DOI: 10.1126/science.1223985
, 563 (2012);337 Science
et al.Zhangquan Peng
Battery2A Reversible and Higher-Rate Li-O
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A Reversible and Higher-Rate
Li-O2 Battery
Zhangquan Peng, Stefan A. Freunberger,* Yuhui Chen, Peter G. Bruce†
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.
Atypical rechargeable nonaqueous Li-O2cell is composed of a Li metal anode(negative electrode), a nonaqueous Li+
conducting electrolyte, and a porous cathode
(positive electrode) (1–6). Operation of the cell
depends critically on O2 being reduced at the
cathode to O2
2–, which combines with Li+ from
the electrolyte to form Li2O2 on discharge, and
the reverse reaction occurring during charging
(1–6). Early investigation of nonaqueous Li-O2
cells focused on the use of organic carbonate–
based electrolytes, which have since been shown
to decompose irreversibly at the cathode on dis-
charge to form products such as lithium formate
(HCO2Li), lithium acetate (CH3CO2Li), lithium
propyl-dicarbonate [C3H6(CO2Li)2], and lithium
carbonate (Li2CO3) with little or no evidence
of Li2O2 formation (7–11). Later work turned
to ethers—while initially promising and certainly
more stable to reduced O2 species than organic
carbonates, ethers exhibit increasing electro-
lyte decomposition upon cycling (figs. S1 to S3)
(11–13). These data show that whether combined
with carbon or nanoporous gold (NPG) elec-
trodes, ethers, including dimethoxyethane (DME),
are increasingly unstable upon cycling. For exam-
ple, in the case of DME-based electrolytes after
only 10 cycles, 20% of the discharge products
arise from electrolyte decomposition (fig. S2).
Such side reactions can be difficult to detect by
x-ray diffraction because of poor crystallinity
of the decomposition products. Similar decom-
position of tetraethylene glycol dimethyl ether
(tetraglyme)–based electrolytes has been re-
ported (12) and is also shown to occur at a NPG
electrode (fig. S3). These results demonstrate
that ethers do not support the necessary re-
versible Li2O2 formation/decomposition upon
cycling that is essential for operation of the Li-O2
cell. A very recent paper comes to a different
conclusion from the papers cited above and from
our own results concerning the cyclability of the
tetraglyme/carbon interface (14).
We constructed a Li-O2 cell that contained
an electrolyte composed of 0.1 M LiClO4 in di-
methyl sulfoxide (DMSO) and a NPG cathode
[for details, see the supplementary materials and
methods section (15)]. The cell was operated in
1 atm of O2. Oxygen reduction electrochemistry
at the DMSO/planar-carbon interface has been
studied previously (16). Discharge/charge curves
for the cell on cycles 1, 5, 10, and 100 are shown
in Fig. 1. Most of the initial capacity (95%) is re-
tained after 100 cycles. However, as is now recog-
nized from the work of many authors, the ability
to recharge a Li-O2 cell is not proof that the reac-
tions occurring at the positive electrode are rever-
sible and involve Li2O2 formation/decomposition
(7–13). To demonstrate that the reaction at the
porous cathode is Li2O2 formation/decomposition,
we collected Fourier transform infrared (FTIR)
spectroscopy data at the end of discharge and
charge as a function of cycle number (1, 5, 10,
and 100) (Fig. 2A). At the end of each discharge,
we observedLi2O2. Its formationwas corroborated
by in situ surface-enhanced Raman spectroscopy
(SERS) carried out on a cell with a sapphire win-
dow for transmission of the Raman laser beam
(Fig. 2B) (17). A few small peaks, in addition to
the peaks arising from Li2O2, are apparent in the
School of Chemistry, University of St. Andrews, North Haugh,
St. Andrews, Fife KY16 9ST, UK.
*Present address: Institute for Chemistry and Technology of
Materials, Graz University of Technology, Stremayrgasse 9,
8010 Graz, Austria.
†To whom correspondence should be addressed. E-mail:
p.g.bruce@st-andrews.ac.uk
Fig. 1. Charge/discharge curves (left) and cycling profile (right) for a Li-O2 cell with a 0.1M LiClO4-DMSO
electrolyte and a NPG cathode, at a current density of 500 mAg−1 (based on the mass of Au). Because the
capacities are given per gram of Au, which is ~10-fold heavier (more dense) than carbon, 300 mAhg−1
(based on the mass of Au) would, for the same porous electrode but formed from carbon, correspond to
~3000 mAhg−1 (based on the mass of carbon). FTIR spectra collected upon charging at points A and B are
shown in fig. S7.
