DOI: 10.1002/celc.201402315
Minireviews
Lithium–Air Batteries: Performance Interplays with
Instability Factors
Luhan Ye,[a] Weiqiang Lv,[a] Junyi Cui,[a] Yachun Liang,[a] Peng Wu,[a] Xiaoning Wang,[a]
Han He,[a] Senjun Lin,[b] Wei Wang,[c] James H. Dickerson,[d, e] and Weidong He*[a, f, g]
Lithium–air batteries are considered to be promising electrochemical storage devices, due to their high specific energy
density. However, instability limits their cyclic performance and
rate capacity and also leads to a high overpotential; lithium–
air batteries are typically characterized by capacity degradation
and short cycle life. Such challenges prevent lithium–air batteries from entering and competing in the battery market. Electrodes, organic solvents, the interface between electrolyte and
cathode, and ambient conditions have all been demonstrated
to impact substantially the stability of the lithium–air battery.
In this Minireview, we focus on electrode and electrolyte decomposition, side reactions, and physical mass transport in
aprotic lithium–air batteries, as well as other types of lithium–
air batteries, and aim to understand comprehensively their performance and association with instability factors.
1. Introduction
The fossil-fuel-based economy has become increasingly unsustainable over the past decades. The emerging energy crisis has
affected various aspects of human society. To meet future
social and industrial energy demands, scientists are facing unprecedented challenges. Much effort has been focused on the
development of sustainable energy devices including solar
cells, and those driven by biomass and hydrogen.[1–4] How to
produce power efficiently has been a central theme throughout the industrial age. In addition, for better energy utilization,
another strategy is to seek new methods for efficient energy
storage.[5, 6] Until now, Li-ion batteries have dominated the bat-
tery market. They have particularly long cycle lives and high
operating voltages, and are thus suitable for certain practical
applications such as powering laptops and cellphones.[7] Although the energy density of Li-ion batteries ( 160 W h kg1)
is relatively high compared to traditional batteries, they are
still far from being applied in electrical vehicles. Cost-effective
and environmentally friendly storage devices with high energy
densities are urgently needed. The lithium–air battery, due to
its excellent energy density, can potentially meet this need.[8, 9]
The first rechargeable nonaqueous lithium–air battery was introduced by Abraham and Jiang in 1996.[10] The theoretical
energy density has been demonstrated to be up to
11 430 W h kg1, which is comparable to fossil fuels and is
much higher compared with other widely used batteries (such
as the Li-ion and Ni–Cd batteries), as shown in Figure 1.[9] The
high energy density of lithium–air batteries is promising for
powering electric vehicles and hybrid electric vehicles. Despite
their high energy density, the applications of lithium–air batteries are still limited significantly by their instability.[11–16] Compared to the lead–acid battery widely used in industry, the performance of lithium–air batteries is affected directly by their instability. Conventional lithium–air batteries, including nonaqueous and aqueous types, are assembled as shown in
Figure 2.[17] Much work has demonstrated their instability
during charge–discharge processes between the application to
internal structures; a subtle disturbance can lead to a large
degradation in capacity and poor cyclic performance.[18–20]
For instance, the main reaction product, Li2O2, is insoluble in
organic solvents, and might consequently clog electrode
pores. Furthermore, Bruce et al.[21] reported that carbon electrodes are unstable as the charge voltage increases beyond
3.5 V. Electrolytes also decompose during operation cycles.[19]
In brief, overall physical transportation problems, decomposition and parasitic reactions cause a sensitive instability in lithi-
[a] L. Ye, W. Lv, J. Cui, Y. Liang, P. Wu, X. Wang, H. He, Prof. W. He
School of Energy Science and Engineering
University of Electronic Science and Technology
Chengdu, Sichuan 611731 (P.R. China)
[b] S. Lin
Department of Industrial Design
Zhejiang University of Technology
Hangzhou, Zhejiang, 310014 (P.R. China)
[c] W. Wang
Department of Material Science and Engineering
Shenzhen Graduate School, Harbin Institute of Technology
Shenzhen 518055 (P.R. China)
[d] J. H. Dickerson
Center for Functional Nanomaterials
Brookhaven National Laboratory
Upton, NY 11973 (USA)
[e] J. H. Dickerson
Department of Physics, Brown University
Providence, RI 02912 (USA)
[f] Prof. W. He
Interdisciplinary Program in Materials Science
Vanderbilt University, Nashville, TN 37234-0106 (USA)
[g] Prof. W. He
Vanderbilt Institute of Nanoscale Science and Engineering
Vanderbilt University, Nashville, TN 37234-0106 (USA)
E-mail: [email protected]
ChemElectroChem 2015, 2, 312 – 323
312
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
2. Electrodes
An electrode is the intermediate bridge between the outside
atmosphere and the inner microstructure of the lithium–air
battery. Electrochemical performance correlates with Li2O2 deposition, air diffusion and decomposition in an electrode. The
cathodes utilized in lithium–air batteries are typically catalystbased porous carbons. Much effort has been devoted to increasing the capacity retention and lowering the kinetic overpotential.[22–26] Existing literature has also demonstrated that
the cyclical stability can be improved by controlling the cathode reaction and optimizing the cathode morphology.[25] Regarding the mechanism, Xiao et al.[27] have investigated various
carbon materials to explore the effect of pore microstructure
as well as porosity and thickness. Among all the studied
carbon sources, the highly porous Ketjen black (KB)-based electrode gave the best performance after assembly (Figure 3 and
Table 1). Due to the high porosity of KB-based electrodes and
their unique chainlike structure, such lithium–air batteries ex-
Figure 1. The gravimetric energy densities (W h kg1) for various types of rechargeable batteries compared with gasoline. Adapted from Ref. [9].
Figure 2. Aprotic (left) and aqueous (right) lithium–air batteries. Adapted
from Ref. [17].
um–air batteries. The stability is correlated with the current
density, specific capacity and cycle life. Instability in a battery
system leads to a high overpotential and consequently reduces
the energy output. To tackle the instability issues, many improved materials and protected structures have been proposed
recently, including the lithium anode with a protective membrane architecture and stable ionic liquid (IL) electrolyte.[11, 12]
Although much attention has been paid to the development
of long-term-stable lithium–air batteries by using various approaches, instability issues remain to be addressed in this field.
As these unstable systems inevitably undergo rapid fading of
capacity and high energy loss, the lithium–air battery will be of
little practical significance unless the overall instability issue is
properly resolved.
In summary, to obtain a high-performance lithium–air battery, instability is the key obstacle to be tackled. In this Minireview, the instability associated with the different parts of lithium–air batteries, including the electrode, the electrolyte and
the ambient environment, is discussed, and insights into the
development of advanced stable materials and structures are
provided. The aim of this Minireview is to give the reader
a comprehensive understanding on the stability of lithium–air
batteries.
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
Figure 3. Comparison of the discharge capacities of lithium–air batteries
using different carbon sources: Ketjen black (KB), Calgon, ball-milled KB,
BP2000, JMC (mesoporous carbon), and Denka. The preparative processes
are presented in Ref. [27].
Table 1. Surface area and pore volume comparison. Pore size distribution
was evaluated using the Barrett–Joyner–Halenda (BJH) method. Adapted
from Ref. [27].
