3CN, has been studied for several decades.2–7 Today, O2 reduction in non-aqueous Li+ electrolytes is receiving considerable attention because it is the reaction on which operation of the Li–air (O2) battery depends.8–29 The Li–O2 battery is generating a great deal of interest because theoretically its high energy density could transform energy storage.8,u20099 As a result, it is crucial to understand the O2 reaction mechanisms in non-aqueous Li+ electrolytes. Important progress has been made using electrochemical measurements including recently by Laoire etu2005al.29 No less than five different mechanisms for O2 reduction in Li+ electrolytes have been proposed over the last 40u2005years based on electrochemical measurements alone.25–29 The value of using spectroelectrochemical methods is that they can identify directly the species involved in the reaction. Here we present in situ spectroscopic data that provide direct evidence that LiO2 is indeed an intermediate on O2 reduction, which then disproportionates to the final product Li2O2. Spectroscopic studies of Li2O2 oxidation demonstrate that LiO2 is not an intermediate on oxidation, that is, oxidation does not follow the reverse pathway to reduction.