The Effect of Water on Quinone Redox Mediators in Nonaqueous Li-O2 Batteries

Article abstract of DOI:10.1021/jacs.7b11007

The parasitic reactions associated with reduced oxygen species and the difficulty in achieving the high theoretical capacity have been major issues plaguing development of practical nonaqueous Li-O₂ batteries. We hereby address the above issues by exploring the synergistic effect of 2,5-di-tert-butyl-1,4-benzoquinone and H₂O on the oxygen chemistry in a nonaqueous Li-O₂ battery. Water stabilizes the quinone monoanion and dianion, shifting the reduction potentials of the quinone and monoanion to more positive values (vs Li/Li⁺). When water and the quinone are used together in a (largely) nonaqueous Li-O₂ battery, the cell discharge operates via a two-electron oxygen reduction reaction to form Li₂O₂, with the battery discharge voltage, rate, and capacity all being considerably increased and fewer side reactions being detected. Li₂O₂ crystals can grow up to 30 μm, more than an order of magnitude larger than cases with the quinone alone or without any additives, suggesting that water is essential to promoting a solution dominated process with the quinone on discharging. The catalytic reduction of O₂ by the quinone monoanion is predominantly responsible for the attractive features mentioned above. Water stabilizes the quinone monoanion via hydrogen-bond formation and by coordination of the Li⁺ ions, and it also helps increase the solvation, concentration, lifetime, and diffusion length of reduced oxygen species that dictate the discharge voltage, rate, and capacity of the battery. When a redox mediator is also used to aid the charging process, a high-power, high energy density, rechargeable Li-O₂ battery is obtained.

Full text of DOI:10.1021/jacs.7b11007

Article
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Cite This: J. Am. Chem. Soc. XXXX, XXX, XXX−XXX
The Effect of Water on Quinone Redox Mediators in Nonaqueous
Li‑O2 Batteries
Tao Liu,† James T. Frith,‡ Gunwoo Kim,†,§ Rachel N. Kerber,† Nicolas Dubouis,†,∥ Yuanlong Shao,§
Zigeng Liu,†,⊥ Pieter C. M. M. Magusin,† Michael T. L. Casford,† Nuria Garcia-Araez,‡
and Clare P. Grey*,†
†Chemistry Department, University of Cambridge, Lensfield Road, Cambridge CB2 1EW, United Kingdom
‡Chemistry Department, University of Southampton, Highfield Campus, Southampton SO17 1BJ, United Kingdom
§Cambridge Graphene Center, University of Cambridge, 9 JJ Thomson Avenue, Cambridge CB3 0FA, United Kingdom
S
* Supporting Information
ABSTRACT: The parasitic reactions associated with reduced
oxygen species and the difficulty in achieving the high
theoretical capacity have been major issues plaguing develop-
ment of practical nonaqueous Li-O2 batteries. We hereby
address the above issues by exploring the synergistic effect of
2,5-di-tert-butyl-1,4-benzoquinone and H2O on the oxygen
chemistry in a nonaqueous Li-O2 battery. Water stabilizes the
quinone monoanion and dianion, shifting the reduction
potentials of the quinone and monoanion to more positive
values (vs Li/Li+). When water and the quinone are used together in a (largely) nonaqueous Li-O2 battery, the cell discharge
operates via a two-electron oxygen reduction reaction to form Li2O2, with the battery discharge voltage, rate, and capacity all
being considerably increased and fewer side reactions being detected. Li2O2 crystals can grow up to 30 μm, more than an order of
magnitude larger than cases with the quinone alone or without any additives, suggesting that water is essential to promoting a
solution dominated process with the quinone on discharging. The catalytic reduction of O2 by the quinone monoanion is
predominantly responsible for the attractive features mentioned above. Water stabilizes the quinone monoanion via hydrogen-
bond formation and by coordination of the Li+ ions, and it also helps increase the solvation, concentration, lifetime, and diffusion
length of reduced oxygen species that dictate the discharge voltage, rate, and capacity of the battery. When a redox mediator is
also used to aid the charging process, a high-power, high energy density, rechargeable Li-O2 battery is obtained.