313
Pore volume BJH pore
Surface
size [nm]
area [m2 g1] [cm3 g1]
Microstructure
from XRD
KB
2672
7.6510
Ballmilled KB
BP2000
342.4
0.4334
poorly crystalline
graphite
amorphous
1567
0.8350
Calgon
1006
0.5460
Denka
black
JMC
102.0
0.5355
548.7
0.2376
2.217–
15.000
no clear
peak
no clear
peak
no clear
peak
2.511 &
6.000
3.0–3.8
poorly crystalline
graphite
crystalline
graphite
poorly crystalline
graphite
amorphous with
ordered
mesopores
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
hibit high specific capacity. The results indicate that the
volume is sufficient to provide extra reaction and storage regions, which is favorable for achieving high battery stability.
Tran et al.[28] also reported that the capacity of batteries is
largely linearly correlated with electrode pore size (Figure 4).
Figure 5. Various growth modes of discharge products: a) cylindrical film,
b) spherical film, and c) planar film. Adapted from Ref. [31].
or adhere to reaction sites, which influences in part the battery
performance. More importantly, the Li2O2 layer restricts the
electrical conductivity of electrodes. It has been demonstrated
that a certain amount of Li vacancies in the layer enable the
surface to be conductive.[37] Radin et al.[38] used first-principles
calculations to explore the surface properties of Li2O2 layers,
and the results indicate that Li2O2 has pronounced conductivity, whereas no conduction was observed in Li2O or LiCO3.
Hummelshøj et al.[39] proposed a model to describe the Li2O2
deposition and its influence on conductivity. Density functional
theory (DFT) was used to identify the origin of the overpotential. In their report, the discharge overpotential was determined to be 0.43 V for discharging and 0.6 V for charging resulting from surface deposition. Albertus et al.[36] showed
a gradual decrease of discharge voltage from an initial 2.6 V,
and a sudden drop to approximately 2.0 V. Furthermore, Viswanathan et al.[35] proposed a first-principles metal–insulator–
metal (MIM) charge-transport model to probe the Li2O2 surface
conductivity (Figure 6). They stated that the tunneling current
should support a high operating current, and that the MIM interface is critical for obtaining high stability. They described
a dramatic drop—termed “sudden death”—that occurs if the
deposition layer thickness is in the range of approximately 5–
10 nm. Once the operating current exceeds the tunneling current, the sudden death occurs.
Figure 4. Discharge time and specific capacity as a function of average pore
diameter. Adapted from Ref. [28].
This work indicates that the performance of the lithium–air
battery is largely affected by the porosity and pore size of electrodes as well as the cathode reaction. An electrode with high
porosity is afforded efficient diffusion, whereas sufficiently
large pores are preferable for Li2O2 accommodation. In this section, we will discuss first product deposition and then carbon
decomposition. Following on, some advanced cathode materials are introduced.
2.1. Li2O2 Deposition
In an organic solvent, the main product Li2O2 is insoluble, and
upon discharge Li2O2 deposits on the electrode pore surface to
form a thin layer (Figure 5).[29–31] The deposition layer can
1) clog a sufficient number of pores to stop the liquid from
permeating, 2) limit the air diffusion rate, and 3) lead to low
charge–discharge efficiency. Among these problems, the first
two result in energy loss and the third leads to a sudden discharge voltage drop as well as large polarization, which makes
the practical capacity much smaller than the theoretical
value.[28–36] Tests on various carbon-based air electrodes have
revealed that expansion of the Li2O2 storage region can deliver
higher performance compared to smaller areas, indicating that
the accommodation of Li2O2 requires abundant volume.[37]
Both experiment and numerical simulation show high overpotentials during charge–discharge cycles, induced mainly by surface deposition. Here, the interface passivation in performance
variation is dominant compared with solvent and O2 transport
issues, and polarization is also mainly due to passivation.[35]
Deposition layers can obstruct the contacts between O2 and
catalyst/electrode. The pore surface provides active reaction
sites for O2 reduction. O2 is difficult to sufficiently transport in
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
2.2. Carbon Electrodes
Carbon is widely used as cathode material because it is cheap
and easily incorporated into a porous electrode.[40–42] However,
its actual effectiveness is often debated, due to its poor stability at the micro-/nanoscale.[19, 21] The electron/oxygen ratio is
near the ideal stoichiometric value of 2.000, based on quantitative differential electrochemical mass spectrometry (DEMS),
and most side reactions occur after the first discharge.[19, 43–44]
The intermediate radical product O2 and dominant reaction
314
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
decomposition of carbon, it begins to accumulate above 3.5 V,
especially during charging.
The reactions of Li2O2, C, and Li2CO3 result in the formation
of a thin layer at the cathode–Li2O2 interface. The formation of
CO2 indicates that Li2CO3 is formed simultaneously and is partially reduced. The electrochemical reduction of Li2CO3 proceeds continuously during charging. However, such reduction
is incomplete. In contrast to Li2O2, Li2CO3 is completely insulating and precipitates on the reaction sites causing electrode
passivation. However, only a small proportion of Li2CO3 formation occurs by the reaction between Li2O2 and carbon during
cycling.[21] In subsequent cycles, the electrolyte and carbon
electrode decompose continuously to form Li2CO3 and lithium
carboxylates. Once carbon starts to form Li2CO3, the charging
voltage rises rapidly. An experiment using a porous gold cathode showed less electrolyte decomposition compared with the
carbon-based cathode in the same electrolyte (DMSO), which
indicates that carbon not only decomposes itself but also promotes the electrolyte decomposition.[21, 45] In addition, the hydrophobicity/hydrophilicity of the surface also impacts the stability of the carbon and electrolyte.[21] Batteries based on hydrophobic carbon are more stable compared to hydrophilic
carbons. This is possibly due to the existence of CO, COOH,
and COH groups. In summary, the direct reaction between
Li2O2 and carbon is destructive when the charging potential is
greater than 3.5 V. Accumulation of Li2CO3 from carbon decomposition leads to high polarization and also stimulates
electrolyte decomposition.
Figure 6. Typical conductivity calculation setup for a) Au j Li2O2 j Au and
b) Au j Li2O2–LiO2 j Au. Adapted from Ref. [35].
product Li2O2 are both reactive, and can attack carbon or electrolyte. McCloskey et al.[44] revealed that Li2CO3 and LiRCO3 (R =
alkyl) are produced at the carbon–Li2O2 interface and Li2O2–
electrolyte interface, respectively. Such “interfacial carbonate
problems” lead to an extra overpotential (Figure 7). When
charged, carbon reacts chemically with Li2O2 in a high-oxidizing environment according to 2 Li2O2 + 2 C + O2 !2 Li2CO3 or
2 Li2O2 + C!Li2O + Li2CO3. 13C was used to investigate the generation of Li2CO3, as it could be distinguished from electrolyte.
Thotiyl et al.[21] suggested that the reactions between Li2O2 and
carbon occur mostly at high charging potentials (U > 3.5 V).
They stated that during the first discharge, electrode decomposition is subtle, whereas the side reactions involving the decomposition of the electrolyte are dominant. Interestingly, following cycles of charge–discharge, the amount of Li2CO3 increases with increasing voltage. As Li2CO3 is formed from the
2.3. Recent Stable Electrodes
As carbon is unstable in a highly oxidizing environment, and
its decomposition promotes the decomposition of the electrolyte, research has focused on new materials to replace carbon.
One example of an air electrode is the nanoporous gold
(NPG) cathode, which was proposed by Bruce et al.[45] After
several cycles, such cathodes show great stability (Figure 8).
