Download A rechargeable room-temperature sodium superoxide

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PUBLISHED ONLINE: 2 DECEMBER 2012 | DOI: 10.1038/NMAT3486
A rechargeable room-temperature sodium
superoxide (NaO2) battery
Pascal Hartmann1 , Conrad L. Bender1 , Miloš Vračar1† , Anna Katharina Dürr2 , Arnd Garsuch2 ,
Jürgen Janek1 * and Philipp Adelhelm1 *
In the search for room-temperature batteries with high energy
densities, rechargeable metal–air (more precisely metal–
oxygen) batteries are considered as particularly attractive
owing to the simplicity of the underlying cell reaction at first
glance1 . Atmospheric oxygen is used to form oxides during discharging, which—ideally—decompose reversibly during charging. Much work has been focused on aprotic Li–O2 cells (mostly
with carbonate-based electrolytes and Li2 O2 as a potential
discharge product), where large overpotentials are observed
and a complex cell chemistry is found2 . In fact, recent studies
evidence that Li–O2 cells suffer from irreversible electrolyte
decomposition during cycling3 . Here we report on a Na–O2
cell reversibly discharging/charging at very low overpotentials
(<200 mV) and current densities as high as 0.2 mA cm−2 using
a pure carbon cathode without an added catalyst. Crystalline
sodium superoxide (NaO2 ) forms in a one-electron transfer
step as a solid discharge product. This work demonstrates that
substitution of lithium by sodium may offer an unexpected
route towards rechargeable metal–air batteries.
Initiated in 1996, research in the field of lithium–air batteries
is particularly motivated by the high energy density of this battery
system4 . The working principle is sketched in Fig. 1. In contrast to
conventional battery systems that are self-containing, a metal–air
battery comprises—like a special type of fuel cell—a gas electrode
that would allow atmospheric oxygen to be used in the cathode
reaction. Assuming Li2 O2 as a solid discharge product, the cell
system operates at up to 2.96 V and exhibits a theoretical energy
density of 3,458 Wh kg−1 , that is, several times higher than the
theoretical energy density of present lithium-ion cells (Supplementary Table S1). In practice, large overpotentials η have been found
on discharge (ηdis ≈ 300 mV) and charge (typically ηchg > 1 V),
leading to relatively low round-trip energy storage efficiencies
down to 60%. Therefore, research has mainly focused on finding
suitable catalysts5–7 . Even though some early results indicated a
more complex cell reaction8–10 , it was only recently evidenced
that the aprotic electrolytes tend to decompose irreversibly during
discharge/charge, mainly owing to the formation of the reactive
superoxide radical (O2 − ; refs 2,11,12). Even worse, the investigated
electrocatalysts promote the undesired decomposition reactions13
and even the poly(vinylidene difluoride) binder was found to be
unstable12 . It was also shown that the desired discharge product
(Li2 O2 ) forms in ether-based electrolytes, however, still leading to
large overpotentials14–16 .
Replacing lithium by sodium to build an analogous Na–O2
cell with sodium peroxide (Na2 O2 ) as the discharge product still
offers the opportunity to construct a cell system with a high energy
density (E 0 = 2.33 V, wth = 1,605 Wh kg−1 (Na2 O2 ), Supplementary
Information). However, as recently reported17 , this system also
suffers from similar high overpotentials and low energy efficiencies
when using carbonate-based sodium electrolytes. In detail, Na2 O2
was the major discharge product, besides Na2 CO3 and NaOCO-R
species from electrolyte decomposition using a solution of NaPF6
in a 1:1 mixture of ethylene carbonate and dimethyl carbonate.
We also found large overpotentials for this electrolyte system
(Supplementary Fig. S1). In 2011, ref. 18 reported on a different
Na–O2 cell with a polymer electrolyte and a molten sodium anode
operating at 100 ◦ C that could be cycled several times, but did not
specify the exact reaction product.
Interestingly, the reactivity of sodium and lithium towards
oxygen is quite different despite their close chemical relation19,20 .
Sodium can form a stable superoxide NaO2 , whereas LiO2 is highly
unstable and is found only as intermediate species in Li–O2 cells21 .
Thus, in a Na–O2 cell, the formation of NaO2 (sodium superoxide)
during discharge will compete with the formation of Na2 O2 .