being characterized by electron spin resonance, ultraviolet−
visible (UV−vis) spectroscopy, and theoretical calcula-
tions.22−24
1. INTRODUCTION
Quinones represent an important class of organic redox
molecules that are involved in energy transduction and storage
in biological systems.1−4 For example, they play a pivotal role in
proton-coupled electron transfer for natural respiratory and
photosynthetic processes.5 This unique charge transfer role of
quinones inspired researchers to explore their applications in a
range of artificial energy harvesting and storage devices,
including dye-sensitized solar cells, artificial photosynthesis,
pseudocapacitors, organic lithium ion batteries, redox flow
batteries, and so on.6−9 This versatility of quinones in part
stems from the ease by which their physicochemical properties
(redox, solubility, optical, and electrical properties) can be
tuned by engineering the molecular structures and through
interactions with their chemical environment.10−12 For
example, the redox potentials of many quinone systems in
nonaqueous media can be shifted to more positive values via
the use of additives containing O−H and N−H bonds.13−17
These observed shifts have been rationalized by hydrogen-bond
formation of the negatively charged carbonyl oxygens on the
reduced quinones with the hydrogen atoms from water,
alcohols, or amines,18−21 the nature of the hydrogen bond
Recently, quinones have also been explored as redox
mediators for the oxygen reduction reaction (ORR) in Li-O2
batteries. The nonaqueous Li-O2 battery is considered the
ultimate battery as it possesses a theoretical energy density
close to gasoline, 10 times higher than the state-of-the-art
lithium ion battery.25−28 Its operation typically involves O2
reduction during discharge, the first step involving a one-
electron electrochemical step to form LiO2, which then
chemically disproportionates to form Li2O2; a solid phase
precipitates out of the liquid electrolyte and deposits on the
porous electrode. On charging, the solid discharge product is
decomposed releasing O2. Realizing the theoretical capacity is,
however, associated with significant challenges, in part because
an electronically insulating discharge product tends to form as
small particles or conformal films that quickly passivate
electrode surfaces,29−31 impeding further interfacial electron
Received: October 23, 2017
© XXXX American Chemical Society
A
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transfer and ion diffusion through the porous electrode. As a
result, the cell discharge tends to finish early before the
discharge product fully takes up the free volume available in the
porous electrode. Furthermore, electrolyte decomposition
promoted by reactions with reduced oxygen species,32−42
particularly LiO2, occurs during discharge. Therefore, enabling
a mechanism that minimizes surface passivation, greatly
promotes crystal growth of the discharge product, and reduces
the amount of side reactions, is key to realizing the full
potential of a Li-O2 battery.
Several redox couples,43−48 including viologens,43,44 phtha-
locyanine,45 and quinones,46,47 have been used to address the
issues associated with the discharge of nonaqueous Li-O2
batteries. A common feature is that these soluble molecules
are able to chemically reduce O2 in solution. The formation of
surface passivation films is inhibited (to some extent), and the
discharge capacity is increased. Another important aspect
concerns their potential ability to reduce the lifetime43,44 and
decrease the free energies47 of the chemically aggressive
reaction intermediates, so that fewer side reactions occur.
In this work, we take advantage of the hydrogen-bonding
properties of quinones13−21 and evaluate the impact of the use
of water with the quinone, 2,5-di-tert-butyl-1,4-benzoquinone
(DBBQ), on the battery rate, capacity, and side reactions. We
show that the interactions with water stabilize the quinone
monoanion, dianion, and reduced oxygen intermediate species,
the discharge becoming dominated by solution-phase pro-
cesses. As a result, the rate and capacity are significantly
improved over the reactions in the dry electrolyte. Importantly,
the extent of side reactions is decreased by around 70%. We
propose potential mechanisms with the aid of supporting
experimental measurements and DFT calculations and discuss
the significance of the work.
Figure 1. Linear voltage sweep measurements in a three-electrode
setup of a cell as a function of the water content (in ppm, as labeled for
each curve) in a 10 mM DBBQ, 0.25 M LiTFSI/DME electrolyte. A
GDL electrode was used as the working electrode, and the sweep rate
is 10 mV/s for all experiments. Because of the high-water content
used, lithium iron phosphate was used as the reference and counter
electrodes. All potentials are referenced against Li/Li+ (−3.04 V versus
standard hydrogen electrode).
extracted and subjected to UV−vis measurements to obtain
reference UV−vis spectra of DBBQ, its anion and dianion
(Figure 2A). Optically, the electrolyte underwent color changes
from green (Q) to brown (Q−•) and light pink (Q2−). The
quinone UV−vis spectrum exhibits a major absorption at 255
nm and a weaker peak at 305 nm. For the quinone monoanion,
in addition to the peaks at 255 and 305 nm, the spectrum also
shows a broad absorption from 350 to 450 nm and peaks at
235, 315, and 325 nm. The dianion spectrum has very distinct
features: the absorption at 320 nm becomes the most intense
peak, followed by the one at 240 nm, and a broad absorption at
around 400 nm.
Moving to the case with 30,000 ppm water, the quinone
spectrum shows absorption peaks at 254 and 300 nm, both
slightly blue-shifted compared to that in the absence of water.
On potentiostatically discharging the cell to only 2.5 V, the
electrolyte became light pink in color indicating that the
dianion is present. The capacity recorded was 90% of the
theoretical value expected for two electrons per quinone
molecule and further discharge at 1.9 V led to negligible
capacity increase, confirming that almost all of the quinone had
been reduced to the dianion at 2.5 V. The corresponding UV−
vis spectra further support this view as the spectra (Figure 2B)
at 2.5 and 1.9 V are nearly identical, showing a major peak at
295 nm and two weaker absorptions at 220 and 230 nm.
Compared to the anhydrous dianion spectrum (Figure 2B),
the weak broad absorption at 400 nm was suppressed and the
other peaks were greatly blue-shifted in the presence of water.
The water-induced blue shifts were confirmed by adding water
to the anhydrous dianion sample (Figure 2B, gray spectra).