NPG cathodes in a DMSO electrolyte can retain 95 % of their
capacity after 100 cycles. During charge and discharge, the reaction is dominated by Li2O2 formation/decomposition. Furthermore, the NPG cathode promotes efficiently Li2O2 formation/decomposition. The kinetics of Li2O2 oxidation was demonstrated to be tenfold faster than that of normal cathodes.
NPG cathodes are known as “concept cathodes” because they
are currently too expensive to be industrially manufactured.
Furthermore, the high mass fraction of gold compromises one
of the most important merits—the high specific energy density—rendering batteries containing these cathodes of little use.
However, NPG cathodes represent significant progress towards
the future of cathode materials by enabling fabrication of
stable Li–O2 batteries with a stable cathode.
In a recent report, TiC-based cathodes proposed by Thotiyl
et al.[46] outperformed a NPG cathode (Figure 9). TiC is four
times lighter than an NPG cathode and shows more robust stability. Fewer side reactions were detected at the cathode–electrolyte interface. The authors showed that the TiC-based Li–O2
battery retained capacity of greater than 98 % after 100 cycles
Figure 7. Upper panels indicate the deposition events during charging that
cause the rising charging potential. An approximately single monolayer of
Li2CO3 forms at the carbon surface, some dispersed carbonate possibly
forms as a result of electrolyte decomposition in the Li2O2 deposit, and an
approximately single monolayer forms at the electrolyte interface during
charging. e represents an unspecified electrochemical reaction that produces carbonate at the interface during charging. The dashed arrows indicate
qualitatively the charging potential appropriate for the three panels. Adapted from Ref. [44].
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
315
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
were collected. At the beginning of discharge, the amount of
TiO2 increases noticeably and it gradually becomes the main
substance on the surface. XPS results indicated that a thin
layer containing TiO2 and TiOC formed on the surface of the
cathode, which was believed to be responsible for the stability
of the cathode.
In addition, the investigation into the application of different
allotropic carbon nanoforms is also a promising strategy. Despite the cathode reaction, advanced carbon electrodes with
special bimodal pore size distributions are promising candidates for long-term-stable lithium–air batteries. Among various
allotropic carbons, carbon nanotubes, carbon nanoballs and
graphene have been widely investigated.[22–24] For instance,
Xiao et al.[22] fabricated hierarchically porous graphene as the
electrode of the lithium–air battery. Using these electrodes
they obtained a capacity of 15 000 mA h g1. The distributions
of pore sizes in this electrode are binomial, which is different
from traditional electrodes. The improved properties were also
demonstrated in high-performance carbon nanoball electrodes.[23] A capacity larger by 30 %–40 % compared with KB was
obtained with a carbon nanoball electrode. Carbon nanotubes,
similarly, due to their large pores and high porosity also gave
superior performance.[24] In these electrodes, large pores can
accommodate the deposition products and small pores can facilitate air diffusion. However, a thorough investigation into
the stability of such electrodes is required to improve the cathode reaction.
Figure 8. Charge–discharge curves (top) and cycling profile (bottom) for
a Li–O2 cell with a 0.1 m LiClO4/DMSO electrolyte and an NPG cathode, at
a current density of 500 mA g1 (based on the mass of Au). Adapted from
Ref. [45].
(compared with 95 % for NPG-cathode-based batteries). The results of Fourier transform infrared (FTIR) spectroscopy demonstrated that Li2O2 is the dominant discharge product during
charge–discharge. Only a minor amount of carboxylates produced by electrolyte decomposition was detected. To investigate the origin of the stability, X-ray photoelectron spectroscopy (XPS) data on the TiC cathode at every stage of the reaction
3. Electrolyte
3.1. Organic Solvents
Figure 9. Cycling curves and capacity retention of TiC cathodes. a) Galvanostatic discharge–charge cycles recorded
in 0.5 m LiClO4 in DMSO at a geometric current density of 1 mA cm2. b) Capacity retention for the same cell as in
(a). c) Galvanostatic discharge–charge cycles recorded in 0.5 m LiPF6 in TEGDME at a geometric current density of
0.5 mA cm2. d) Capacity retention for the same cell as in (c). Adapted from Ref. [46].
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
316
Lithium metal is a reactive material that needs to be protected
in batteries. The direct contact
of lithium with an aqueous solution results in the reaction 2 Li +
Lithium
2 H2O!H2› + 2 LiOH.
metal reacts rapidly if aqueous
electrolytes are used (discussed
in following section) unless
a membrane separator is used.
Current aprotic lithium–air batteries usually use organic salts
dissolved in organic solvents as
the electrolytes. The first-generation electrolyte applied in the
lithium–air battery was ethylene
carbonate associated with LiPF6,
introduced by Abraham in
1996.[10] For the first time, successful recharge performance
and a relatively high capacity
(800 mA h g at 1 mA cm2) were
achieved. However, poor cyclic
performance and low coulombic
efficiency limited the widespread
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
application of such a carbonate solvent. Propylene carbonate
(PC), ethylene carbonate (EC), dimethyl carbonate (DMC), ethyl
methyl carbonate (EMC), or their mixtures are substitutes proposed by others.[47] For example, using Super P as the air electrode and PC as solvent, Read[43] obtained the capacity of
1934 mA h g at a current density of 0.05 mA cm1. Sun et al.[41]
loaded CoO nanoparticles onto a mesoporous carbon (CMK-3)
cathode, with PC as the solvent. The cycle performance was
better than previously reported results, and a 95 % capacity retention was observed after 15 cycles. Cyclic performance and
voltaic efficiency were increased correspondingly, but were still
far from the standard required for practical application. Further
research on organic carbonate solvents revealed the reaction
mechanisms during the oxygen reduction reaction (ORR) and
oxygen evolution reaction (OER), and eventually showed the
drawbacks of using such solvents. The mechanism was shown
to involve attack on the electrode or electrolyte by superoxide
radical anions or peroxide anions produced during discharge
attacks the electrode or electrolyte (Scheme 1). Such a mecha-
Figure 10. FTIR spectrum of a composite electrode (SuperP/R-MnO2/Kynar)
discharged in 1 m LiPF6/PC to 2 V in O2. The reference spectra for Li2O2 (small
impurity peaks at 1080, 1450, and 1620 cm1), Li2CO3, and the electrode
before discharge are also shown. Adapted from Ref. [51].
Scheme 1. Molecular oxygen reduction in aprotic systems. Adapted from
Ref. [48] with permission. Copyright 2011 Wiley-VCH.
nism was then accepted widely to be the reaction principle for
electrolyte decomposition, which also demonstrates that alky
carbonate should not be used in long cycle lithium-air batteries. It has been widely proposed that during discharge, radical
O2 is the intermediate product.[14, 48] Due to the reactivity of O2
and the subsequent radical product O2
2 , the solvent decomposes rapidly during the subsequent discharge process.[19, 49, 50]
Freunberger et al.[51] reported side reactions with the formation
of C3H6(OCO2Li)2, HCO2Li, CH3CO2Li, CO2, and water during discharge (Figure 10). After several cycles, side products accumulate rapidly. The same results were also observed by
others.[52, 53] This stability issue over highly active radicals is pivotal for the electrolyte test. As the organic carbonates have
been associated with disastrous drawbacks, much effort has
been focused on other solvents.