Moreover, even though peroxide formation is thermodynamically
favoured (E 0 (Na2 O2 ) = 2.33 V versus E 0 (NaO2 ) = 2.27 V), the
formation of NaO2 requires the transfer of only one electron
per formula unit and will be kinetically preferred relative to the
two-electron transfer towards the peroxide (Fig. 1). Indeed, in an
ether-based electrolyte we found solid NaO2 to be formed reversibly
and exclusively (within the precision of our analytical methods) as
a crystalline product at very low overpotentials.
Our cell comprised a metallic sodium anode, a glass fibre
separator and a carbon-fibre gas diffusion layer (GDL) as
the cathode. As the electrolyte we used a 0.5 M solution of
sodium triflate salt (NaSO3 CF3 ) in diethylene glycol dimethyl
ether (C6 H14 O3 , diglyme, DEGDME). Figure 2a,b shows the first
galvanostatic discharge/charge cycles of the Na–O2 cell together
with results for an analogous Li–O2 cell (LiSO3 CF3 /DEGDME)
that was built for comparison. For the Li–O2 cell, we observed
a voltage plateau of 2.6 V during discharge, corresponding to an
overpotential (ηdis ) of about 300 mV. On charging, the voltage
quickly increased to around 4.2 V (ηchg ≈ 1.3 V), which is typical
for Li–O2 cells and agrees well with previous reports3,22 . Even at
a comparably small current density of 20 µA cm−2 we obtained
a discharge capacity of only 0.06 mAh, reflecting the use of the
pure GDL carbon cathode with low surface area. In contrast, the
Na–O2 cells could be cycled at much higher current densities with
a maximum discharge capacity of about 3.3 mAh at 120 µA cm−2
and 0.13 mAh at 500 µA cm−2 . Even at 500 µA cm−2 the capacity is
still two times higher than for the Li–O2 cell cycled at 20 µA cm−2 .
For low current densities a wide voltage plateau of about 2.2 V,
1 Physikalisch-Chemisches
Institut, Justus-Liebig-Universität Gießen, Heinrich-Buff-Ring 58, 35392 Gießen, Germany, 2 BASF SE, 67056 Ludwigshafen,
Germany.
address: BELLA, Institut für Nanotechnologie, Karlsruher Institut für Technologie, Hermann-von-Helmholtz-Platz 1, 76344
Eggenstein-Leopoldshafen, Germany. *e-mail: [email protected]; [email protected].
† Present
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NATURE MATERIALS DOI: 10.1038/NMAT3486
LETTERS
e¬
¬
e
Lithium¬oxygen
Sodium¬oxygen
e¬
e¬
n
yge
n
¬
O2
A+
e¬
yge
e¬
e¬
Ox
Ox
A+
Li+
¬
¬
A+
A+
Alkali metal
anode
Electrolyte/
separator
Carbon
air cathode
Na + O2 → NaO2
(E0 = 2.27 V)
2Na + O2 → Na2O2
(E0 = 2.33 V)
2Na + 1/2O2 → Na2O (E0 = 1.95 V)
2Li + O2 → Li2O2
2Li + 1/2O2 → Li2O
(E0 = 2.96 V)
(E0 = 2.91 V)
Figure 1 | The general function principle of an alkali metal–oxygen battery. During discharge, metal A is oxidized to a soluble A+ cation at the
anode/electrolyte interface and the electron is transferred to the outer circuit. At the cathode side, oxygen is reduced to an O2 − species (superoxide
radical) that may form an alkali metal superoxide (AO2 ) in the presence of A+ cations. As LiO2 is highly unstable, it further reacts to form lithium peroxide
(Li2 O2 ) either by reduction (LiO2 + e− + Li+ → Li2 O2 ) or by disproportion (2LiO2 → Li2 O2 + O2 ). In the case of sodium, the same reactions may occur, but
sodium superoxide is thermodynamically more stable and is indeed found as discharge product in the present NaO2 cell. Other possible cell reactions that
have not been observed experimentally are shown in grey.
corresponding to ηdis < 100 mV, is observed whereas for high
current densities the cell voltage continuously decreases, in line with
expected kinetic limitations at high currents. The charging process,
at low current densities, occurs at a voltage plateau between 2.3 and
2.4 V, which is close to the potential expected for the decomposition
of NaO2 to form Na and oxygen (E 0 = 2.27 V). These findings are
also supported by results from cyclic voltammetry (Supplementary
Fig. S2). Charging at such low overpotentials is in stark contrast
to the results found for the analogous Li–O2 cell, indicating a
different reaction mechanism in the Na–O2 cell. We note that the
high-temperature sodium cell in ref. 18 was cycled at a maximum
current density of 100 µA cm−2 , and that the overpotential on
charging was about 700 mV at a temperature of 100 ◦ C.