2. RESULTS AND DISCUSSION
2.1. Stabilization of DBBQ Monoanion and Dianion by
Water. The effect of added water on the redox chemistry of
DBBQ is first evaluated. Figure 1 shows the linear voltage
sweep experiments of DBBQ as a function of the water content
in a lithium bis(trifluoromethanesulfonyl)imide/dimethoxy-
ethane (LiTFSI/DME) electrolyte. Under nominally dry
conditions (<10 ppm water), two reduction peaks of DBBQ
were observed, one at 2.5 V and the other at 2.15 V. These
peaks are attributed to sequential one-electron reductions of
the quinone (Q) to the quinone monoanion (Q−•) (i) and
then to the quinone dianion (Q2−) (ii), each DBBQ molecule
eventually taking up two electrons. This type of redox behavior
is commonly observed for quinones in nonaqueous media.15,18
As the water concentration was increased, the Q2− peak shifted
positively by up to +0.4 V, whereas the Q−• peak position
shifted only slightly by +0.05 V. Above 20,000 ppm (1 M) H2O
content, the quinone monoanion and dianion peaks merged
together into a single peak; the peak area being the sum of
those measured under anhydrous conditions (Figure S1 details
the quantitative analysis.). A similar behavior was observed in
an electrolyte with diglyme as the solvent and with another
hydrogen-bonding donor molecule, methanol (Figure S1).
To investigate the effect of water on the quinone
electrochemistry further, we measured the UV−vis spectra of
the quinones in anhydrous and wet conditions (Figure 2). In
the anhydrous case, the cell was potentiostatically discharged at
2.5 and 1.9 V until the current dropped to approximately zero.
Electrolyte samples at the corresponding voltages were
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Figure 2. UV−vis spectra of DBBQ at different states of oxidation in an (A) anhydrous and (B) 3% H2O (30,000 ppm) added 0.25 M LiTFSI/DME
electrolyte. The monoanion and dianions were obtained by potentiostatically reducing the quinone at 2.5 and 1.9 V, respectively, and their
corresponding spectra are color coded, green for the quinone, brown for the quinone monoanion, and pink for the dianion. For the case with added
water, the quinone species generated potentiostatically at 2.5 and 1.9 V are presented with similar color codes. The effect of added water on the UV−
vis spectrum of the dianion was investigated by adding an equivalent amount of H2O (3%) to the dianion sample obtained in the anhydrous
electrolyte (gray). The sharp peak at 208 nm is very close to the spectrum cutoff and does not shift on addition of water. Its origin is unclear.
Figure 3. DFT calculations for the Gibbs free energy of formation for the monoanion LiQ• (a) and the dianion Li2Q (b) in the presence of 0, 2, 4,
and 6 coordinating water molecules. 20,000 ppm water in the electrolyte, as used experimentally in this work, corresponds to a water/Li+ molar ratio
of 4. The relevant energy minimized structures that result from these calculations are shown below, where the number of coordinating water used in
the models is color coded from light to dark blue with the increasing H2O content. The green, dark gray, red, and white balls correspond to Li, C, O,
and H, respectively. The hydrogen bonds are illustrated with black dashed lines. In the absence of water, the Li+ ions strongly coordinate to the
oxygen atoms (OQ) of the reduced quinone anions. The addition of water results in a gradual elongation of the r(Li-OQ) bond, and ultimately, in the
case of the dianion with six H2O molecules, complete shielding of the Li+ from Q2− occurs, each Li+ ion being solvated by three water molecules. The
effect of the dispersion force correction (Grimme’s dispersion with the original D3 damping function, see Materials and Methods in the Supporting
Information) is stronger for the case of the monoanion than for the dianion structural models, which is ascribed to the presence of more extensive
hydrogen-bonding (H−O) interactions in the monoanion models.
These shifts and the positive shifts of the half-wave potentials of
reduced DBBQ due to added water are consistent with the
hydrogen-bonding effects reported for many other quinones in
nonaqueous media.13,14,18,19,21,24
Figure 3 shows the density functional theory (DFT)
calculations for the Gibbs free energy of formation of the
monoanion and the dianion in the presence of increasing
numbers of coordinating water molecules. These calculations
indicate that with increasing water concentrations, the energies
of formation become more negative (more favorable),
consistent with the positively shifted reduction potentials.
Moreover, as illustrated in structures shown in Figure 3A,B, this
increased thermodynamic stability of the reduced quinone
anions appears to result from hydrogen-bond formation
between the negatively charged monoanion/dianion oxygens
and the protons of water, hydrogen-bonding between water
molecules (black dashed lines) and the coordination of the
lithium ions by water. Because of the higher negative charge
density in Q2− versus Q−• and the presence of two Li+ ions per
anion, stronger interactions with water and thus larger positive
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Figure 4. Electrochemistry of Li-O2 batteries with different electrode structures and electrolytes (A−C) and SEM characterization of discharged
electrodes (D−K). Linear voltage sweep measurements (A) of cells using GDL working electrodes with different electrolytes under O2 (as labeled:
water content in ppm and 10 mM for DBBQ concentration). Galvanostatic discharge curves of cells made of super P (B) and reduced graphene
oxide (C) electrodes either with neat 0.25 M LiTFSI/DME electrolyte, with only DBBQ added, or with both DBBQ and H2O added to the neat
electrolyte (as labeled in the figures). D and E represent pristine super P and rGO electrodes. F (G), H (I), J (K), respectively, represent super P
(rGO) electrodes discharged in an anhydrous neat electrolyte, a neat electrolyte with 10 mM DBBQ, and an electrolyte with both 10 mM DBBQ
and 20,000 ppm water added. All cells in (B and C) were discharged at 0.1 mA/cm2; 15 mAh/cm2 for SP and rGO electrodes is equivalent to 15,000
mAh/gc and 150,000 mAh/gc, respectively.
potential shifts are usually observed for the dianion than for the
monoanion.15,18
with no DBBQ) leads to a peak current of −2.2 mA/cm2 and
there is no shift in the onset reduction potential (green curve).