Ethers, used widely after the alkyl carbonates were proposed, were investigated by others.[54, 55] Because of their stability at a high charge potential, ethers were considered originally
as promising electrolyte solvents. However, as research continued, the instability issues of solvents were demonstrated to
still exist even for ethers. Although ether moieties are generally
more stable than carboxyl groups, the cyclic performance and
charge–discharge efficiency with ethers is far from the desired
level. McCloskey et al. and Bruce et al. and many other studies[19, 56, 57, 58] investigated the stability of the widely used ether
dimethoxyethane (DME; Figure 11). Although the major discharge product in the ORR was Li2O2, side products such as
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
Figure 11. Raman spectra of discharged carbon cathodes from a pure DMEbased cell, a 1:1 (v/v) EC/DME-based cell, and a 1:2 (v/v) PC/DME-based cell.
The spectrum of neat P50 carbon paper is included for comparison. Discharge conditions: 0.09 mA cm2 under O2 to 2 V. Adapted from Ref. [19].
Li2CO3, HCO2Li, and C2H4(OCO2Li)2 formed at the initial cycles
and continued to accumulate. The use of isotope tracer methods also indicated that the side products were mainly from
electrolyte decomposition. In addition to DME, tetraethylene
glycol dimethyl ether (TEGDME) is a promising alternative organic solvent and is now commonly used. Scrosati et al.[42] reported a high-performance lithium–air battery with the capacity of 5000 mA h gcarbon1 at a relatively high current density of
3 A gcarbon1. The high performance is believed to be associated
with the chemical inertia of the optimized TEGDME–LiOTf electrolyte. The lifetime of O2 is relatively short, and it is nearly instantly involved in the reaction that forms Li2O2.
As well as TEGDME, DMSO is another stable solvent according to several recent reports. Stable cathode-based examination shows that the quantity of decomposition product with
TEGDME is double that of the DMSO-based electrolyte. Although both electrolytes undergo little decomposition,
TEGDME showed a higher polarization increase on cycling.[45]
In addition, Mozhzhukhina et al.[59] tested the stability of DMSO
317
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
using infrared spectroscopy and demonstrated that it was
stable during discharge. However, under a high potential, the
DMSO undergoes electrochemical oxidation and the formation
of dimethyl sulfone was observed.[59] The experiment was conducted using the porous gold cathode, which eliminates the
influence of carbon decomposition. The formation of dimethyl
sulfone was detected at a charge potential above approximately 4.3 V. Furthermore, DMSO has a high volatility. The
same challenges occur with the use of diethyl ether. Metal lithium undergoes oxidization in DMSO, which is another limitation for the stability of this solvent. For further information on
electrolyte, the reader is referred to selected books and reviews.[60–62]
ILs, which have been used widely in the field of traditional
batteries, are also considered to be promising electrolytes for
lithium–air batteries. Their high specific heat capacity allows
them to operate under special conditions. The use of ILs improves stability of the battery, and few side reactions and
products are observed.[63–66] However, their high viscosity and
low conductivity limit battery performance. Xu et al.[66] revealed
the suitability of ILs as solvents, and showed that a cell containing N-methyl-N-butylpyrrolidinium bis(trifluoromethanesulfonyl)imide (Pyr14TFSI) can be only discharged at low operating
currents. Further investigation of IL-based lithium–air batteries
is needed.
explore the stability of different common lithium salts in lithium–air batteries, and their effects (Figures 12 and 13). X-ray
diffraction (XRD), NMR, and XPS methods were used to investigate the decomposition products. Cyclic performance was associated with salt decomposition by discharge–charge voltage
profiles and XPS results. This research indicates that salts, and
not only solvents, are also an important influence on the performance of lithium–air batteries. Of the seven lithium salts
tested, LiClO4 is the most stable. Other salts form LiF (LiPF6,
LiOTf, LiBF4, LiTFSI) or Li2C2O4 (LiBOB), as detected by XPS
(Figure 13). LiBr, analogous to LiNO3-based electrolyte, due to
its low conductivity in a glyme-based solvent, showed a somewhat low capacity (0.2 Ah g1). However, this is still a novel
idea, because no adverse reaction was observed using LiNO3 in
N,N-dimethylacetamide.[71] For traditional lithium salts, an interesting phenomenon is that the LiClO4-based battery does not
deliver the highest capacity during the cycle, which indicates
that the degradation of salts is only a small factor in overall
performance. The decomposition of solvent contributes most
to the instability.
4. Ambient Factors
The high energy density of lithium–air batteries is partially the
result of using ambient air. The environment provides a sustainable cathode-active material, O2, allowing higher energy storage compared with other modern batteries. However, ambient
air is a double-edged sword, and also introduces undesired
gases into the battery interior. Among the destructive gas species, CO2 and the humidity in ambient air are the primary contaminations.[72, 73, 76] First, water and CO2 attack the discharge
product Li2O2 to form Li2CO3 on its surface, according to the re2 Li2O2 + 2 CO2 !
actions
2 H2O + 2 Li2O2 !4 LiOH + O2›,
2 Li2CO3 + O2 and 2 LiOH + CO2 !Li2CO3 + H2O. After a few
cycles, Li2CO3, instead of Li2O2, is predominant even without
electrolyte or electrode decomposition. The potential for
charging Li2CO3 to evolve CO2 (> 4 V) is much higher than
charging Li2O2 ( 3.0–3.5 V), which enhances the potential gap
between charge and discharge.[74] Second, the anode is con-
3.2. Lithium Salts
Organic salts such as lithium tetrafluoroborate (LiBF4), lithium
bis(oxalato)borate (LiBOB), lithium trifluoromethanesulfonate
(LiOTf), lithium bis(trifluoromethanesulfonyl)imide (LiTFSI), lithium perchlorate (LiClO4), lithium hexafluorophosphate (LiPF6),
lithium bromide (LiBr) and lithium nitrate (LiNO3) are highly associated with the ionic conductivity of electrolytes, and greatly
influence the formation of the solid–electrolyte interphase. The
decomposition of salts is also a problem impacting the stability. The capability of anti-O2 (and O2
2 ) is also necessary for the
lithium salts. However, the decomposition of lithium salts is
slightly less than of the solvent, which still causes a reduction
in capacity and irreversibility.[67]
LiPF6 is widely used in both Liion and lithium–air batteries,
and undergoes hydrolysis to
produce HF and LiF. In the initial
study, Oswald et al.[67] stated that
LiPF6, LiClO4, and LiBOB all decompose upon contact with reactive Li2O2. Other studies also
confirmed that LiPF6 decomposed to form LiF.[55, 68] LiBOB decomposes to form boron oxide
and lithium carbonates.[69] The
degradation of organic salts is
attributed to both chemical and
electrochemical effects. A sysFigure 12. a) First-cycle discharge–charge voltage profiles for Li–O2 batteries with various electrolytes at
tematic investigation was con- 0.05 mA cm2 current density. b) Cycling stability of Li–O2 batteries with various electrolytes at 0.05 mA cm2
ducted by Nasybulin et al.[70] to current density. Adapted from Ref. [70].
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
318
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
effect of water was investigated
by Meini et al.[75] As shown in
Figure 14, in the LiClO4/DME
electrolyte, with a small amount
of water present in the air, the
capacity was increased largely at
the first discharge. As it is proposed that the capacity of the
lithium–air battery is limited by
the resistivity of Li2O2, the consumption of Li2O2 by water (to
form soluble LiOH or H2O2)
might be the reason for the increase of capacity in the initial
cycles. However, assuming that
the permeated water corrodes
the lithium anode and salts, the
lithium–air battery should be
kept free of water in the long
term. On the other hand, water
and CO2 react to form LiCO3 on
the surface, and the large
charge–discharge potential gap
and the origin of CO2 are shown
in Figure 15.