The charge efficiency of the Na–O2 cells is, at least after the
first cycle, between 80% and 90%. Even though the cells can
be discharged and charged several times, the absolute capacities
decreased with increasing cycle number, with negligible energy
storage after 8 cycles (Supplementary Fig. S3). The lower efficiency
of the first cycle compared with the subsequent ones might be
related to a first activation of the electrode, that is, irreversible
reaction of the electrolyte or NaO2 with the surface of the carbon
cathode, for example. The formation of NaO2 as the discharge
product is further supported by recording the open-circuit voltage
(Eoc ) value that instantly approaches the accordant potential after
discharge (Supplementary Fig. S4).
Convincing evidence for the reversible formation/decomposition of NaO2 was obtained from pressure measurements of the
oxygen gas phase with a pressure gauge during cycling (Fig. 2c).
The pressure decreased linearly by about 40 mbar when discharged
by 0.36 mAh, and increased during re-charge back to a value close
to the starting pressure, indicating reversibility of the cell reaction.
Under the open-circuit condition, that is, zero current, the pressure
remained constant at constant potential values. We estimate the
total volume of oxygen and the starting pressure as 8.9 cm3 and
1,053 mbar, and thus, can determine the number of electrons
transferred to oxygen by applying the ideal gas and Faraday’s laws.
The number of transferred electrons per O2 molecule is 0.94±0.09,
which strongly supports the assumption of NaO2 formation.
The results obtained directly from the electrochemical measurements and the corresponding oxygen consumption/release
clearly indicate reversible NaO2 formation in the present Na–O2
cell, but unequivocal evidence is obtained by post-mortem anal2
ysis (Raman spectroscopy, X-ray diffraction (XRD) and scanning
electron microscopy (SEM)) of the cathode after discharge and
charge. Figure 3c shows a SEM image of the cathode (GDL) in
the pristine state consisting of interwoven carbon fibres with an
average diameter of about 10 µm. The carbon fibres themselves are
mostly non-porous with a specific surface area below 1 m2 g−1 as
determined by nitrogen physisorption. In Li–O2 cells described in
the literature typically porous carbons with specific surface areas
exceeding >50 m2 g−1 are used, with electrodes prepared by using
poly(vinylidene difluoride)-, polytetrafluoroethylene - or Nafionbased binders. We also found high capacities for Li–O2 cells when
using high-surface-area carbons (Supplementary Fig. S1). However,
the advantage of using a free-standing, binder-free GDL with a low
surface area as the electrode is that unwanted side reactions due
to binder decomposition12 are excluded and unknown influences
from the carbon surface chemistry are minimized. In addition,
the open, macroporous network allows the accommodation of the
discharge product. It is a major result of this work that the present
Na–O2 cell runs already well even with such a low-surface-area
carbon and without any added catalyst. Figure 3a shows the cathode
at the oxygen interface after discharge to 2 mAh, with a solid
discharge product filling, at distinct areas, almost the whole volume
between the carbon fibres. The typical particle size is estimated to
be about 1–50 µm. Figure 3b shows an image from the opposite
side of the cathode, that is, the side of the cathode that was facing
the separator. Significant amounts of solid discharge products can
also be clearly seen here and we conclude that the cell reaction
took place in the whole cathode volume. The preferable formation of NaO2 at the oxygen side is due to the known transport
limitation of gaseous oxygen through the electrolyte-filled cathode
structure (see also the discussion on the theoretical discharge
capacity in the Supplementary Information)23,24 . Figure 3d shows a
high-resolution image of the formed particles. The particle shows
a morphology with a well-defined macroscopic cubic symmetry.
We note the large size of many product particles, which indicates
at least a certain degree of electrical conductivity of NaO2 . Li2 O2
is considered as an electric insulator25 , with, according to recent
calculations, conducting crystal surfaces26 . As the electrical properties of NaO2 have not been reported yet (a bandgap of 3.3 eV
is calculated)27 , this has to be part of future investigations. For
KO2 a significant electric conductivity at room temperature is
reported in the literature28 .
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NATURE MATERIALS DOI: 10.1038/NMAT3486
a
3.0
120 μA cm¬2
200 μA cm¬2
300 μA cm¬2
500 μA cm¬2
Na¬O2 cell
2.8
E/V versus Na/Na+
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2.6
2.4
Charge
2.2
Discharge
2.0
1.8
b
0.0
0.5
1.0
4.4
1.5
2.0
Q (mAh)
2.5
3.0
3.5
20 μA cm¬2
Charge
E/V versus Li/Li+
4.0
3.6
3.2
ηchg
Eoc
ηdis
2.8
Discharge
2.4
Li¬O2 cell
2.0
0.00
0.02
0.04
0.06
Q (mAh)
c
3.0
Discharge
200 μA cm¬2
Eoc
Charge
200 μA cm¬2
0
Pre
¬10
re d
2.6
¬20
rop
2.4
¬30
2.2
2.0
Δp (mbar)
ssu
E/V versus Na/Na+
2.8
¬40
300
360
420
t (min)
480
540
Figure 2 | Electrochemical characterization of Li–O2 and Na–O2 cells with
a GDL cathode. a, Discharge–charge cycles of Na–O2 cells at various
current densities (that is, the rate capability). Cutoff potentials were set to
1.8 V for discharge and 3.6 V for charge. Dotted line: E0 (NaO2 ) = 2.27 V.
b, Typical discharge/charge hysteresis of an otherwise identical Li–O2 cell
(E0 (Li2 O2 ) = 2.96 V) that was prepared in parallel to the Na–O2 cells. In
comparison with the Li–O2 cell, the Na–O2 cells show a more than ten
times higher discharge capacity at more than ten times higher current
densities, with a ten times lower overpotential for the charging process.
c, Pressure change in the O2 gas reservoir (green line) during discharge and
re-charge of the cell, which allows the determination of the number of
electrons transferred to O2 . Current densities refer to the geometric area of
the electrode. Values referring to the specific surface area and carbon
weight can be found in Supplementary Table S2.
Energy-dispersive X-ray spectroscopy (EDS) spectra of
micrometre-sized product particles were collected to obtain
information on the local chemical composition and uniformity of
the discharge product (Fig. 4a). The strongest signals are related to
characteristic spectral lines for the elements oxygen, sodium and
carbon with a relative atomic composition of about 61%, 33% and
6%, respectively. The carbon signal is caused by the carbon fibre
in the vicinity of the investigated particle. Trace amounts of the
conducting salt (NaSO3 CF3 ) were found indicated by very weak
signals of sulphur and fluorine. As exact quantification of light
elements with EDS is known to be difficult, it is not possible to
extract the actual composition and therefore the chemical identity
of the discharge product by this method. Thus, Raman spectroscopy
and powder XRD were used to evidence NaO2 as the discharge
product. The Raman spectrum of a cathode after discharge is
shown in Fig. 4b and compared with sodium superoxide (NaO2 ),
sodium peroxide (Na2 O2 ) and sodium oxide (Na2 O). Clearly, the
appearance of the intense Raman band at 1,156 cm−1 (ref. 29)
proves that NaO2 is the discharge product. Moreover, the XRD
pattern of a discharged cathode evidences NaO2 being the discharge
product (Fig. 4c). Both the positions of the diffraction lines and the
intensities agree well with JCPDS reference card No. 01-077-0207,
a calculated diffraction pattern of sodium superoxide, on the basis
of experimental data for the room-temperature modification of
NaO2 , a distorted pyrite-type structure with cubic (NaCl, Fm3̄m)
symmetry30,31 . All other reflection maxima can be related to the
gas-tight sample holder (Supplementary Fig. S6) or the carbon
cathode as shown by the corresponding XRD of the pristine
cathode. Here we note that phase-pure NaO2 is chemically very
difficult to synthesize (for example, at 150 bar oxygen pressure and
temperatures of 450 ◦ C (ref. 32), or in liquid ammonia33 ) and is
not available commercially. Instead, here our electrochemical cell
shows an alternative and comparably mild synthesis route.