To evaluate the effect of wet DBBQ on the discharge
capacity, Li-O2 batteries made of super P carbon and reduced
graphene oxide (rGO) electrodes were investigated. Figure
4B,C shows the galvanostatic discharge curves of the Li-O2
batteries with different electrolyte compositions. For both super
P and graphene electrodes, the combined use of DBBQ with
water leads to large capacity increases (up to 40 times with
20,000 ppm of H2O) compared with those of using neat
electrolyte or only DBBQ; the respective discharge plateaus
also shift to higher voltages, consistent with positively shifted
reduction potentials of DBBQ in the presence of water (Figures
1 and 4A). The corresponding SEM images (Figure 4D−K) of
the discharged electrodes reveal that in the neat electrolyte, the
discharge product forms as small disc-like particles (F) or
conformal films (G) covering the electrode surface. This is
consistent with the previous results obtained in weak solvating
electrolytes, where Li2O2 formation is dominated by a surface
mechanism.29,30 With added DBBQ, thicker layers of
agglomerates comprising small particles form on the discharged
super P (H) and rGO (I) surfaces. In the case of DBBQ and
water, large 10−20 μm sized particles were observed for both
discharged super P (J) and rGO electrodes (K), and bare super
P and rGO electrode surfaces can be still seen (more SEM
images in Figure S2). This result supports an enhanced
solution-mediated mechanism during discharge, which accounts
2.2. Improved Rate and Capacity. Having established the
effect of water on the redox chemistry of DBBQ, we next
explore the role of water and DBBQ in a Li-O2 battery. Figure
4A compares the linear voltage sweeps of cells with different
electrolytes and atmospheres. To ensure a consistent electrode
surface area for all tests, commercial gas diffusion layers (GDL)
were used as the working electrode for all cells. Without DBBQ
(blue curve), the ORR shows a peak current of −0.8 mA/cm2,
with an onset potential at 2.6 V. When only DBBQ was added
to the electrolyte (red curve), both reduction peaks associated
with the quinone monoanion (Q−•) and dianion (Q2−)
increased in the presence of O2, the former (−2 mA/cm2)
being around 4 times higher in O2 than in Ar. These
observations suggest that the quinone monoanion catalyzes
the ORR, consistent with previous reports.46,47 Adding a water
content of 20,000 ppm (black curve) resulted in a greatly
enhanced reduction current, the corresponding reduction peak
(−5.2 mA/cm2) being around an order of magnitude higher
than that without DBBQ and water (blue curve), and increased
by more than 2.5 times than that observed when using DBBQ
alone (red curve). Furthermore, the onset voltage for reduction
current is increased from 2.70 to 2.85 V, consistent with the
positively shifted reduction potential of DBBQ (Figure 1) due
to water. Notably, the aforementioned effects cannot be solely
due to added water, because the addition of water only (and
D
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Figure 5. XRD (A), solid-state 7Li and 1H MAS NMR spectra of discharged electrodes (B, C) and operando DEMS measurements (D−E) of the Li-
O2 battery system. (A) The XRD patterns of electrodes fully discharged in anhydrous and wet (20,000 ppm of H2O) 0.01 M DBBQ 0.25 M LiTFSI/
DME electrolyte (as labeled). The XRD patterns of Li2O2 and LiOH reference powders are plotted for comparison. 7Li NMR (B) and quantitative
1H NMR spectra (C) comparing the amount of side reactions involved in cells made of super P electrodes and different electrolytes (0.25 M
LiTFSI/DME without any additive, with only DBBQ, and with DBBQ and 20,000 ppm water, as labeled). DEMS signals measured in a continuous
flow mode for cells with a nominally anhydrous (D, discharging; F, charging) and 2% H2O added (E, discharging; G, charging) electrolytes (0.01 M
DBBQ 0.25 M LiTFSI/diglyme). Diglyme was used because DME evaporates too fast in a continuous flow mode, even at a slow flow rate of 80 μL/
min. Red, black, blue, and pink colors represent the cell voltage, O2, H2, and CO2 mass spectrometry signals, respectively. The green broken lines
show the oxygen signals expected for an ideal 2 electrons per O2 consumption/evolution.
for the large capacities observed. Similar phenomena were
observed with GDL electrodes, as used in a previous study.47
The discharge capacity was considerably increased when DBBQ
was used together with water, compared to DBBQ alone; larger
discharge particles and agglomerates of particles were formed in
the discharged GDL electrodes (Figure S3). Of note, the bare
surfaces (Figure S2) and open space (H) available in the rGO
electrodes suggest that the cell can in principle continue to
discharge; the termination of its discharge was in fact limited by
the loading of the LiFePO4 (LFP) counter electrode (around
80 mg), used so as to avoid problems with lithium metal in the
presence of excess water. Compared with a mesoporous
structure, a macroporous electrode structure enables the
growth of larger crystals and appears to allow unhindered
diffusion of redox species even at deep discharges.