A high charge potential is destructive to the electrolyte and
electrode. However, most current lithium–air batteries have
been examined under ideal conditions, and many excellent results were obtained by operation
under pure O2. If ambient factors
are not taken into account, the
high
performance
obtained
under pure O2 conditions is
viable for further applications. As
research continues, both undesired polluters and means of
protection should be carefully
considered.[75]
To improve the stability of lithium–air batteries operated under
the ambient conditions, early approaches reduced the electrode
diffusivity to decrease the diffusion of both undesired and
useful species by other porous
membranes.[77] However, ineffiFigure 13. XPS results for a) O 1s scan of the discharge products, b) Li 1s scan of the discharge products, c) F 1s
cient air diffusion in lithium–air
scan of the discharge products and the pure salts, d) S 2p scan of the discharge products and the pure salts,
batteries is a key challenge, as
e) P 2p scan of the discharge products and the pure salts, f) C 1s scan of the discharge products, g) B 1s scan of
the discharge products and the pure salts, and h) Cl 2p scan of the discharge products and the pure salt. The y
the primary diffusivity is relativeaxis stands for the intensity (counts) of photon electrons. Adapted from Ref. [70].
ly low due to the nanostructured
pores.[33, 78] More importantly, although the electrode diffusivity is decreased, a molecule of
sumed by dissolved water, according to the reactions 2 Li +
water is smaller than O2, and water therefore permeates into
2 CO2 + O2 !2 Li2CO3 and 4 Li + 2 H2O!2 Li2OH + H2›. Third,
salts are also hydrolyzed with permeated water. The overall
the electrolyte. An optimized selective membrane was then
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
319
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
Figure 15. Galvanostatic discharge and charge (0.47 mA cm2) of cells discharged under a pure 18O2 headspace (a1–c1) and a 10:90 C18O2/16O2 mixture
(a2–c2). The (a) panels present the discharge profiles. The (b) panels present
the O2 evolution rate (n’O2) during the cell charge. Only single isotopes, that
is, 18O2 (b1) and 16O2 (b2), were evolved from each cell and corresponded to
the isotope used in the discharge headspace. The (c) panels present the isotopic CO2 evolution rate (n’) during the cell charge. Cells were charged
under an Ar headspace. Adapted from Ref. [74].
Figure 14. a) Discharge capacity (1st cycle) comparison of Li–O2 cells with
0.1 m LiClO4 in DME with water-free, or water-contaminated, oxygen
(100 kPa(abs)), and Li–O2 cells with water introduced by means of a small leak
between the cell and ambient air (a) in the “leaker cell”, or by connecting
a water reservoir to produce H2O-saturated O2 inside the cell (g) in the
“water vapor cell”. b) Discharge capacity (1st cycle) comparison of sealed Li–
O2 cells using water-free, or deliberately water-contaminated electrolyte
(0.1 m LiClO4 in DME) with 250 (a), 500 (g), and 1000 (d) ppm of
water. The cells were discharged galvanostatically at 120 mA gcarbon1
(0.05 mA cm2 electrode) after a 30 min rest period at open-circuit voltage
(OCV) in pure 100 kPa(abs) O2. c) Nyquist impedance plots of Li–O2 cells using
water-free (*) and water-contaminated (1000 ppm, *) electrolytes registered after the 30 min OCV rest period (100 kHz to 0.1 Hz at an AC perturbation of 5.0 mV). The water vapor and leaker cells were discussed in Ref. [75].
Adapted from Ref. [75].
in Li-ion batteries, which are closed systems, are relatively
stable. In contrast, lithium–air batteries are semi-open systems
as the cathode exchanges gas species directly with the external environment. The high volatility of electrolytes apparently
reduces the solvent in these systems, and as the cycle number
increases the loss of electrolytes is substantial. Selective membranes or low-diffusivity electrodes can reduce solvent evaporation, but can cause other negative issues. Zhang et al.[80] introduced a blend of tris(2,2,2-trifluoroethyl) phosphate (TFP)/
PC organic solvents as a low-volatility electrolyte. The evaporation rate of liquid blend decreases after the addition of TFP,
making such batteries safe to operate under ambient conditions. However, the ion conductivity is also reduced, leading to
a high overpotential. Interestingly, most current solutions to
stability issues sacrifice the energy utilization efficiency.
proposed by Zhang et al.[77] O2-selective silicone oils were
loaded on porous membranes or polytetrafluoroethylene films;
due to synergetic effects on moisture blocking, O2 must be
carefully selected. Another air-diffusion layer inevitably decreases the diffusivity of electrodes. With increasing cathode
thickness, the maximum current density decreases. As a result,
one more selective layer, due to its harmful obstruction is also
unrealizable unless an advanced and high-performance material is synthesized for this purpose in the future. In addition, N2
can permeate easily into the selective membranes, which demands further investigation.
Furthermore, as presented above, O2 is of low solubility in
the organic electrolytes currently used. Read et al.[32] reported
the critical role of improving the solubility of O2 for increasing
the capacity. The diffusion of O2 depends partly on the attractive force generated from the solubility. The O2 transport rate
can further decrease once the reactive product (Li2O2) deposits
on the porous electrode surface. The pores of electrodes can
also be blocked by organic electrolyte,s as the attractive force
between O2 and the electrolyte is rather small.[79] Early work indicated that organic solvents such as diethyl ether and DMSO
volatilize in the exposed environment.[45] The organic solvents
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
5. Aqueous and Nonaqueous systems
Currently available lithium–air batteries consist of four electrolytic types—aqueous, nonaqueous, hybrid, and solid state
(Figure 16).[8] In general, an organic liquid is used in the nonaqueous electrolyte type and the reaction product is Li2O2.
However, the application of Li + -conducting membranes makes
it possible to replace unstable nonaqueous electrolytes with
aqueous electrolytes. Wang et al.[82] have proposed a new type
of lithium–air rechargeable battery, by applying a special aqueous electrolyte. The all-organic electrolyte was replaced with
320
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
an organic liquid j Li + -conducting membrane j aq. electrolyte
system. The membrane cannot contact the metal lithium directly, and thus an organic layer or IL is typically used to protect the anode. O2 is consumed continuously during discharge
to form LiOH. Aqueous lithium–air batteries are of higher stability compared with their nonaqueous counterparts in many
aspects. This type of lithium–air battery can be operated safely
in the ambient air as the dissolved water and CO2 affects metal
lithium only slightly. Only Li + is capable of being transported
through conducting membranes, and CO2 and water are confined to the cathode side. Due to the solubility of LiOH, electrode passivation is alleviated compared with the nonaqueous
lithium–air battery. Side reactions in the aqueous electrolyte
are much fewer. Compared with the organic liquid, the aqueous electrolyte is rather stable, even though radical ions attack
it. The boiling point of water is higher than most organic solvents, therefore solvent evaporation occurs much less than
with an organic liquid. O2 dissolves in aqueous electrolytes
more easily, which produces a more stable energy output.
However, an aqueous system has its own problems of instability. Separated by the Li + -conducting membranes, the aqueous
electrolyte containing LiCl or LiNO3 is usually applied in the
oxygen half-cell. A Li + -conducting glass or ceramic (e.g.