Fig. 2. Vibrational spectra of a NPG cathode at the end of discharge and charge in 0.1 M LiClO4-DMSO.
(A) FTIR and (B) SERS spectra.
www.sciencemag.org SCIENCE VOL 337 3 AUGUST 2012 563
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FTIR spectra at the end of discharge, at ~880,
1420, 1490, and 1600 cm−1. These peaks could
be assigned to a mixture of Li2CO3 and HCO2Li,
with no other species being detected, such as from
S containing decomposition products (Fig. 2A).
The presence of HCO2Li was confirmed by wash-
ing the NPG electrode at the end of discharge
with D2O and examining the resulting solution
by 1H nuclear magnetic resonance (NMR), fol-
lowing the procedure described previously (7, 12).
HCO2D in the
1H NMR indicated the presence
of HCO2Li in the discharged electrode before
washing.
Batteries and chemical/electrochemical re-
actions in general exhibit some degree of side
reaction, particularly on the first cycle (e.g., Li-ion
batteries). The key question is the extent of such
side reactions: whether this is sufficiently small
compared with the amount of electrolyte used in
practical cells and whether the extent increases
with cycling. We prepared mechanical mixtures
of Li2O2 with Li2CO3 and Li2O2with HCO2Li of
varying ratios, collected their FTIR spectra, and
constructed calibration curves (figs. S4 and S5);
from these curves, we determined the fractions
of Li2CO3 and HCO2Li in the FTIR spectra in
Fig. 2 to be <1%. The proportion of Li2O2 at the
end of discharge exceeds 99%, and there is no
evidence of this proportion decreasing on cy-
cling.We used 1H and 13CNMR to investigate the
presence of any solution-soluble decomposition
products. Sensitivity to detection of such species
depends on the ratio between the amount of
electrolyte and the amount of discharge product
(15). We collected spectra after 100 cycles to con-
centrate any decomposition products, but we did
not detect evidence of any such species (fig. S6).
We used differential electrochemical mass spec-
trometry (DEMS) to obtain further confirmation
that discharge was overwhelmingly dominated
by Li2O2 formation. TheDEMSprocess involves
in situ mass spectrometric analysis of the gases
consumed/evolved during a slow-sweep (0.1
mVs–1) linear potential scan (Fig. 3A) (15). The
only gas detected on discharge was O2. There was
no evidence of CO2, SO2, or SO3 (i.e., no evidence
of electrolyte decomposition), in contrast to other
electrolytes. The high purity of Li2O2 formation
implies that for every two electrons (e–) passed,
one O2 molecule should be consumed; that is, the
charge-to-mass ratio should be 2e–/O2. The O2
consumption on discharge follows the cell
current (Fig. 3A), and the charge-to-mass ratio is
2e–/O2 on each discharge (Table 1).
The FTIR spectra collected at the end of
charge on cycles 1, 5, 10, and 100 are shown in
Fig. 2A, from which it is clear that the product
formed on discharge has been removed upon
charging. This observation was confirmed by the
SERS data in Fig. 2B, where the characteristic
peak for Li2O2 at ~800 cm
−1, observed at the
end of discharge, is absent from the spectrum at
the end of charge. To probe the oxidation in
more detail, we used DEMS on charging for
cycles 1, 5, 10, and 100 (Fig. 3B). Only O2 was
detected, confirming that Li2O2 had formed on
the previous discharge and also that the electro-
lyte, even in the presence of Li2O2, is stable on
oxidation. Upon examining the linear voltam-
metry (current-voltage curve) in Fig. 3B, several
peaks are evident, corresponding well with the
peaks for O2 evolution. A similar heterogeneous
oxidation process spanning a range of potentials
has been observed previously in porous electrodes
and has been ascribed to oxidation of Li2O2 being
easier in certain pores than in others (11). We col-
lected FTIR spectra (fig. S7) during charging,
at the points shown in Fig. 1. The spectra in-
dicate that the quantity of Li2O2 is diminishing
with increasing state of charge, but that some
Li2O2 is still present at point B. The ratio of charge
passed to O2 evolved on charging is given in
Table 1. As was the case for discharge, the ratio
is close to 2e–/O2 on each cycle, in accord with
charging involving oxidation of Li2O2 without
electrolyte degradation. Over the collection of
up to 100 cycles, the results from FTIR, SERS,
NMR, and DEMS all demonstrate that the cell
cycles by the reversible formation/decomposition
of Li2O2.
To investigate whether the dominance of
Li2O2 formation/decomposition is due to the salt,
solvent, or electrode substrate, we constructed cells
in which LiClO4 was replaced by LiTFSI [lithium
bis(trifluoromethanesulfonyl)imide] and separately
in which the NPG electrode was replaced by car-
bon black (Super P, Timcal, Bodio, Switzerland).
In the former case, the load curves and FTIR spec-
tra at the end of discharge and charge on cycling
are the same as those for LiClO4 (fig. S8),
demonstrating that changing the salt does not in-
fluence the results. In contrast, replacing the NPG
electrode with carbon does adversely affect the
results (Fig. 4). The FTIR at the end of discharge
on carbon shows a greater proportion of side re-
action, Li2CO3, and HCO2Li (Fig. 4). Using cal-
ibration plots, as before, we estimate the total
Fig. 3. DEMS of a NPG cathode during (A) discharge and (B) charge in 0.1 M LiClO4-DMSO. Linear potential
scans at 0.1 mVs−1 (corresponding to a low rate of discharge/charge) between 2.3 and 4.0 V were used.
n’ indicates the gas-consumption/-generation rates during discharge and charge.
Table 1. Ratios of the number of electrons to
oxygen molecules upon reduction (discharge) and
oxidation (charge).
Cycle number
Discharge
e–/O2
Charge
e–/O2
1 2.01 1.98
5 1.99 2.04
10 2.02 1.98
100 2.03 2.01
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proportion of side-reaction products to be ~15%.
The carbon itself may be unstable, as suggested
recently (18), although the HCO2Li formation
is likely to involve DMSO. Further work is re-
quired to investigate the origin of the side pro-
ducts formed at the DMSO/carbon interface. The
charging curve (Fig. 4) is also different from the
NPG electrode (Fig. 1). The voltage rises rapidly,
passes through a very small step at 3.3 V to
~3.75 V, then slowly to 4 V. The higher charging
voltage for carbon versus NPG occurs despite
the current density (based on the true surface
area of the electrode) being less for the carbon
electrode than for NPG: 0.1 mAcm−2 (true surface
area of carbon) compared with 1 mAcm−2 (true
surface area of NPG). Note that the kinetics of
the different electrodes is discussed below. The
DEMS data in Fig. 4 confirm a very minor de-
gree of O2 evolution at 3.3 to 3.4 V, with most
of the O2 being evolved above 4 V and a sub-
stantial amount above 4.5 V, where it is accom-
panied by CO2 evolution, which is indicative of
electrolyte oxidation. The DEMS data for the
Super P carbon cathode in Fig. 4 contrast strong-
ly with those for the NPG electrode in Fig. 3B,
where O2 evolution commences at ~3.2 Vand all
of the O2 is evolved below 4 V (Table 1 confirms
that all of the O2 expected from the Li2O2 present
is evolved). These results indicate that NPG low-
ers the charging voltage (i.e., NPG is more effec-
tive than carbon at promoting Li2O2 oxidation).
The DEMS results for the Super P cathode
are in accord with the difficulty in cycling a cell
with a carbon electrode. Incorporation of a-MnO2
nanowires into a porous carbon electrode proved
effective in promoting Li2O2 oxidation in previ-
ous studies (19). However, reduction of O2 in
the DMSO electrolyte at a Super P electrode in-
corporating a-MnO2 nanowires resulted in the
formation of LiOH on the first discharge, as
noted in previous studies in ethers, possibly aris-
ing from –OH groups on the surface of the
oxide (12). Therefore, we constructed a com-
posite electrode made of Super P with nano-
particulate gold (15). The results are shown in
Fig. 5. As for Super P alone, the side products
are Li2CO3 and HCO2Li, which together ac-
count for ~15% of the discharge products. Charg-
ing occurs at a somewhat lower voltage than
without the Au, as noted previously (20), but
overall nanoparticulate Au/carbon composite
electrodes are less effective at promoting Li2O2
oxidation than NPG electrodes. This is espe-
cially evident when comparing the DEMS data
in Figs. 3, 4, and 5: Whereas only a small pro-
portion of O2 is evolved at the carbon electrode
below 4 V (Fig. 4), the proportion increases
somewhat with the addition of nanoparticulate
Au to the electrode (Fig. 5), but it is much greater
for NPG (Fig. 3).