2.3. A Li2O2 Chemistry, but with Fewer Side Reactions.
Because of the large amount of water used, it is important to
verify the chemical nature of the discharge product. XRD
measurements for discharged super P electrodes (Figure 5A)
confirm that Li2O2 is the only crystalline discharge product in
DBBQ electrolytes with and without added wateressentially
no crystalline LiOH is formed;31,49 the former case leads to
much higher Li2O2 crystallinity, as also evidenced by SEM
(Figure 4J). 7Li solid-state magic angle spinning nuclear
magnetic resonance (MAS NMR) measurements (Figure 5B)
show a single resonance at around 0.35 ppm, suggesting that
Li2O2 is the dominant discharge product.31,50 Differential
electrochemical mass spectrometry (DEMS) experiments were
also conducted to confirm the e−/O2 molar ratio during the
course of cycling. It can be seen for both cases that during
discharge, the e−/O2 molar ratio stayed close to 2 and no
obvious CO2 or H2 evolution was observed. This suggests a
process that consumes 2 electrons per O2 reduction and is
consistent with Li2O2 formation. Operando pressure measure-
ments performed on discharging also strongly support that the
quinone-mediated ORR closely follows 2e− per reacting O2, at
rates up to 1 mA/cm2 (Figure S4). To determine the extent of
any side reactions of the quinone-involved chemistry, we
1
performed quantitative H MAS NMR measurements by using
a spin flip angle of 30° (π/6) and a sufficiently long recycle
delay of 200 s. The 1H NMR spectrum (Figure 5C) for the neat
electrolyte exhibits resonances at 1.9, 3.5, 8, and −1.5 ppm,
which represent the typical side reaction products of lithium
acetate, methoxide, formate, and hydroxide, respectively, in an
ether-based electrolyte. These side reactions are mainly caused
by nucleophilic attack of ether by the superoxide and
peroxide.31−38 The use of DBBQ led to little change in the
nature or the quantity of the side reaction products. In the
presence of both DBBQ and 20,000 ppm water, however,
1
around 70% reduction in the total H signal intensity was
observed, demonstrating that all of the aforementioned side
reaction products were reduced. This conclusion is consistent
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with a corresponding narrower 7Li NMR resonance centered at
0.35 ppm, that is, fewer types of lithium-containing chemical
environments are present (Figure 5B).
The rechargeability of the battery was also assessed (Figure
S5), charging being associated with a plateau at 3.6 V. In the
subsequent cycling, however, the 3.6 V plateau shortens, and an
additional charging plateau at 4.2 V appeared that dominated
the charging process; the Coulombic efficiency decreases
steadily on cycling. A high-voltage plateau (above 4.0 V) is
typically observed in anhydrous Li2O2-based Li-O2 batteries. To
explore this further, DEMS measurements were performed on
charging: Figure 5F,G compares gas evolutions following a
discharge, as presented in Figure 5D,E, respectively. For the
case with DBBQ alone, the battery can only recharge to around
one-third of the prior discharge capacity (i.e., 33% Coulombic
efficiency) under the conditions used in the DEMS setup, with
the charging voltage rapidly rising from 3.4 to 4.5 V. The
corresponding O2 signal (F) is considerably less than that
expected for a process involving 2 electrons per O2 evolved, and
there is an accompanying CO2 evolution, suggesting that side
reactions occur on charging. The negative H2 intensity is a
result of a decrease in the H2 background signal and then a
baseline correction of the data. This H2 intensity decrease could
also be due to side reactions that consume electrolyte
components contributing to the H2 background signal (e.g.,
from diglyme and H2O). On the other hand, the case with a
combined use of DBBQ and water is associated with a rapidly
increasing voltage until a charging plateau at 4.2 V (G) is
observed, the recharging finishing with 69% Coulombic
efficiency at 4.5 V. The corresponding O2 signal on charging
rose to the expected level of 2 electrons per O2 evolved. It then
dropped down by half and afterward gradually climbed up again
to the expected level. It is clear in this case that the e−/O2
molar ratio stays closer to 2, compared to the case without
water, and the rechargeability is improved with water. The dip
in the O2 DEMS signal at the beginning of the charge is
tentatively ascribed to a competing water oxidation reaction,
which involves 4 electrons per O2 evolved, and thus a decrease
in the O2 evolution signal. This O2 signal decrease coincides
with degradation reactions evolving H2. As the water oxidation
reactions and other side reactions involving CO2 evolution (G)
continued, the water content was reduced and/or the carbon
surface became passivated, which slowly shifts the charging
reaction closer to 2 electrons per O2 evolution, the ratio
expected for Li2O2 decomposition. These observations are
consistent with our proposal that the gradual increase in the cell
charging voltages (Figure S5) to potentials similar to those seen
in the absence of added water are due to the electrochemical
loss of water in the electrolyte in the DEMS cell.
Figure 6. Electrochemistry of Li-O2 batteries made using a 0.25 M
LiTFSI/diglyme electrolyte with DBBQ (10 mM), TEMPO (20 mM)
mediators, and 20,000 ppm added water. Comparison of cells with and
without both redox mediators are shown in (A). The cycling stability
of a cell with both mediators at 0.5 mA/cm2 and with its capacity
limited to 1 mAh/cm2is shown in (B). Carbonized polyacrylonitrile
electrodes were used here, where the macroporous electrode structure
facilitates diffusion of redox species at higher rates.
(compared to the case without TEMPO). It is clear that by
using DBBQ, TEMPO, and water, the energy efficiency and
rechargeability of the battery have been improved.