Li1.35Ti1.75Al0.25Si0.1P0.9O12, LTAP) is typically used as the membrane. The solubility of LiOH is only approximately 13 g in
100 g of room-temperature water, and most of the material adheres to the Li + -conducting membrane and positive electrode
pores, thereby clogging them. Capacity reduces rapidly after
approximately five cycles (in 10 m LiCl with only a protected
cathode).[81] For precipitating LiOH in the electrolyte instead of
at the cathode or ceramic separator, a special design is now
widely used (Figure 17).[81, 91]
Current lithium-conducting membranes are unstable in
strongly acidic or basic solutions.[83–85] After ORR, the product
LiOH increases the alkalinity of the electrolyte. The development of Li+-conducting membranes is considered key to increasing the stability of the aqueous lithium–air battery.
Hasegawa et al.[87] investigated the stability of NASICONtype glass ceramics
(Li1 + z + yAlxTi2xSiyP3yO12), and LATP in aqueous 1 m LiOH and
0.1 m HCl. XRD and scanning electron microscopy results indicated that after three weeks, either LiOH- or HCl-saturated
LTAP membranes were destroyed to different extents. Follow-
Figure 17. The bi-electrode design. The composite air electrode contains an
anionic polymer electrolyte and the ceramic separator has been developed
with a cationic polymer. Adapted from Ref. [91].
Figure 16. Types of lithium–air batteries: a) aqueous electrolyte, b) aprotic/nonaqueous electrolyte, c) mixed
(aprotic-aqueous), and d) solid state. Adapted from Ref. [8].
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
321
ing this work, Shimonishi et al.[83]
investigated the stability of an
NASICON-type glass ceramic
(LTAP) in an aqueous solution of
a weak acid. After immersion in
an acetic acid/lithium acetate solution at 50 8C for three weeks,
no apparent change in conductivity was observed. The test results suggest that LTAP is stable
in both neutral and weakly
acidic solutions, but not in
strongly acidic or alkaline electrolytes. To protect the Li + -conducting membrane, a lithium
acetate/acetic acid solution is
proposed as the cathodic electrolyte. However, the catalyst
choice is limited by the acidic
solution and only acid-resistant
materials such as noble metals
(e.g. Pt, Au) can be used. Utilizing such catalysts notably improves cost effectiveness. Another approach is to utilize a con-
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
developed further.[11, 89] Third, suitable cathode materials and
catalysts have the promise to reduce the overpotential and
keep the system stable.[94–99] Future solutions to these issues
will facilitate the development of the lithium–air battery, and
be source of inspiration across the entire field of lithium
batteries.[100]
centrated lithium salt as the electrolyte, such as 10 m LiCl solution. Due to the common-ion effect, the reversible reaction,
LiOHQLi + + OH , occurs and, therefore, the pH of decreases to
8.14.[92] In addition, the ion conductivity of the lithium-conducting membrane is low, and we can expect advances in the
conducting membrane field.
In addition to the separators, many other issues still exist.
During the charging process, the OER forms LiOH instead of
Li2O2 by the reaction 4 Li + + O2 + 2 H2O + e !4 LiOH. The LiOH
in the electrolyte also reacts with CO2 from the air. Water has
a narrow electrochemical window, leading to dramatic solvent
consumption after several cycles. More importantly, an inextricable problem in all Li-based batteries is the formation of lithium dendrites.[86] Huge capacity degradation is predicted as
a result of the formation of lithium dendrites. The dendrite formation typically occurs in an aqueous lithium–air battery due
to the special Li-separator design. Lithium dendrites enhance
the possibility of attaching the separator to bring about degradation of the fragile membrane. Like the aqueous systems, the
solid-state lithium–air battery structure is also attracting increased attention due to its evident stability.
Kumar et al.[88] presented a solid-state rechargeable lithium–
air battery. The liquid electrolyte was replaced with a waterproof glass ceramic and a polymer ceramic to deliver stable recharge performance. The solid membrane in the lithium air
battery is non-volatile and stable. In nonaqueous batteries, the
electrolyte saturates the air electrode and reacts with the
active radicals O2 or O2
2 in an undesired side reaction. Whereas the solid electrolyte facilitates Li + transport, less decomposition means it is stable even when being operated under ambient air conditions. However, solid electrolytes have a large resistance. Studies also show that battery conductivity increases
only if the operating temperature increases.[88] Based on those
factors, current solid-state lithium–air batteries are typically operated at high temperatures. Interface stability is also a big
issue, as the system is assembled in a solid-solid fashion in the
cathodic cell. In summary, although solid-state lithium–air batteries are stable, the overall performance is limited by the materials. Ceramics or polymers are of great stability but are less
effective; the high internal resistance of such a material is the
critical drawback and remains to be addressed.
7. Summary
The recently uncovered instability issues largely limit the application of the lithium–air battery. In this Minireview, the origin
of these instabilities and the effectiveness of a few recently
proposed solutions have been discussed. Different types of stability have been reviewed at the end of this article. The latest
reports have shown promising prospects for the application of
the lithium–air battery for solving the major issues that include
electrolyte decomposition, ambient operation, and anode
protection.
Keywords: electrochemistry · electrode
instability factors · lithium–air batteries
For further advances in the field of lithium–air batteries, stability issues must be addressed. First, stable electrodes and electrolyte materials are in urgent demand. Advanced electrolytes
and electrodes are expected to undergo less decomposition
during cycles.[90, 93] The stability of electrodes of allotropic
carbon nanoforms, such as graphene, deserves a more detailed
investigation. Second, a more systematic examination of ambient operation should be conducted in the future and developed. The O2-selective membranes represent significant progress towards this, and in this respect, high-selectivity membranes are worth investigating. On the inside of the system,
protective measures, including the use of LTAP membranes in
the aprotic lithium–air battery, also have the possibility to be
www.chemelectrochem.org
electrolyte
·
[1] M. Liu, M. B. Johnston, H. J. Snaith, Nature 2013, 501, 395 – 398.
[2] T. Abbasi, S. A. Abbasi, Renewable Sustainable Energy Rev. 2010, 14,
919 – 937.
[3] S. P. S. Badwal, S. Giddey, A. Kulkarni, C. Munnings, J. Aust. Ceram. Soc.
2014, 50, 23 – 37.
[4] R. BaÇos, F. Manzano-Agugliaro, F. G. Montoya, C. Gil, A. Alcayde, J.
Gmez, Renewable Sustainable Energy Rev. 2011, 15, 1753 – 1766.
[5] S. Han, D. Wu, S. Li, F. Zhang, X. Feng, Adv. Mater. 2014, 26, 849 – 864.
[6] P. Poizot, F. Dolhem, Energy Environ. Sci. 2011, 4, 2003 – 2019.
[7] G. Zhou, D. W. Wang, F. Li, L. Zhang, N. Li, Z. S. Wu, W. Lei, G. Q. Lu,
H. M. Cheng, Chem. Mater. 2010, 22, 5306 – 5313.
[8] M. A. Rahman, X. Wang, C. Wen, J. Appl. Electrochem. 2014, 44, 5 – 22.
[9] G. Girishkumar, B. McCloskey, A. C. Luntz, S. Swanson, W. Wilcke, J.
Phys. Chem. Lett. 2010, 1, 2193 – 2203.
[10] K. M. Abraham, Z. Jiang, J. Electrochem. Soc. 1996, 143, 1 – 5.
[11] S. J. Visco, Y. S. Nimon, B. Katz, US Patent 8652692, 2014.
[12] S. J. Visco, B. D. Katz, Y. S. Nimon, L. C. De Jonghe, US Patent 8828580,
2014..