An important challenge for Li-O2 cells is to
increase the kinetics of the electrode reaction,
which is generally observed to be relatively low,
especially for the charging process (1–6, 21–33).
The rate used in Fig. 1 is 500 mAg−1 of gold
(equivalent to ~5000 mAg−1 for a carbon elec-
trode of the same volume), which translates into
1.0 mAcm−2 based on the total active surface area
of the NPG electrode (50 m2g−1) (15). The rate
used for the carbon-based electrodes (Figs. 4 and
5) is 70 mAg−1, a typical value from the literature
(19, 24), which translates into a true current
density of 0.1 mAcm−2, based on a surface area
for Super P of ~60 m2g−1. Therefore, the true
rate at the electrode surface is 10 times greater
in the case of NPG than is typical for carbon
electrodes. Yet, this is still a relatively low rate
overall. The discharge potential is hardly af-
fected by the change in rate, but as noted above, a
substantial proportion of the charging occurs at
lower voltages for NPG than for carbon or Super
P/nanoparticulate Au, despite the rate being 10-
fold higher for NPG. This result underlines the
fact that oxidation of Li2O2 on NPG is much
more facile than on carbon. Other factors, such
as electrode porosity, can also affect rate per-
formance, and this will differ between NPG and
Super P. Recent studies of the electrocatalysis of
O2 evolution on charging Li2O2 suggest that there
is little evidence of true electrocatalysis (24). We
do not claim electrocatalysis is necessarily
taking place here, but we simply observe that
the charging voltage is lower and kinetics is faster
compared with a carbon electrode. Although the
capacity obtained with NPG in Fig. 1 may look
relatively modest at ~300 mAhg−1, it must be
noted that this value is normalized to the mass of
gold and is equivalent to 3000 mAhg−1 of carbon.
In conclusion, we have shown that a Li-O2
cell composed of a DMSO-based electrolyte and
a NPG electrode can sustain reversible cycling,
retaining 95% of its capacity after 100 cycles
and having >99% purity of Li2O2 formation at
the cathode, even on the 100th cycle, and its
complete oxidation on charge. The charge-to-
Fig. 4. (A) Discharge-charge curve of a Li-O2 cell employing a composite
carbon cathode at 70 mAg−1 (normalized to the mass of carbon). (B) FTIR
at the end of discharge. (C) DEMS of the porous carbon cathode during
charging in 0.1 M LiClO4-DMSO; scan rate 0.1 mVs
–1. The composition of
the cathode is Super P carbon:polytetrafluoroethylene (PTFE) 8:2 (m/m).
n’ indicates the gas-generation rates during the charging process.
Fig. 5. (A) Discharge-charge curve of a Li-O2 cell employing a gold-loaded
composite carbon cathode at 70 mAg−1 (normalized to the mass of car-
bon). (B) FTIR at the end of discharge. (C) DEMS of gold-loaded porous
carbon cathode [Super P:PTFE:Au 8:1:1 (m/m)] during charging in 0.1 M
LiClO4-DMSO; scan rate 0.1 mVs
–1. n’ indicates the gas-generation rates
during the charging process. Note the electrode area is ¼ of that in Fig. 4.
www.sciencemag.org SCIENCE VOL 337 3 AUGUST 2012 565
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mass ratio on discharge and charge is 2e–/O2,
confirming that the reaction is overwhelmingly
Li2O2 formation/decomposition. We have also
shown that such electrodes are particularly effec-
tive at promoting the decomposition of Li2O2,
with all the Li2O2 being decomposed below 4 V
and ~50% decomposed below 3.3 V, at a rate ap-
proximately one order of magnitude higher than
on carbon. Although DMSO is not stable with
bare Li anodes, it could be used with protected Li
anodes. Nanoporous Au electrodes are not suit-
able for practical cells, but if the same benefits
could be obtained with Au-coated carbon, then
low-mass electrodes would be obtained, although
cost may still be a problem. A cathode reaction
overwhelmingly dominated by Li2O2 formation
on discharge, its complete oxidation on charge
and sustainable on cycling, is an essential pre-
requisite for a rechargeable nonaqueous Li-O2
battery. Hence, the results presented here encour-
age further study of the rechargeable nonaqueous
Li-O2 cell, although many challenges to practical
devices remain.
References and Notes
1. K. M. Abraham, Z. Jiang, J. Electrochem.
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