In summary, we have shown that with a combined use of
DBBQ and water, the discharge rate, capacity, and over-
potentials of the Li-O2 battery are all improved. The oxygen
reduction becomes dominated by a solution-phase process.
When a charging mediator is also added, the rechargeability of
the battery is further improved. This opens up the possibility of
developing a more efficient, high-power Li-O2 battery by
choosing alternative charging mediators with lower redox
potentials than TEMPO and by further optimizing the
mediator concentrations. Importantly, fewer side reactions are
involved during both cell discharge and charge in the battery. In
the next section, we discuss the potential mechanisms
responsible for the attractive features of the current oxygen
chemistry.
The water oxidation issue can be resolved by introducing the
soluble charging mediator 2,2,6,6-tetramethyl-1-piperidinyloxy
(TEMPO) together with DBBQ (Figure 6). TEMPO was
chosen since it has previously been demonstrated as an effective
charging mediator in an ether electrolyte.51 By using just 20
mM TEMPO, the charging voltage plateau (A) decreases to
around 3.6 V; operando pressure measurements confirm oxygen
evolution on charging the battery (Figure S6). At 1 mA/cm2,
the cell still has a capacity of more than 4 mAh/cm2 (or 4000
mAh/gc), whereas without DBBQ and TEMPO, the cell has
very little capacity on charging even at a lower rate of 0.5 mA/
cm2; the voltage rapidly rising to the cutoff value of 4.5 V.
When the discharge capacity is curtailed at 1 mAh/cm2 (or
1000 mAh/gc) (Figure 6B), the cell cycles much more stably
2.4. Mechanism Interpretation. To enable a fast-rate,
solution dominated discharge process with fewer side reactions
in lithium oxygen batteries, the following factors are important:
F
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Figure 7. UV−vis spectra performed to determine the ability of the quinone monoanion and dianion to chemically reduce O2 at dry and wet
conditions. UV−vis spectra of the monoanion (Q−•), dianion (Q2−) generated under anhydrous conditions (0.25 M LiTFSI/DME, 0.01 M DBBQ),
and the dianion generated under wet conditions (30,000 ppm of H2O) (A), after exposing them to an excess of pure O2; the spectra prior to O2
exposure are plotted for comparison. The time-dependent reaction of the quinone dianion with O2 under wet conditions (30,000 ppm of H2O) is
shown in (B). UV−vis spectra of quinone monoanions mixed with wet electrolytes (4000, 20,000 and 40,000 ppm water) and then exposed to O2
(C). The anhydrous monoanion and dianion react immediately to reform the quinone (A). In the wet case, the dianion spectrum remained
unchanged after O2 exposure (A and B), the extent of reaction only becoming appreciable after many hours, the absorption (254 nm) due to Q
gradually increasing, while those associated with O2 and Q2− (295 nm) decrease (arrows in B). The general increase in the spectral background in
the UV region is due to absorption of excess O2 (blue curve in (A), obtained by exposing an anhydrous DME solvent to O2). Quinone monoanion
(Q−•) in the presence of water (4000−40,000 ppm) can still readily react with O2, itself being oxidized back to the quinone state (Q).
(1) the concentration of dissolved O2 in the electrolyte, (2) the
concentration and the thermodynamic stability of the reaction
intermediates (which control their lifetime, diffusion length in
the electrolyte, and chemical reactivity), and (3) the solubility
and thermodynamic stability of the final discharge product(s).
The ensuing discussion focuses on the impacts of water on the
above aspects in the current DBBQ-mediated Li-O2 battery
system.
In the absence of water, a quinone molecule is first
electrochemically reduced to form a lithium quinone
monoanion, LiQ•, as a soluble solvated species (sol) in the
electrolyte (reaction 1). The diffusion length of the soluble
monoanion away from the surface can potentially be very long,
until it meets an O2 molecule and chemically reduces it. This
chemical reduction of O2 to Li2O2 could proceed via two
pathways: One involves LiQ•
reacting with O2 to form
(sol)
In the anhydrous case, the galvanostatic discharge (Figure
4B,C) occurs between 2.3 and 2.6 V, where only the quinone
monoanion can be formed according to the CV obtained under
anhydrous conditions (Figure 1). Since the reduction potential
of Q/Q−• is higher than that of O2/O2− (Figures 1 and 2), the
monoanion will necessarily be generated prior to the formation
of superoxide anions during discharging. Previous studies36,37
suggested that the quinone monoanion can chemically reduce
O2, where the formation of Li2O2 was confirmed in the solid
precipitate after mixing the monoanion with O2. To verify the
proposed reactions of the monoanion Q−• with O2, the reduced
quinone species generated in anhydrous conditions was
exposed to pure O2. The corresponding UV−vis spectrum
(Figure 7) shows that the quinone monoanion can indeed
chemically reduce O2, itself being reoxidized to quinone again.
The increased oxygen reduction current (Figure 4A, at above
2.3 V) is hence attributed to the catalysis of O2 by the
electrochemically generated monoanion.
LiQO2(sol) and two LiQO2(sol) then disproportionate to form
Li2O2 (reactions 2 and 3) or LiQO2(sol) reacts with LiQ• (sol) to
form Li2O2 (reaction 4). The other path involves the
•
dissociation of LiQO2(sol) to form LiO2 (sol) and Q(sol); the
former further decomposing to form Li2O2 (reactions 5 and 6).
In a previous study, Gao et al.47 proposed a very similar
mechanism as described in reactions 1−4. It was further
suggested that the reactive LiO2 species was circumvented, that
is, LiQO2(sol) does not dissociate into LiO2 and Q, which was in
part supported by the observation of a small reduction in the
concentration of parasitic reaction products in their batteries:
4% and 5% fewer side reaction products were seen in DME and
TEGDME electrolytes, respectively, with DBBQ than without
it. In our experiments (Figure 5C), the quantities of
degradation products formed in the DME electrolytes with or
without DBBQ were extremely similar so that the small
difference is within the error limit of our quantification
experiments. Given that no substantial reduction of the side
reactions was seen with the use of DBBQ, we suggest that
reactions 4 and 5 are likely to occur in addition to the
disproportionation (reaction 3). In other words, LiO2 still exists
as an intermediate species, being responsible for a significant
amount of side reactions observed during discharge. In the
anhydrous electrolyte with DBBQ, LiQO2(sol) complex
formation and the high solubility of LiQ•(sol), to some extent,
help to generate/transport reduced oxygen species (LiO2 and
Li2O2) further away from the surface, promoting the formation
of thicker layers of Li2O2 as compared to the case without
DBBQ, as illustrated in Figure 4F−I. From the SEM images
(Figure 4H,I), the discharge process is still largely confined to
regions near electrode surfaces (within 200 nm). This
observation supports the view that LiQ(sol) is chemically
oxidized by nearby O2(sol) before it diffuses further into the
The possible elementary steps are summarized in reactions
1−6.
Q (sol) + Li+ + e− → LiQ•
(sol)
(1)
(2)
(3)
(4)
LiQ•(sol) + O2(sol) → LiQO2(sol)
2LiQO
→ Li2O2(s) + O2(sol) + 2Q (sol)
2(sol)
LiQO
+ LiQ•(sol) → Li2O2(s) + 2Q (sol)
2(sol)
•
LiQO
→ LiO2
+ Q (sol)
(sol)
2(sol)
(5)
(6)
•
2LiO2
→ Li2O2(s) + O2(sol)
(sol)
G
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Article
bulk electrolyte and that the complex LiQO2(sol) and Li2O2(sol)
do not diffuse very far into the electrolyte either.
reduction at 20,000 ppm water to the catalysis of quinone
monoanion.
The potential reaction pathways are summarized in reactions
7−11). The quinone molecule is first electrochemically reduced
to form the lithium quinone monoanion (reaction 7). The
monoanion then chemically reacts with dissolved O2 to form a
water-coordinated monoanion-oxygen complex, LiQO2-
(nH2O)(sol) (reaction 8), where n =1−4 is the number of the
coordinating water with the complex. Subsequently, the
monoanion-oxygen complex disproportionates or reacts with
another quinone monoanion to form Li2O2, reactions 9−11.
Given the fewer side reactions observed in the presence of
water (Figure 5B,C), we tentatively propose that the LiO2
intermediate has been circumvented in the wet case.
Turning to the wet case with 20,000 ppm water, the redox
chemistry of hydroquinones needs to be taken into
consideration, because it is well-known that many hydro-
quinone monoanions are able to reduce oxygen.16,17 Indeed,
millions of tons of H2O2 are produced annually via the
oxidation of hydroantraquinone by O2.52 One thus needs to
consider whether in the presence of such a large amount of
H2O (H2O to DBBQ molar ratio being equal to 100:1 at
20,000 ppm), the ORR in the current system is actually
mediated by the equivalent hydroquinone instead of DBBQ.
This possibility was ruled out by performing CV measurements
of the 2,5-di-tert-butyl-hydroquinone in the same water with
added DME electrolyte (Figure S7), which showed a redox
potential at around 3.5 V, considerably higher than the redox
process seen in Figure 4. Clearly the hydroquinone is not
formed under the present conditions, and it is not responsible
for the oxygen chemistry observed in this work.
Q (sol) + Li+(sol) + e− + nH2O → LiQ• − (nH2O)(sol)
(7)
LiQ• − (nH2O)(sol) + O2(sol) → LiQO − (nH2O)(sol)
2
(8)
Indeed, Gupta and Linschitz demonstrated that the
mechanism of reduction of quinones in wet organic solvents
does not involve protonation of the quinone monoanion or
dianion.19 Protonation of the reduced quinones could be ruled
out in their work, because it was found that their pKa’s were
much lower than those of water or alcohols, meaning that it was
much easier to deprotonate reduced quinones than it is to
extract a proton from water or alcohols. Therefore, the
extraction of a proton from water by the quinone monoanion
or dianion can be ruled out on thermodynamic grounds, that is,
Q−• + H2O = HQ• + OH− is thermodynamically unfavorable,
whereas stabilization of the quinone monoanion by hydrogen
bonding with water is thermodynamically favorable, again as
supported by DFT studies of small clusters (Figure 3).
In the presence of more than 20,000 ppm water (Figures 1
and 4), both the benzoquinone monoanion and dianion can be
formed during the galvanostatic discharge at 2.3−2.8 V, and
CV/DEMS/pressure measurements (Figures 4, 5, and S4)
show that the reduced quinone definitely mediated the two-
electron reduction reaction. To verify its reactivity with O2, the
dianion species generated in either an anhydrous or wet
electrolyte were investigated with UV−vis spectroscopy after
being exposed to an excess of O2. As shown in Figure 7,
although the dianion under anhydrous conditions can readily
reduce O2 (A), the reaction of quinone dianion with O2 in the
presence of 20,000 ppm water is sluggish. There was no
appreciable change in the spectrum within 1 h following the O2
exposure, and there was still around one-third of the quinone
dianion unreacted (Figure 7B) even after 50 h. This
observation is inconsistent with the fast DEMS response that
supports two-electron ORR activity in Figures 5E and S4,
suggesting that oxygen reduction mediated by the quinone
dianion is unlikely to be the dominant process on discharging.
On the other hand, the quinone monoanion was found to react
readily with oxygen in the presence of 4000−40,000 ppm water.
Figure 7C shows that the characteristic absorption bands
between 300 and 450 nm associated with the quinone
monoanion disappeared after exposing the sample to O2, and
the resulting spectra support that the oxidized quinone state, Q,
is formed. In situ Raman measurements further suggest that in
the presence of 20,000 ppm water, the quinone monoanion is
indeed an intermediate species during the reduction to form the
dianion (Figure S8). We therefore ascribe the enhanced O2
2LiQO − (nH2O)(sol) → Li2O2(sol) + 2Q + 2nH2O + O2
2
(9)
LiQO − (nH2O)(sol) + LiQ• − (nH2O)(sol)
2
→ Li2O2(sol) + 2Q + 2nH2O + O2
(10)
Li2O2(sol) → Li2O2(s)
(11)
Despite the fact that the quinone dianion reacts slowly
toward O2 reduction, its formation is not detrimental for the
battery operation, because it can react with the neutral quinone
via the following comproportionation reaction 1253,54 to form
the monoanion that readily reduces O2. This hypothesis is
consistent with the observation (Figure S4) that the battery can
operate via two electrons per reduced O2 at high rates.
Q (sol) + Li2Q − (nH2O)(sol) → 2LiQ• − (nH2O)(sol)
(12)
From the SEM images (Figures 4J,K and S2), the Li2O2
crystals grow as large as 30 μm, suggesting that reduced oxygen
species, Li2O2(sol) and/or LiQO2-(nH2O)(sol), were generated at
a location or had diffused by a distance up to 30 μm away from
the electrode surface, much longer than that in the absence of
water. The final product Li2O2 is much more soluble in water
than in ether, the added water thus helping Li2O2 to diffuse
further away from the electrode surfaces. However, cells
discharged with 20,000 ppm added water but no DBBQ
show that Li2O2 particles can only grow up to 3 μm at the end
of the discharge (Figure S9), which implies that the water
solvation alone cannot explain the phenomena seen with both
DBBQ and water added. Given that added water tends to slow
down the reaction of reduced quinone with O2 (e.g., for the
dianion), it is possible that in a O2 saturated electrolyte, the
water-coordinated monoanion can diffuse for a longer distance
away from the electrode surfaces before it reduces O2
(compared to the anhydrous case), effectively enlarging the
reaction zone in the electrolyte and allowing reactions to occur
deeper into the electrolyte. In addition, the water induced
hydrogen-bond formation and Li+ solvation via the quinone
monoanion can help increase the concentration and diffusion
length of LiQO2-(nH2O)(sol) in the electrolyte. All of these
H
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Journal of the American Chemical Society
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ASSOCIATED CONTENT
* Supporting Information
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/jacs.7b11007.
■
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Materials and methods and figures S1−S10 (PDF)
AUTHOR INFORMATION
Corresponding Author
*cpg27@cam.ac.uk
ORCID
Tao Liu: 0000-0002-6515-0427
Gunwoo Kim: 0000-0001-9153-3141
Nicolas Dubouis: 0000-0001-8698-7202
Nuria Garcia-Araez: 0000-0001-9095-2379
Clare P. Grey: 0000-0001-5572-192X
■
Present Addresses
∥Chimie du Solide et de I’Energie, College de France, UMR
8260, 75231, Paris Cedex 05, France
⊥Forschungszentrum Juelich GmbH, Institute of Energy and
Climate Research Fundamental Electrochemistry (IEK-9)
52425, Juelich and Max-Planck-Institute for Chemical Energy
Conversion, Stiftstrasse 34−36, 45470 Muelheim an der Ruhr,
Germany
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Energy 2016, 1, 16128.
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P. C.; Praveen, B. B.; Dholakia, K.; Tarascon, J. M.; Bruce, P. G. Nat.
Chem. 2014, 6, 1091.
Notes
The authors declare no competing financial interest.
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Viswanathan, V.; Luntz, A. C. Nat. Chem. 2015, 7, 50.
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M.; Kim, G.; Grey, C. P. Science 2015, 350, 530.
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ACKNOWLEDGMENTS
■
The authors thank EPSRC-EP/M009521/1 (T.L., G.K.,
C.P.G.), Innovate UK (T.L.), Darwin Schlumberger Fellowship
(T.L.), EU Horizon 2020 GrapheneCore1-No.696656 (G.K.,
C.P.G.), EPSRC - EP/N024303/1, EP/L019469/1 (N.G.-A.,
J.T.F.), Royal Society - RG130523 (N.G.-A.), and the European
Commission FP7-MC−CIG Funlab, 630162 (N.G.-A.) for
research funding.
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