[13] A. Kraytsberg, Y. Ein-Eli, J. Power Sources 2011, 196, 886 – 893.
[14] V. S. Bryantsev, V. Giordani, W. Walker, M. Blanco, S. Zecevic, K. Sasaki,
G. V. Chase, J. Phys. Chem. A 2011, 115, 12399 – 12409.
[15] V. S. Bryantsev, J. Uddin, V. Giordani, W. Walker, D. Addison, G. V. Chase,
J. Electrochem. Soc. 2013, 160, A160 – A171.
[16] V. S. Bryantsev, F. Faglioni, J. Phys. Chem. A 2012, 116, 7128 – 7138.
[17] D. Capsoni, M. Bini, S. Ferrari, E. Quartarone, P. Mustarelli, J. Power
Sources 2012, 220, 253 – 263.
[18] I. Kowalczk, J. Read, M. Salomon, Pure. Appl. Chem. 2007, 79, 851 – 860.
[19] B. D. McCloskey, D. S. Bethune, R. M. Shelby, G. Girishkumar, A. C.
Luntz, J. Phys. Chem. Lett. 2011, 2, 1161 – 1166.
[20] R. Padbury, X. Zhang, J. Power Sources 2011, 196, 4436 – 4444.
[21] M. M. Ottakam Thotiyl, S. A. Freunberger, Z. Peng, P. G. Bruce, J. Am.
Chem. Soc. 2013, 135, 494 – 500.
[22] J. Xiao, D, Mei, X, Li, W. Xu, D. Wang, G. L. Graff, J. G. Zhang, Nano
Lett. 2011, 11, 5071 – 5078.
[23] J. Kang, O. L. Li, N. Saito, J. Power Sources 2014, 261, 156 – 161.
[24] Y. Li, J. Wang, X. Li, J. Liu, D. Geng, J. Yang, X. Sun, Electrochem.
Commun. 2011, 13, 668 – 672.
[25] N. Zhao, C. Li, X. Guo, Energy Technol. 2014, 2, 317 – 324.
[26] S. H. Oh, R. Black, E. Pomerantseva, J. H. Lee, L. F. Nazar, Nat. Chem.
2012, 4, 1004 – 1010.
[27] J. Xiao, D. Wang, W. Xu, D. Wang, G. L. Graff, J. G. Zhang, J. Electrochem. Soc. 2010, 157, A487 – A492.
[28] C. Tran, X. Q. Yang, D. Qu, J. Power Sources 2010, 195, 2057 – 2063.
6. Prospects
ChemElectroChem 2015, 2, 312 – 323
·
322
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Minireviews
[29] R. Younesi, M. Hahlin, K. Edstrçm, ACS Appl. Mater. Interfaces 2013, 5,
1333 – 1341.
[30] W. Xu, V. V. Viswanathan, D. Wang, S. A. Towne, J. Xiao, Z. Nie, J. G.
Zhang, J. Power Sources 2011, 196, 3894 – 3899.
[31] Y. Wang, Electrochim. Acta 2012, 75, 239 – 246.
[32] J. Read, K. Mutolo, M. Ervin, W. Behl, J. Wolfenstine, A. Driedger, D.
Foster, J. Electrochem. Soc. 2003, 150, A1351 – A1356.
[33] N. Ding, S. Chien, T. S. A. Hor, R. Lum, Y. Zong, Z. Liu, J. Mater. Chem. A
2014, 2, 12433 – 12441.
[34] C. Tran, J. Kafle, X. Q. Yang, D. Qu, Carbon 2011, 49, 1266 – 1271.
[35] V. Viswanathan, K. S. Thygesen, J. S. Hummelshøj, J. K. Nørskov, G. Girishkumar, B. D. McCloskey, A. C. Luntz, J. Chem. Phys. 2011, 135,
214704.
[36] P. Albertus, G. Girishkumar, B. McCloskey, R. S. Sanchez-Carrera, B. Kozinsky, J. Christensen, A. C. Luntz, J. Electrochem. Soc. 2011, 158, A343 –
A351.
[37] J. Chen, J. S. Hummelshøj, K. S. Thygesen, J. S. G. Myrdal, J. K. Nørskov,
T. Vegge, Catal. Today 2011, 165, 2 – 9.
[38] M. D. Radin, J. F. Rodriguez, F. Tian, D. J. Siegel, J. Am. Chem. Soc. 2011,
134, 1093 – 1103.
[39] J. S. Hummelshøj, J. Blomqvist, S. Datta, T. Vegge, J. Rossmeisl, K. S.
Thygesen, A. C. Luntz, K. W. Jacobsen, J. K. Nørskov, J. Chem. Phys.
2010, 132, 071101.
[40] R. Mi, H. Liu, H. Wang, K. W. Wong, J. Mei, Y. Chen, H. Yan, Carbon
2014, 67, 744 – 752.
[41] B. Sun, H. Liu, P. Munroe, H. Ahn, G. Wang, Nano Res. 2012, 5, 460 –
469.
[42] H. G. Jung, J. Hassoun, J. B. Park, Y. K. Sun, B. Scrosati, Nat. Chem.
2012, 4, 579 – 585.
[43] J. Read, J. Electrochem. Soc. 2002, 149, A1190 – A1195.
[44] B. D. McCloskey, A. Speidel, R. Scheffler, D. C. Miller, V. Viswanathan,
J. S. Hummelshøj, A. C. Luntz, J. Phys. Chem. Lett. 2012, 3, 997 – 1001.
[45] Z. Peng, S. A. Freunberger, Y. Chen, P. G. Bruce, Science 2012, 337,
563 – 566.
[46] M. M. Ottakam Thotiyl, S. A. Freunberger, Z. Peng, Y. Chen, Y. Liu, P. G.
Bruce, Nat. Mater. 2013, 12, 1050 – 1056.
[47] S. A. Freunberger, Y. Chen, F. Bard, K. Takechi, F. Mizuno, P. G. Bruce in
The Lithium Air Battery (Eds.: N. Imanishi, A.C. Luntz, P. Bruce), Springer,
New York, 2014, pp. 23 – 58.
[48] J. Hassoun, F. Croce, M. Armand, B. Scrosati, Angew. Chem. Int. Ed.
2011, 50, 2999 – 3002; Angew. Chem. 2011, 123, 3055 – 3058.
[49] D. Sharon, V. Etacheri, A. Garsuch, M. Afri, A. A. Frimer, D. Aurbach, J.
Phys. Chem. Lett. 2013, 4, 127 – 131.
[50] E. J. Calvo, N. Mozhzhukhina, Electrochem. Commun. 2013, 31, 56 – 58.
[51] S. A. Freunberger, Y. Chen, Z. Peng, J. M. Griffin, L. J. Hardwick, F. Bard,
P. G. Bruce, J. Am. Chem. Soc. 2011, 133, 8040 – 8047.
[52] B. D. McCloskey, R. Scheffler, A. Speidel, G. Girishkumar, A. C. Luntz, J.
Phys. Chem. C 2012, 116, 23897 – 23905.
[53] Y. C. Lu, Y. Shao-Horn, J. Phys. Chem. Lett. 2013, 4, 93 – 99.
[54] H. Wang, K. Xie, Electrochim. Acta 2012, 64, 29 – 34.
[55] Z. Zhang, J. Lu, R. S. Assary, P. Du, H. H. Wang, Y. K. Sun, K. Amine, J.
Phys. Chem. C 2011, 115, 25535 – 25542.