Even though NaO2 has been clearly confirmed as the discharge
product one could question whether Na and O2 are actually
reversibly formed on charging. For comparison with Li–O2
cells it is important to note that experimental evidence for
Li2 O2 formation and decomposition after subsequent cycling
has been found only recently using nanoporous gold electrodes
and a dimethylsulphoxide based electrolyte, but not with carbon
electrodes34 . To prove the reversibility of the reaction in the case
of Na–O2 , we also collected the XRD pattern after charging and
after the fourth discharge cycle. After charging, it can be clearly
seen that the NaO2 diffraction lines almost disappeared, whereas
after the fourth discharge, NaO2 can still be identified as the
discharge product. This proves that the cell reaction, that is, the
formation and decomposition of NaO2 , is at least partially reversible
over several cycles. We add that we never found evidence for
an unwanted decomposition of discharged cathodes with NaO2 ,
neither under cell conditions nor during later analysis. First results
from differential electrochemical mass spectroscopy (not shown
here) indicate that O2 is the only gaseous product during charging.
We were successful in constructing a room-temperature
sodium–oxygen cell with an ether-based electrolyte that achieved
discharge capacities of over 300 mAh g−1 (carbon), corresponding
to roughly 3.3 mAh cm−2 (electrode area). Cells could be cycled
several times at current densities as high as 0.2 mA cm−2 using
carbon with a specific surface area orders of magnitude smaller than
in studies of Li–O2 cells. As a major breakthrough we consider the
very low overpotential of less than 200 mV during charging, which
is at least a factor of 3–4 times lower than for any other Li–O2 or
Na–O2 cell reported in the literature. The discharge product was
unequivocally identified by XRD and Raman spectroscopy to be
sodium superoxide (NaO2 ). The oxygen reduction reaction occurs
as a single-electron transfer process (O2 + e− → O2 − ) and seems
to be kinetically highly favoured, which explains the reversibility
of the cell reaction. The severe difference in the cathodic reaction
and the reaction products changing from carbonate-based to
ether-based electrolytes may be explained by the stability of NaO2
and the improved stability of ethers against the highly nucleophilic
superoxide anion O2 − (refs 9–11). The results demonstrate that
the sodium-based cell chemistry might offer—compared with
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NATURE MATERIALS DOI: 10.1038/NMAT3486
LETTERS
a
b
100 µm
100 µm
c
d
100 µm
10 µm
Figure 3 | SEM images of the Na–O2 cell cathodes. a,b, SEM images of a cathode structure after discharge to 2 mAh at 80 µA cm−2 at the oxygen/GDL
and GDL/separator side, respectively. c, For comparison, a pristine cathode structure. d, Solid products formed on the carbon fibres during discharge at
higher magnification. For a discussion on the theoretical capacity achievable, see Supplementary Fig. S5.
a
b
1,156 cm¬1
Focus point
Discharge product
O¬K
l (a.u.)
l (a.u.)
Na¬K
NaO2
Na2O
Na2O2
C¬K
0.0
0.5
1.0
1,500
1.5
1,200
300
ν (cm¬1)
(422)
(331)
(420)
(400)
(311)
(222)
∗
∗
∗
l (a.u.)
(220)
(200)
c
600
900
E (keV)
First discharge
(6 mAh)
First charge
Fourth discharge
(1 mAh)
Pristine cathode
30
40
50
60
2θ (°) (Cu-Kα)
70
80
Figure 4 | Analysis of the discharge products in Na–O2 cells. a, The local distribution of elements in a discharged cathode as measured by EDS and SEM.
The inset shows the SEM image (about 75 µm × 45 µm) and the cross indicates the location of the EDS measurement. b, Raman spectra, taken in a
microscope, of a particle in a discharged cathode and for commercial material of sodium peroxide (Na2 O2 ) and sodium oxide (Na2 O). Reference material
of NaO2 was provided by M. Jansen (MPI-FKF, Stuttgart, Germany). c, XRD pattern of a pristine, a discharged, a re-charged and a four-times-cycled Na–O2
cell cathode. The Miller indices on the reflections correspond to NaO2 according to JCPDS reference card No. 01-077-0207. The asterisks denote
reflections from the sample holder. All diffraction patterns were background-corrected and normalized to the reflection at 2θ = 41.5◦ .
4
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NATURE MATERIALS DOI: 10.1038/NMAT3486
lithium-based cells—unexpected opportunities in the search for
reversible energy storage devices.
Methods
Cell assembly. The used cell hardware was a modified 1/2 inch Swagelok design,
consisting of lithium (Chemetall) or sodium metal foil (donated by BASF SE)
as the anode, glass microfibre filters (GF/A, Whatman) and a binder-free GDL
(Freudenberg H2315, Quintech) as the cathode. The porosity of the material was
estimated to be around 80% on the basis of the geometric area weight of around
95 g m−2 provided by the manufacturer. The Brunauer–Emmett–Teller surface
area of the GDL was determined to be below 1 m2 g−1 . No catalyst was used.
The average cathode area, thickness and mass were 1.13 cm2 , 210 µm and 11 mg,
respectively. DEGDME (anhydrous, 99.5% Sigma Aldrich) was used as the solvent
for the electrolyte. Lithium triflate (LiSO3 CF3 , 99.995% trace metal basis, Aldrich)
and sodium triflate (NaSO3 CF3 , 98%, Aldrich) were used as conducting salts and
dried under vacuum at 75 ◦ C for 24 h. The electrolyte solutions of 0.5 M LiSO3 CF3
or NaSO3 CF3 in DEGDME were prepared in a glove box. The final water contents
of the electrolytes were determined with an 831KF Karl Fischer coulometer
(Metrohm) to be less than 20 ppm. The amount of electrolyte in the cell was 85 µl,
and a further 30 µl was added for the connection to the reference electrode. Cell
assembly was carried out in argon-filled glove boxes (Labmaster, MBraun and
GST4, Glovebox Systemtechnik) at water and oxygen contents below 1 ppm.
Electrochemical cell testing. Cell tests were performed galvanostatically at room
temperature, using battery cycling systems from Maccor (4300) and Biologic
(VMP3). Before measurements, the cells were flushed with oxygen (purity
5.0, Praxair) for 10 s at 105 Pa. The total volume of the oxygen reservoir was
approximately 8.9 cm3 . The lower and upper cutoff limits for the voltage were
2.0 V versus Li/Li+ and 4.2 V versus Li/Li+ and 1.8 V versus Na/Na+ and 3.6 V
versus Na/Na+ , respectively. Every measurement started with a 2 h recording of
the open-circuit potential to ensure equilibrium in the cell. The pressure of the
oxygen gas reservoir was measured with a pressure transducer, PAA-33X (Omega
Engineering, Supplementary Fig. S7).
Structural characterization. Structural characterization of the discharge products
was carried out using an X’Pert Pro (PANalytical) powder X-ray diffractometer
(Cu-Kα source, 40 kV, 40 mA) and a Senterra Raman microscope (Bruker, 2 mW
laser power at 523 nm). For both experiments, a self-made gas-tight sample holder
was used. The specific surface area (Brunauer–Emmett–Teller) was determined by
nitrogen physisorption at 77 K using an Autosorb-1 machine from Quantachrome
Instruments. SEM investigations were performed on a Merlin high-resolution
Schottky field-emission electron microscope (Zeiss SMT) equipped with an X-Max
EDS detector (Oxford Instruments).
Received 25 April 2012; accepted 9 October 2012; published online
2 December 2012
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Acknowledgements
The research was supported by the BASF International Scientific Network for
Electrochemistry and Batteries. P. Hartmann is grateful to Fonds der chemischen
Industrie (FCI) for a scholarship. The authors thank M. Ante, B. Jache and C. Raiß for
experimental support. We further thank H. Heidt, H. Weigand, G. Pfeiffer and S. Lember
for technical support. We are indebted to M. Jansen (Max-Planck-Institute for Solid State
Research) for providing phase-pure bulk NaO2 as a reference material.
Author contributions
P.A., P.H. and J.J. designed this study. P.H. and C.L.B. carried out the electrochemical
experiments and XRD analysis. M.V. developed the metal–air cell set-up for the battery
tests. P.H. developed the gas pressure set-up and conducted the SEM, EDS and Raman
spectroscopy experiments. P.H., P.A. and J.J. analysed and discussed the results and
wrote the manuscript. A.K.D. and A.G. contributed to the scientific discussion. P.A. and
J.J. supervised the research project.
Additional information
Supplementary information is available in the online version of the paper. Reprints and
permissions information is available online at www.nature.com/reprints. Correspondence
and requests for materials should be addressed to J.J. or P.A.
Competing financial interests
A US-Provisional Patent Application (US 61/615901) directed to sodium oxygen cells as
described in the manuscript has been filed by BASF SE with the USPTO. A.K.D. and A.G.
are employees of BASF SE.
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