[56] S. A. Freunberger, Y. Chen, N. E. Drewett, L. J. Hardwick, F. Bard, P. G.
Bruce, Angew. Chem. Int. Ed. 2011, 50, 8609 – 8613; Angew. Chem. 2011,
123, 8768 – 8772.
[57] S. Meini, N. Tsiouvaras, J. Herranz, M. Piana, H. Gasteiger, ECS Meeting
Abstracts 2013, MA2013-02, 429.
[58] D. G. Kwabi, N. Ortiz-Vitoriano, S. A. Freunberger, Y. Chen, N. Imanishi,
P. G. Bruce, Y. Shao-Horn, MRS Bull. 2014, 39, 443 – 452.
[59] N. Mozhzhukhina, L. P. Mndez De Leo, E. J. Calvo, J. Phys. Chem. C
2013, 117, 18375 – 18380.
[60] A. M. Stephan, Eur. Polym. J 2006, 42, 21 – 42.
[61] Z. L. Wang, D. Xu, J. J. Xu, X. B. Zhang, Chem. Soc. Rev. 2014, 43, 7746 –
7786.
[62] T. R. Jow, K. Xu, O. Borodin, M. Ue, Electrolytes for Lithium and LithiumIon Batteries, Springer, New York, 2014.
[63] J. Herranz, A. Garsuch, H. A. Gasteiger, J. Phys. Chem. C 2012, 116,
19084 – 19094.
ChemElectroChem 2015, 2, 312 – 323
www.chemelectrochem.org
[64] K. Takechi, S. Higashi, F. Mizuno, H. Nishikoori, H. Iba, T. Shiga, ECS Electrochem. Lett. 2012, 1, A27 – A29.
[65] G. A. Elia, J. B. Park, Y. K. Sun, B. Scrosati, J. Hassoun, ChemElectroChem.
2014, 1, 47 – 50.
[66] W. Xu, J. Hu, M. H. Engelhard, S. A. Towne, J. S. Hardy, J. Xiao, J. G.
Zhang, J. Power Sources 2012, 215, 240 – 247.
[67] S. Oswald, D. Mikhailova, F. Scheiba, P. Reichel, A. Fiedler, H. Ehrenberg,
Anal. Bioanal. Chem. 2011, 400, 691 – 696.
[68] G. M. Veith, N. J. Dudney, J. Howe, J. Nanda, J. Phys. Chem. C 2011, 115,
14325 – 14333.
[69] S. H. Oh, T. Yim, E. Pomerantseva, L. F. Nazar, Electrochem. Solid-State
Lett. 2011, 14, A185 – A188.
[70] E. Nasybulin, W. Xu, M. H. Engelhard, Z. Nie, S. D. Burton, L. Cosimbescu, J. G. Zhang, J. Phys. Chem. C 2013, 117, 2635 – 2645.
[71] W. Walker, V. Giordani, J. Uddin, V. S. Bryantsev, G. V. Chase, D. Addison,
J. Am. Chem. Soc. 2013, 135, 2076 – 2079.
[72] J. G. Zhang, D. Wang, W. Xu, J. Xiao, R. E. Williford, J. Power Sources
2010, 195, 4332 – 4337.
[73] M. Mirzaeian, P. J. Hall, F. B. Sillars, I. Fletcher, M. M. Goldin, G. O. Shittabey, H. F. Jirandehi, J. Electrochem. Soc. 2013, 160, A25 – A30.
[74] S. R. Gowda, A. Brunet, G. M. Wallraff, B. D. McCloskey, J. Phys. Chem.
Lett. 2013, 4, 276 – 279.
[75] S. Meini, M. Piana, N. Tsiouvaras, A. Garsuch, H. A. Gasteiger, Electrochem. Solid-State Lett. 2012, 15, A45 – A48.
[76] M. Mirzaeian, P. J. Hall, M. M. Goldin, ECS Meeting Abstracts 2014,
MA2014-1, 205.
[77] J. Zhang, W. Xu, X. Li, W. Liu, J. Electrochem. Soc. 2010, 157, A940 –
A946.
[78] J. Zhang, W. Xu, W. Liu, J. Power Sources 2010, 195, 7438 – 7444.
[79] K. H. Xue, T. K. Nguyen, A. A. Franco, J. Electrochem. Soc. 2014, 161,
E3028 – E3035.
[80] S. S. Zhang, K. Xu, J. Read, J. Power Sources 2011, 196, 3906 – 3910.
[81] S. Sunahiro, M. Matsui, Y. Takeda, O. Yamamoto, N. Imanishi, J. Power
Sources 2014, 262, 338 – 343.
[82] Y. Wang, H. Zhou, J. Power Sources 2010, 195, 358 – 361.
[83] Y. Shimonishi, T. Zhang, P. Johnson, N. Imanishi, A. Hirano, Y. Takeda, N.
Sammes, J. Power Sources 2010, 195, 6187 – 6191.
[84] Y. Shimonishi, A. Toda, T. Zhang, A. Hirano, N. Imanishi, O. Yamamoto,
Y. Takeda, Solid State Ionics 2011, 183, 48 – 53.
[85] P. He, Y. Wang, H. Zhou, J. Power Sources 2011, 196, 5611 – 5616.
[86] R. Khurana, J. Schaefer, L. A. Archer, G. W. Coates, J. Am. Chem. Soc.
2014, 136, 7395 – 7402.
[87] S. Hasegawa, N. Imanishi, T. Zhang, J. Xie, A. Hirano, Y. Takeda, O. Yamamoto, J. Power Sources 2009, 189, 371 – 377.
[88] B. Kumar, J. Kumar, R. Leese, J. P. Fellner, S. J. Rodrigues, K. M. Abraham, J. Electrochem. Soc. 2010, 157, A50 – A54.
[89] T. Zhang, H. Zhou, Nat. Commun. 2013, 4, 1817.
[90] Z. K. Luo, C. S. Liang, F. Wang, Adv. Funct. Mater. 2014, 24, 2101 – 2105.
[91] P. Stevens, G. Toussaint, L. Puech, P. Vinatier, ECS Trans. 2013, 50, 1 – 11.
[92] Y. Shimonishi, T. Zhang, N. Imanishi, D. Imb, D. Lee, A. Hiranoa, Y.
Takeda, O. Yamamoto, N. Sammesc, J. Power Sources 2011, 196, 5128 –
5132.
[93] W. He, X. Wang, L. Ye, Y. Pan, Y. Song, A. Liu, J. H. Dickerson, ChemElectroChem. 2014, 1, 2052 – 2057.
[94] W. He, B. Wang, Adv. Energy Mater. 2012, 2, 329 – 333.
[95] W. He, J. Zou, B. Wang, S. Vilayurganapathy, M. Zhou, X. Lin, K. H. L.
Zhang, J. Lin, P. Xu, J. H. Dickerson, J. Power Sources 2013, 237, 64 – 73.
[96] W. He, Bin Wang, J. Power Sources 2013, 232, 93 – 98.
[97] W. He, B. Wang, J. H. Dickerson, Nano Energy 2012, 1, 828 – 832.
[98] W. He, B. Wang, H. Zhao, Y. Jiao, J. Power Sources 2011, 196, 9985.
[99] W. He, W. Lv, J. H. Dickerson, Gas Transport in Solid Oxide Fuel Cells
Springer 2014.
[100] X. Wang, K. Wen, Y. Song, J. Power Sources 2014, in press.
Received: September 17, 2014
Published online on December 16, 2014
323
2015 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim