Download Investigating the kinetic mechanisms of the oxygen

Universidade de São Paulo
Biblioteca Digital da Produção Intelectual - BDPI
Departamento de Física e Ciências Materiais - IFSC/FCM
Artigos e Materiais de Revistas Científicas - IFSC/FCM
2014-08
Investigating the kinetic mechanisms of the
oxygen reduction reaction in a nonaqueous
solvent
Journal of Physical Chemistry C, Washington, DC : American Chemical Society - ACS, v. 118, n. 38,
p. 21995-22002, Aug. 2014
http://www.producao.usp.br/handle/BDPI/50697
Downloaded from: Biblioteca Digital da Produção Intelectual - BDPI, Universidade de São Paulo
Article
pubs.acs.org/JPCC
Investigating the Kinetic Mechanisms of the Oxygen Reduction
Reaction in a Nonaqueous Solvent
Nelson A. Galiote,† Dayse C. de Azevedo,‡ Osvaldo N. Oliveira, Jr.,‡ and Fritz Huguenin*,†
†
Departamento de Química, Faculdade de Filosofia, Ciências e Letras de Ribeirão Preto, Universidade de São Paulo, 14040-901
Ribeirão Preto (SP), Brazil
‡
Instituto de Física de São Carlos, Universidade de São Paulo, CP 369, 13560-970 São Carlos (SP), Brazil
S Supporting Information
*
ABSTRACT: The high theoretical energy density of lithium−oxygen batteries brings
the promise of higher performance than existing batteries, but several technological
problems must be addressed before actual applications are made possible. Among the
difficulties to be faced is the slow oxygen reduction reaction (ORR), which requires a
suitable choice of catalysts and electrolytic solution. This can only be achieved if the
kinetics and mechanism of this reaction are known in detail. In this study, we
determined the rate constants for each elementary step of ORR for a platinum electrode
in 0.1 mol·L−1 LiClO4/1,2-dimethoxyethane (DME), using a kinetic model in the
frequency domain. We found that the energy storage capacity of lithium−air batteries
can be increased by converting a large amount of lithium superoxide into lithium
peroxide during the electrochemical step in comparison with chemical disproportionation. The mechanisms for ORR were supported by data from an electrochemical quartz
crystal microbalance (EQCM): ORR could be distinguished from parasitic reactions
induced by solvent degradation, and agglomerates of LixO2 (1 ≤ x ≤ 2) were adsorbed on the electrode. The rate-limiting step
for ORR was the electron transfer to the oxygen molecules strongly adsorbed onto platinum sites, particularly as a large amount
of reaction product (Li2O2) adsorbed onto the electrode. Even though Pt sheets are likely to be impracticable for real
applications due to their low surface area, they were useful in making it possible to determine the kinetics of ORR steps. This can
now be employed to devise more involved electrodes, such as those containing dispersed Pt nanoparticles.
■
transfer.14−16 Materials like Pd and Pt appear to exhibit this
intermediate adsorption strength.14 Even though platinum is
expensive as electrocatalyst in energy sources and batteries, it is
still the best ORR catalyst in aqueous media and one of the
most efficient in aprotic media.15,17−20 Therefore, studying the
kinetics in these electrochemical systems remains relevant.
ORR has been studied in aprotic media where the superoxide
anion arises.21−24 ORR has been suggested to occur via a oneelectron, two-electron, or four-electron mechanisms in nonaqueous systems, to afford LiO2, Li2O2, and Li2O, respectively,
according to the global reactions:25,26
INTRODUCTION
Population growth has pushed the energetic demand and the
global dependence on oil, with the ensuing rise in emission of
pollutant gases. The use of suitable devices to store the excess
energy originating from “green sources” or a convenient
substitute for combustion engines in transportation is a possible
solution.1−3 However, huge challenges must be addressed
before these technologies become successful. Among the few
alternatives is the lithium−air battery, which has high
theoretical specific energy density (11.7 kW·h·kg−1),4−7 with
a practical specific energy density close to 3.5 kW·h·kg−1 being
reported.8 These batteries have been studied extensively since
the report on the first nonaqueous secondary battery.9 Catalysts
in the positive electrode can enhance the oxygen reduction
reaction (ORR) rate to deliver and store energy efficiently.
Several materials may be applied for this purpose, including
carbonaceous materials (e.g., graphene sheets), transition-metal
oxides (manganese spinel and cobalt oxides), inorganic−
organic composites (metal macrocycle derivatives), metal
oxides (Fe2O3, Fe3O4, and CuO), noble metals, and metallic
alloys (Pt, Pd, and Au).10−13 Electrocatalysts should allow for
suitable oxygen adsorption at the electrode surface, which must
be strong enough to decrease the activation energy barrier for
electro-reduction while being sufficiently weak to allow these
anions to bind with lithium ions after the first electron
© 2014 American Chemical Society
Li+ + e− + O2 → LiO2 ,
E° = 3.00 V (vs Li/Li+)
2Li+ + 2e− + O2 → Li 2O2 ,
(1)
E° = 2.96 V (vs Li/Li+)
(2)
4Li+ + 4e− + O2 → 2Li 2O,
E° = 2.91 V (vs Li/Li+)
(3)
In practical systems, additional electroreduction steps occur
at the positive electrode, with intermediates emerging and
Received: May 30, 2014
Revised: August 27, 2014
Published: August 29, 2014
21995
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
decomposing, and competition with solvent degradation also
exists. The formation of Li2O is kinetically limited, and the
major discharge product is Li2O2, even at standard potentials
more negative than 2.91 V. In theoretical and experimental
studies using density functional theory (DFT) and spectroscopic measurements, respectively, Li2O2 production was
reported. According to these studies, one-electron transfer to
the O2 molecule initiates ORR, thus leading to the superoxide
anion (O2−) being weakly adsorbed on the electrode
surface.4,27−30 Solvent molecules can solvate this adsorbed
species, which can also react with cations such as Li+ in the
electrolytic solution.31 Several organic solvents undergo
nucleophilic attack by superoxide ions, yielding discharge
products that are different from the expected Li x O 2
species.32−36 Other possible solvent degradation reactions are
proton abstraction by the superoxide and autoxidation in the
presence of the solution saturated with oxygen at high
potentials.37,38 These issues limit the practical potential
window, making the choice of a suitable solvent for ORR
crucial.32−34,37,38 Several degradation products (as lithium
carbonate, acetates, formates, and hydroperoxides) have been
detected for ORR in ether solutions,27,39 but a theoretical
approach to predict solvent stability has suggested the use of
ethers as aprotic solvent for lithium−air batteries.36,37
Since ORR in nonaqueous solvents and lithium salts is not
yet fully understood, here we investigated the kinetics of ORR
using a LiClO4/1,2-dimethoxyethane (DME) electrolyte
solution and a platinum electrode. Platinum sheets are not
suitable for batteries due to their surface area, but they are
suitable to study ORR mechanisms in detail since effects from
other variables, such as size distribution in nanoparticles and
contribution from more than one catalyst in alloys and
composites, are eliminated. To estimate the kinetic constants
of the reaction steps, we conducted measurements in the
frequency domain and developed a kinetic model.
+
− k4
Pt−LiO2(ads) + Li + e → Pt−Li 2O2(ads)
k5′
v2 = k 2θO2 exp( −bE 0) exp(bE)
K1K5[O2 ] + 1 + K5
(15)
K5K1[O2 ](1 − θLiO2)
K1K5[O2 ] + 1 + K5
⎛1 ⎞
⎜ ⎟
K1K5[O2 ] + 1 + K5 ⎝ K5 ⎠
(16)
K5(1 − θLiO2)
(17)
where K5 = (k5/k5′ ).
Equation 18 gives the temporal variation of θLiO2 which goes
to zero in the steady-state, necessary to relate the reaction steps
and θLiO2:
dθLiO2
dt
= v2 − v3 − v4 = 0
(18)
In the transient state during the perturbation ac potential,
Taylor expansion helps to express the step rates as42,43
⎛ ∂v ⎞
⎛ ∂v ⎞
2 ⎟
ΔθLiO2e(jωt ) + ⎜ 2 ⎟ ΔEe(jωt )
v2 = v20 + ⎜⎜
⎟
⎝ ∂E ⎠θ
⎝ ∂θLiO2 ⎠ E
LiO2
(19)
⎛ ∂v ⎞
3 ⎟
(jωt )
v3 = v30 + ⎜⎜
⎟ ΔθLiO2e
∂
θ
⎝ LiO2 ⎠ E
(5)
(6)
(20)
⎛ ∂v ⎞
⎛ ∂v ⎞
v4 = v40 + ⎜⎜ 4 ⎟⎟ ΔθLiO2e(jωt ) + ⎜ 4 ⎟ ΔEe(jωt )
⎝ ∂E ⎠θ
∂
θ
⎝ LiO2 ⎠ E
LiO2
(7)
(21)
where j = (−1)1/2, ω is the angular frequency, and v0i is the
steady-state rate for reaction i. The terms ΔθLiO2 and ΔE are
associated with the oscillation amplitude of the coverage degree
and with the sinusoidal ac potential. They obey the following
relation:
(8)
where the subscript “ads” means species adsorbed onto the
platinum sites. The electrochemical steps have been considered
out of equilibrium due to overpotential. On the basis of these
reaction steps, the corresponding rates are obtained:
v1 = k1θPt[O2 ] − k1′θO2
(14)
K5(1 − θLiO2)
θLi 2O2 =
k5
Pt−Li 2O2(ads) ⇄ Pt + Li 2O2
(13)
θO2 =
k3
k 3′
v5 = k5θLi 2O2 − k5′θPt
θPt =
(4)
2Pt−LiO2(ads) ⇄ Pt 2−Li 2O2(ads) + O2
(12)
θPt + θO2 + θLiO2 + θLi 2O2 = 1
K1
Pt−O2(ads) + Li+ + e → Pt−LiO2(ads)
v4 = k4θLiO2 exp( −bE 0) exp(bE)
where b = (αnF/RT), E is the equilibrium potential, and E is
the interfacial potential. θPt is the surface coverage degree
associated with the platinum free sites, and θO2,θLiO2, and θLi2O2
are the surface coverage degrees for O2, LiO2, and Li2O2,
respectively. The term [O2] represents the oxygen molar
concentration dissolved in DME. θPt, θO2 and θLi2O2 can be
related to the surface coverage degree of the adsorbed
intermediate θLiO2 (eqs 14−17) through the mass balance,
which makes its substitution in eq 9−13 possible.
■
− k2
(11)
0
KINETIC MODEL
The elementary steps are based on the ORR mechanism
proposed in the literature.25,31,40,41 The exception is the last
chemical equilibrium, which considers partial dissolution of
Li2O2 adsorbed onto the platinum electrode in DME:
Pt + O2 ⇄ Pt−O2(ads)
v3 = k 3(θLiO2)2 − k 3′θLi 2O2[O2 ]
0
θLiO2 = θLiO
+ ΔθLiO2 exp(jωt )
2
(9)
(22)
θ0LiO2is
where
the coverage degree of adsorbed LiO2 at the
steady-state. The temporal variation of θLiO2 at the transient
(10)
21996
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
state allows one to associate ΔθLiO2 with ΔE as expressed
below:
Γ
dθLiO2
dt
= ΓFjωΔθLiO2 exp(jωt ) = v2 − v3 − v4
Equation 25 shows the total impedance (ZT), which depends
on the capacitive reactance (Zc) associated with the doublelayer capacitance (Cdl) in parallel with the faradaic impedance
(23)
where Γ is the maximum surface coverage. Hence, eqs 19 and
21 provide the oscillatory faradaic current density (Δif), which
is the sum of faradaic contributions Δν2 and Δν4 (Δν2 = ν2 −
ν20 and Δν4 = ν4 − ν40) multiplied by the Faraday constant, to
determine the impedance values for each angular frequency.
ΔE
Zf =
Δi f
Zf − 1
and in series with the electrolytic solution resistance (Rs):
⎛ Zf Zc ⎞
⎟⎟ + R s
ZT = ⎜⎜
⎝ Zf + Zc ⎠
(24)
(−
=−
k 2K5K1[O2 ]ebE
K1K5[O2 ] + 1 + K5
k 2K5[O2 ]K1ebE
K1K5[O2 ] + 1 + K5
+ k4ebE
)⎛⎝−
⎜
+ 2k 3θLiO2 +
where
k 2K5K1[O2 ](1 − θ LiO2)bebE
K1K5[O2 ] + 1 + K5
k 3′[O2 ]
K1K5[O2 ] + 1 + K5
(25)
⎞
− k4θLiO2bebE⎟
⎠
+ k4e
bE
+ Iω ΓF
+
k 2K5K1[O2](1 − θLiO2)bebE
K1K5[O2 ] + 1 + K5
− k4θLiO2bebE
(26)
Equation 26 was used to fit the electrochemical impedance
spectroscopy data. The rate constants were scanned; the
selected ones were based on the lowest difference between the
experimental and theoretical impedance modulus.43 The
methodology used to determine the errors for each rate
constant is described in the Supporting Information.
electrochemically active area of 0.361 cm2 and an integral
sensitivity constant of 0.0815 cm2 s−1 g−1. The resonance
frequency shift was measured with a HP-5370B Universal
Timer/Counter. Changes in the resonance frequency of the
EQCM crystal were transformed into mass changes using the
Sauerbrey equation.46
EXPERIMENTAL SECTION
All the reagents were purchased from Sigma-Aldrich and used
as received: 1,2-dimethoxyethane containing molecular sieves
(with water content lower than 0.005%); lithium perchlorate
salt grade (99.99% purity); and tetrabutylammonium perchlorate (TBAClO4, purity higher than 99.0%). The electrolyte
solutions were prepared prior to use and stored in a glovebox
(MBRAUN) with water and oxygen concentration lower than
10 ppm.
A platinum sheet was used as the working and counter
electrode, and an Ag/Ag+ saturated in electrolyte solution was
employed as a quasi-reference electrode. This latter electrode
had a potential of 2.94 V versus Li/Li+, so all the potentials
referred to hereafter are based on the reference Li/Li+. A
LiClO4 (DME) or TBAClO4/DME electrolytic solution (0.1
mol·L−1) was employed in the electrochemical experiments. All
the experiments were carried out using an Autolab PGSTAT30
potentiostat/galvanostat. Ac electrochemical impedance spectroscopy was employed between 10 kHz and 10 mHz. Several
dc potentials were tried using 5 mV of superimposed ac
amplitude. This impedance was measured after pretreatment
for at least 2000 s, to achieve a pseudosteady state. Here, these
states were defined as a 10% change in current density within
half an hour. Ultrapure N2 (99.999% purity) or O2 (99.998%
purity) from Linde-AGA were purged into the solution for 20
min before the experiments; the gas flow was maintained
during the electrochemical measurements. Oxygen solubility in
DME is 9.57 mmol.cm−3.26 Cyclic voltammetry was performed
in acid medium (0.5 M H2SO4) to determine the value of Γ
(=2.46 × 10−9 mol of sites per cm2), considering an oxidation
charge associated with atomic hydrogen of 210 μA per square
centimeter of real area for polycrystalline platinum.44,45 The
program Maple v13.0 was used to fit the experimental data. For
EQCM experiments, the substrates were 6 MHz AT-cut quartz
crystals coated with platinum. They have a piezoelectrically and
RESULTS AND DISCUSSION
Cyclic voltammetry experiments in an electrochemical cell
using electrolytic solution deaerated with nitrogen gas and
saturated with oxygen gas indicate that the parasitic reaction in
the electrolytic solution in the absence of oxygen for the
potential window used in this work can be neglected, as shown
in Figure 1. These data show that the reduction process in 0.1
■
■
Figure 1. Cyclic voltammograms using electrolytic solution (− −)
deaerated with nitrogen gas and () saturated with oxygen gas at 20
mV s−1. Inset Figure shows the third (red -•-) and fifth (blue line)
voltammetric cycles.
M LiClO4/DME saturated with O2 happens at approximately
2.7 V, as expected for ORR in ether/lithium ion based
electrolytes.14,47−49 The electrooxidation process takes place at
potentials more positive than 3.5 V, confirming the reaction
irreversibility in the conditions used, with low cyclability of the
electrochemical system being illustrated in the third and fifth
voltammetric cyles in the inset in Figure 1. This is a
consequence of the insolubility of the reaction products in
21997
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
DME, which tends to passivate the electrode surface with
increasing number of voltammetric cycles. The potentiodynamic profile of the potential changed along the cycles,
probably owing to other reaction pathways originating from
alterations in the chemical environment due to partial
passivation of the platinum surface.35,47,50,51
The potential, charge (q), and mass change (Δm) measured
with an electrochemical quartz crystal microbalance (EQCM)
during the linear negative potential scan at 20 mV s−1 and in 0.1
M LiClO4/DME saturated with O2 are shown in Figure 2. This
Figure 3. Chronoamperometric curves at 2.34, 2.10, and 1.90 V, using
electrolytic solution saturated with oxygen gas.
to minimize the effects of these oscillations, which last about
100 s. Moreover, the impedance data did not change after
several consecutive measurements in this frequency range.
DME, molecular oxygen, and platinum catalysts participate in
these current oscillations, because the latter do not occur
during chronoamperometric experiments performed with other
electrodes and with aerated or deaerated solvents (for instance,
propylene carbonate or glass carbon electrode). Studies about
these oscillations will be presented elsewhere.
Figure 4 shows the Nyquist and Bode diagrams (in the inset)
obtained in oxygen saturated electrolytic solution containing
0.1 M LiClO4/DME, as well as in oxygen-saturated 0.1 M
TBAClO4/DME, at (a) 2.34, (b) 2.10, and (c) 1.90 V. The
diagrams are superimposed with the ac potential perturbation
of 5 mV from 10 kHz to 10 mHz. With extrapolation of the
semicircle at 2.34 V, the charge transfer resistance (Rct) is
determined as 155 kΩ·cm2 in the presence of oxygen dissolved
in 0.1 M LiClO4/DME. Rct decreases as the dc potential
becomes less positive, and two semicircles emerge at 2.10 and
1.90 V. The semicircles at high and low frequencies are
probably associated with ORR and solvent degradation,
respectively. In the Bode diagrams, two distinct processes are
identified which display different time constants, at less positive
potentials. The Rct values associated with ORR are 20 and 7
kΩ·cm2 at 2.10 and 1.90 V, respectively.
Electrochemical impedance spectroscopy in oxygen-saturated
electrolytic solution consisting of 0.1 M TBAClO4/DME may
confirm ORR and DME degradation reactions in 0.1 M
LiClO4/DME, since tetrabutylammonium cations (TBA+) can
sustain ORR via one-electron transfer without any solvent
degradation26TBA+ ions are poorly solvated and tend to
form a stable TBA−O2 complex.39,55,56 According to the hard−
soft acid−base theory (HSAB), O2− is a soft base with higher
affinity for TBA+ (poorly solvated soft acid) than for the Li+ ion
(hard acid).26 Note that a second semicircle (or second time
constant) does not appear in the Nyquist diagrams at 2.34,
2.10, or 1.90 V when 0.1 M TBAClO4/DME is the electrolytic
solution. Therefore, the second semicircle at low frequencies in
0.1 M LiClO4/DME at 2.10 and 1.90 V (Figures 4b and 4c)
refers to solvent degradation by O2− or Li2O2.39,57
Solvent degradation reactions (not included in the reaction
mechanism) at 2.10 and 1.90 V modify coverage degrees of
reaction products and intermediates of ORR at the steady-state,
and this may compromise determination of the rate constants.
However, we opted to consider kinetic constants as estimated
at these dc potentials, because the inclusion of degradation
Figure 2. Bidimensional projections for E − q (a), E − Δm (b), and q
− Δm planes during ORR in 0.1 M LiClO4/DME saturated with O2 at
20 mV s−1.
experiment was aimed at determining the chemical species
taking part in the electrochemical reactions. The mass change is
small at potentials more positive than 2.10 V, but it increases
significantly at more negative potentials owing to solvent
degradation. The slope in the mass change as a function of the
electroreduction charge is ca. 3.4 × 10−4g C−1 at 2.34 V, which
is an intermediate value between those expected for formation
of Li2O2 (2.4 × 10−4 g C−1) and LiO2 (4.0 × 10−4 g C−1). It
seems that agglomerates of Li2O2 and LiO2 were formed, as
predicted using density functional theory (DFT) calculations
and Raman spectroscopy.16,52,53 Differently from results in
aqueous medium, these EQCM data support a nondissociative
adsorption for oxygen molecules, as Li2O is formed during
dissociative adsorption and the ratio for mass change/charge
expected for Li2O (7.8 × 10−5 g C−1) was not observed.
Nondissociative adsorption for oxygen molecules has also been
supported by isotope labeling experiments.54
The chronoamperometric curves at 2.34, 2.10, and 1.90 V in
Figure 3 are evidence for ORR in oxygen-saturated 0.1 M
LiClO4/DME. The current density decreased fast in the
beginning, but the steady state was not reached even after a
long time because the reaction products partially blocked the
platinum surface, as mentioned above. These experiments
helped to determine the time that would be necessary to obtain
pseudosteady states, which is important when conducting
electrochemical impedance spectroscopy at dc potentials. These
data also provide important information about how the current
oscillates during the electrochemical reactions. These oscillations correspond to ca. 5% of the total current density at 2.34
V, and increase to ca. 10% at 2.10 and 1.90 V. Since the kinetic
model above involves the steady state for ORR, the impedance
data had to be fitted at high frequencies (from 2 kHz to 3 Hz)
21998
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
Figure 4. Nyquist diagrams in oxygen saturated electrolytic solution
containing (○) 0.1 M LiClO4/DME and (•) 0.1 M TBAClO4/DME
at (a) 2.34 V, (b) 2.10 V, and (c) 1.90 V. Bode diagrams are shown in
the inset: (blue ○) 0.1 M LiClO4/DME and (red ●) 0.1 M
TBAClO4/DME.
Figure 5. (•) Experimental and (○) theoretical impedance data for
ORR at (a) 2.34, (b) 2.10, and (c) 1.90 V. () Theoretical
impedance data using CPE. Complex diagrams of capacitance are
shown in the inset figures: (red ●) Experimental and (blue ○)
theoretical impedance data.
reaction in the reaction mechanisms would increase the number
of variables (degrees of freedom), thus turning the rate
constants unreliable. Moreover, we did not observe solvent
degradation for measurements at 2.34 V, although parasitic
reactions were reported for open circuit potential in the
presence of molecular oxygen.54 Therefore, the coverage degree
for the chemical species in the proposed reaction mechanism
can be taken as (almost) independent of parasitic reactions.
Figure 5a through 5c illustrate the fitting of the experimental
impedance data from 2 kHz to 5 Hz at (a) 2.34 V, (b) 2.10 V,
and (c) 1.90 V, using eq 26. First, we fitted the impedance data
using a double-layer capacitance of 40 μF.cm−2 at 2.34 V and of
20 μF.cm−2 at 2.10 and 1.90 V. Table 1 lists the rate constants
and their errors, determined from parabolic fitting around the
values of each kinetic constant (see Supporting Information S1,
S2, and S3). For the sake of completeness in analyzing the
impedance data, we fitted them using a constant-phase element
(CPE) with the kinetic constants of Table 1, since frequency
dispersions associated with the nonfaradaic processes are
observed and are not included in the kinetic model. The
impedance of 1/Q(jω)n replaced the double-layer capacitance,
where Q is a constant with dimension Fsn‑1. The insets in Figure
5 show these fittings in the complex capacitance diagram, to
allow for better visualization at higher frequencies. Moreover,
the beginning of a second semicircle associated with other
processes (solvent degradation reactions) arises at low
frequencies, but we have not fitted it because it is not part of
our kinetic model. The most significant differences (at most
four times) between rate constants at 2.34 V and those
estimated at less positive potentials may be ascribed to solvent
degradation reactions (as mentioned above), in addition to
changes in the chemical environment due to formation of
insoluble Li2O2 product.
21999
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
Table 1. Kinetic and Equilibrium Constants as a Function of Potential
constants
K1
k2,
k3,
k′3,
k4,
K5
mol cm−2·s−1
mol cm−2·s−1
cm·s−1
mol cm−2·s−1
2.34 V
2.10 V
1.90 V
(2.30 ± 0.13) × 102
(1.52 ± 0.15) × 10−6
4.33 ± 0.38
2.54 ± 0.02
(2.36 ± 0.10) × 102
1.90 ± 0.02
(2.18 ± 0.01) × 102
(1.42 ± 0.3) × 10−6
4.760 ± 0.024
0.620 ± 0.007
(9.00 ± 0.007) × 102
7.40 ± 0.01
(2.36 ± 0.08) × 102
(1.59 ± 0.5) × 10−6
9.76 ± 0.10
0.70 ± 0.11
(2.45 ± 0.49) × 102
7.37 ± 0.02
Notes
These kinetic constants suggest that electron transfer to the
oxygen molecules adsorbed onto the platinum sites (eq 5) is
the rate-limiting step. The k2 values are close to those reported
in the literature;24,25 the difference may lie in the solvent
employed and in kinetic model adopted. K1 values are high
because molecular oxygen strongly adsorbs onto platinum
sites.14,16,20 The low Li2O2 solubility in organic medium and the
K5 values in Table 1 suggest that Li2O2 formed during ORR
remained almost totally adsorbed onto the platinum sites,
contributing to electrode surface passivation. This conclusion is
consistent with the decreased current density during the
chronoamperommetric experiments and as a function of the
number of voltammetric cycles, already mentioned. High k4
values indicate larger lithium ion affinity for O22− as compared
with O2−.26,58 The difference between k3 and k4 suggests that
Li2O2 emerges preferentially during the electrochemical step, to
enhance the current density, the energy density and the power
of lithium−air batteries. It is worth pointing out that though the
number of fitting parameters is large, the values inferred from
the fitting were consistent with expectation.
The authors declare no competing financial interest.
■
ACKNOWLEDGMENTS
We are grateful to FAPESP (Projects 2012/21629-1, 2011/
12668-0, and 2011/21545-0), nBioNet (CAPES), and INCTADAPTA/FAPEAM/CNPq (573976/2008-2) for the financial
support. We are also grateful to Prof. Luis Gustavo Dias
(FFCLRP/USP) for the statistical treatment.
■
■
CONCLUSIONS
ORR was found highly irreversible for a platinum electrode in
an electrolytic solution of LiClO4 and DME, according to the
cyclic voltammetry experiments. The low cyclability of this
electrochemical system results from the low solubility of the
reaction products in DME. The EQCM data suggested
formation of Li2O2 and LiO2 during ORR in an organic
medium, whose rate of elementary steps was investigated using
a kinetic model in the frequency domain. The kinetic constants
are consistent with the literature, pointing to high adsorption
energy for oxygen molecules, slow electron transfer to
molecules adsorbed onto the platinum sites, and partial
blockage of these sites due to formation of reaction products
that are insoluble in the organic solvent. Most significantly, the
reaction product (Li2O2) may preferentially arise during the
electrochemical step, thus increasing the number of electrons
transferred to each oxygen molecule adsorbed onto the
platinum sites and enhancing the power and energy storage
capacity of lithium−air batteries. Because these findings are
likely to be independent of the form of Pt, ORR can now be
studied in detail for platinum nanoparticles that are more
suitable for practical devices.
■
ASSOCIATED CONTENT
S Supporting Information
*
Parabolic fittings to determine the error of the rate constants.
This material is available free of charge via the Internet at
http://pubs.acs.org.
■
REFERENCES
(1) Herranz, J.; Garsuch, A.; Gasteiger, H. A. Using Rotating Ring
Disc Electrode Voltammetry to Quantify the Superoxide Radical
Stability of Aprotic Li−Air Battery Electrolytes. J. Phys. Chem. C 2012,
116, 19084−19094.
(2) Wagner, F. T.; Lakshmanan, B.; Mathias, M. F. Electrochemistry
and the Future of the Automobile. J. Phys. Chem. Lett. 2010, 1, 2204−
2219.
(3) Alotto, P.; Guarnieri, M.; Moro, F. Redox Flow Batteries for the
Storage of Renewable Energy: A Review. Renew. Sust. Energy Rev. 2014,
29, 325−335.
(4) McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Girishkumar,
G.; Luntz, A. C. Solvents’ Critical Role in Nonaqueous Lithium−
Oxygen Battery Electrochemistry. J. Phys. Chem. Lett. 2011, 2, 1161−
1166.
(5) Electrochemistry: Rechargeable Li-Air Battery. NatureRes.
Highlights 2012, 488, 8−8.
(6) Scrosati, B. Bottled Lightning: Superbatteries, Electric Cars, and
the New Lithium Economy. Nature 2011, 473, 448−449.
(7) Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J.-M.
Li-O2 and Li-S Batteries with High Energy Storage. Nat. Mater. 2012,
11, 19−29.
(8) Kwabi, D. G.; Ortiz-Vitoriano, N.; Freunberger, S. A.; Chen, Y.;
Imanishi, N.; Bruce, P. G.; Shao-Horn, Y. Materials Challenges in
Rechargeable Lithium-Air Batteries. MRS Bull. 2014, 39, 443−452.
(9) Abraham, K. M. A Polymer Electrolyte-Based Rechargeable
Lithium/Oxygen Battery. J. Electrochem. Soc. 1996, 143, 1−5.
(10) Cheng, F.; Chen, J. Metal-Air Batteries: From Oxygen
Reduction Electrochemistry to Cathode Catalysts. Chem. Soc. Rev.
2012, 41, 2172−2192.
(11) Ren, X. M.; Zhang, S. S.; Tran, D. T.; Read, J. Oxygen
Reduction Reaction Catalyst on Lithium/Air Battery Discharge
Performance. J. Mater. Chem. 2011, 21, 10118−10125.
(12) Débart, A.; Bao, J.; Armstrong, G.; Bruce, P. G. An O2 Cathode
for Rechargeable Lithium Batteries: The Effect of a Catalyst. J. Power
Sources 2007, 174, 1177−1182.
(13) Shao, Y.; Park, S.; Xiao, J.; Zhang, J.-G.; Wang, Y.; Liu, J.
Electrocatalysts for Nonaqueous Lithium−Air Batteries: Status,
Challenges, and Perspective. ACS Catal. 2012, 2, 844−857.
(14) Lu, Y. C.; Gasteiger, H. A.; Shao-Horn, Y. Catalytic Activity
Trends of Oxygen Reduction Reaction for Nonaqueous Li-Air
Batteries. J. Am. Chem. Soc. 2011, 133, 19048−19051.
(15) Lima, F. H. B.; Zhang, J.; Shao, M. H.; Sasaki, K.; Vukmirovic,
M. B.; Ticianelli, E. A.; Adzic, R. R. Catalytic Activity-D-Band Center
Correlation for the O2 Reduction Reaction on Platinum in Alkaline
Solutions. J. Phys. Chem. C 2007, 111, 404−410.
AUTHOR INFORMATION
Corresponding Author
*(F.H.) E-mail: fritz@ffclrp.usp.br.
22000
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
(16) Dathar, G. K. P.; Shelton, W. A.; Xu, Y. Trends in the Catalytic
Activity of Transition Metals for the Oxygen Reduction Reaction by
Lithium. J. Phys. Chem. Lett. 2012, 3, 891−895.
(17) Liang, Y.; Li, Y.; Wang, H.; Dai, H. Strongly Coupled Inorganic/
Nanocarbon Hybrid Materials for Advanced Electrocatalysis. J. Am.
Chem. Soc. 2013, 135, 2013−2036.
(18) Lima, F. H. B.; de Castro, J. F. R.; Santos, L. G. R. A.; Ticianelli,
E. A. Electrocatalysis of Oxygen Reduction on Carbon-Supported Pt−
Co Nanoparticles with Low Pt Content. J. Power Sources 2009, 190,
293−300.
(19) Shao, Y. Y.; Ding, F.; Xiao, J.; Zhang, J.; Xu, W.; Park, S.; Zhang,
J. G.; Wang, Y.; Liu, J. Making Li-Air Batteries Rechargeable: Material
Challenges. Adv. Funct. Mater. 2013, 23, 987−1004.
(20) Lu, Y. C.; Gallant, B. M.; Kwabi, D. G.; Harding, J. R.; Mitchell,
R. R.; Whittingham, M. S.; Shao-Horn, Y. Lithium-Oxygen Batteries:
Bridging Mechanistic Understanding and Battery Performance. Energy
Environ. Sci. 2013, 6, 750−768.
(21) Maricle, D. L.; Hodgson, W. G. Reducion of Oxygen to
Superoxide Anion in Aprotic Solvents. Anal. Chem. 1965, 37, 1562−
1565.
(22) Sawyer, D. T.; Chlericato, G.; Angelis, C. T.; Nanni, E. J.;
Tsuchiya, T. Effects of Media and Electrode Materials on the
Electrochemical Reduction of Dioxygen. Anal. Chem. 1982, 54, 1720−
1724.
(23) Sawyer, D. T.; Valentine, J. S. How Super Is Superoxide? Acc.
Chem. Res. 1981, 14, 393−400.
(24) Chiericato, G.; Rodrigues, S.; Da Silva, L. A. Comparative Study
between Two Methods of Resistance Compensation in the Cyclic
Voltammetric Determination of the Standard Electrochemical Rate
Constant of Dioxygen. Electrochim. Acta 1991, 36, 2113−2117.
(25) Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.;
Hendrickson, M. A. Elucidating the Mechanism of Oxygen Reduction
for Lithium-Air Battery Applications. J. Phys. Chem. C 2009, 113,
20127−20134.
(26) Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.;
Hendrickson, M. A. Influence of Nonaqueous Solvents on the
Electrochemistry of Oxygen in the Rechargeable Lithium−Air Battery.
J. Phys. Chem. C 2010, 114, 9178−9186.
(27) Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwick, L. J.;
Barde, F.; Bruce, P. G. The Lithium-Oxygen Battery with Ether-Based
Electrolytes. Angew. Chem., Int. Ed. 2011, 50, 8609−8613.
(28) Ottakam Thotiyl, M. M.; Freunberger, S. A.; Peng, Z.; Bruce, P.
G. The Carbon Electrode in Nonaqueous Li-O2 Cells. J. Am. Chem.
Soc. 2013, 135, 494−500.
(29) Hummelshoj, J. S.; Luntz, A. C.; Norskov, J. K. Theoretical
Evidence for Low Kinetic Overpotentials in Li-O2 Electrochemistry. J.
Chem. Phys. 2013, 138, 034703.
(30) Viswanathan, V.; Thygesen, K. S.; Hummelshoj, J. S.; Norskov,
J. K.; Girishkumar, G.; McCloskey, B. D.; Luntz, A. C. Electrical
Conductivity in Li2O2 and Its Role in Determining Capacity
Limitations in Non-Aqueous Li-O2 Batteries. J. Chem. Phys. 2011,
135, 214704.
(31) Lu, Y. C.; Gasteiger, H. A.; Crumlin, E.; McGuire, R.; ShaoHorn, Y. Electrocatalytic Activity Studies of Select Metal Surfaces and
Implications in Li-Air Batteries. J. Electrochem. Soc. 2010, 157, A1016−
A1025.
(32) Bryantsev, V. S.; Giordani, V.; Walker, W.; Blanco, M.; Zecevic,
S.; Sasaki, K.; Uddin, J.; Addison, D.; Chase, G. V. Predicting Solvent
Stability in Aprotic Electrolyte Li-Air Batteries: Nucleophilic
Substitution by the Superoxide Anion Radical (O2(*−)). J. Phys.
Chem. A 2011, 115, 12399−12409.
(33) Bryantsev, V. S.; Uddin, J.; Giordani, V.; Walker, W.; Addison,
D.; Chase, G. V. The Identification of Stable Solvents for Nonaqueous
Rechargeable Li-Air Batteries. J. Electrochem. Soc. 2013, 160, A160−
A171.
(34) Bryantsev, V. S. Predicting the Stability of Aprotic Solvents in
Li-Air Batteries: Pk(a) Calculations of Aliphatic C-H Acids in
Dimethyl Sulfoxide. Chem. Phys. Lett. 2013, 558, 42−47.
(35) Freunberger, S. A.; Chen, Y.; Peng, Z.; Griffin, J. M.; Hardwick,
L. J.; Barde, F.; Novak, P.; Bruce, P. G. Reactions in the Rechargeable
Lithium-O2 Battery with Alkyl Carbonate Electrolytes. J. Am. Chem.
Soc. 2011, 133, 8040−8047.
(36) Sahapatsombut, U.; Cheng, H.; Scott, K. Modelling of
Electrolyte Degradation and Cycling Behaviour In a lithium−Air
Battery. J. Power Sources 2013, 243, 409−418.
(37) Bryantsev, V. S.; Faglioni, F. Predicting Autoxidation Stability of
Ether- and Amide-Based Electrolyte Solvents for Li-Air Batteries. J.
Phys. Chem. A 2012, 116, 7128−7138.
(38) Zhu, D.; Zhang, L.; Song, M.; Wang, X.; Mei, J.; Lau, L. W. M.;
Chen, Y. G. Solvent Autoxidation, Electrolyte Decomposition, and
Performance Deterioration of the Aprotic Li-O−2 Battery. J. Solid State
Electrochem. 2013, 17, 2865−2870.
(39) Younesi, R.; Hahlin, M.; Bjorefors, F.; Johansson, P.; Edstrom,
K. Li-O−2 Battery Degradation by Lithium Peroxide (Li2O2): A Model
Study. Chem. Mater. 2013, 25, 77−84.
(40) Girishkumar, G.; McCloskey, B.; Luntz, A. C.; Swanson, S.;
Wilcke, W. Lithium−Air Battery: Promise and Challenges. J. Phys.
Chem. Lett. 2010, 1, 2193−2203.
(41) McCloskey, B. D.; Scheffler, R.; Speidel, A.; Girishkumar, G.;
Luntz, A. C. On the Mechanism of Nonaqueous Li−O2electrochemistry on C and Its Kinetic Overpotentials: Some Implications for
Li−Air Batteries. J. Phys. Chem. C 2012, 116, 23897−23905.
(42) Armstron, Rd; Henderso, M. Impedance Plane Display of a
Reaction with an Adsorbed Intermediate. J. Electroanal. Chem. 1972,
39, 81−90.
(43) Orazem, M. E.; Tribollet, B. Electrochemical Impedance
Spectroscopy; John Wiley & Sons: NJ, 2008; p 533.
(44) Santos, V. P. d.; Tremiliosi Filho, G. Correlaçaõ Entre a
Estrutura Atômica Superficial E O Processo De Adsorçaõ -Dessorçaõ
́ De Hidrogênio Em Eletrodos Monocristalinos Pt(111),
Reversivel
Pt(100) E Pt(110). Quim. Nova 2001, 24, 856−863.
(45) Gasparotto, L. H. S.; Ciapina, E. G.; Cantane, D. A.; Gomes, J.
F.; Lima, F. H. B.; Ticianelli, E. A.; Tremiliosi, G. Altering the
Adsorptive and Electronic Properties of Pt through Poly(Vinyl
Alcohol) Adsorption. Electrochim. Acta 2013, 104, 358−366.
(46) Sauerbrey, G. Verwendung Von Schwingquarzen Zur Wägung
Dünner Schichten Und Zur Mikrowägung. Z. Phys. 1959, 155, 206−
222.
(47) Meini, S.; Piana, M.; Tsiouvaras, N.; Garsuch, A.; Gasteiger, H.
A. The Effect of Water on the Discharge Capacity of a Non-Catalyzed
Carbon Cathode for Li-O2 Batteries. Electrochem. Solid. State Lett.
2012, 15, A45−A48.
(48) Laoire, C. O.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M. A.;
Abraham, K. M. Rechargeable Lithium/Tegdme-Lipf6/O2 Battery. J.
Electrochem. Soc. 2011, 158, A302−A308.
(49) Gallant, B. M.; Kwabi, D. G.; Mitchell, R. R.; Zhou, J.;
Thompson, C. V.; Shao-Horn, Y. Influence of Li2O2 Morphology on
Oxygen Reduction and Evolution Kinetics in Li−O2 Batteries. Energy
Environ. Sci. 2013, 6, 2518−2528.
(50) Aurbach, D.; Daroux, M.; Faguy, P.; Yeager, E. The
Electrochemistry of Noble-Metal Electrodes in Aprotic OrganicSolvents Containing Lithium-Salts. J. Electroanal. Chem. 1991, 297,
225−244.
(51) Cota, L. G.; de la Mora, P. On the Structure of Lithium
Peroxide, Li2O2. Acta Crystallogr., B 2005, 61, 133−136.
(52) Lau, K. C.; Assary, R. S.; Redfern, P.; Greeley, J.; Curtiss, L. A.
Electronic Structure of Lithium Peroxide Clusters and Relevance to
Lithium-Air Batteries. J. Phys. Chem. C 2012, 116, 23890−23896.
(53) Yang, J.; Zhai, D.; Wang, H. H.; Lau, K. C.; Schlueter, J. A.; Du,
P.; Myers, D. J.; Sun, Y. K.; Curtiss, L. A.; Amine, K. Evidence for
Lithium Superoxide-Like Species in the Discharge Product of a Li-O2
Battery. Phys. Chem. Chem. Phys. 2013, 15, 3764−3771.
(54) McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.;
Shelby, R. M.; Luntz, A. C. On the Efficacy of Electrocatalysis in
Nonaqueous Li−O2 Batteries. J. Am. Chem. Soc. 2011, 133, 18038−
18041.
22001
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002
The Journal of Physical Chemistry C
Article
(55) Wang, H.; Xie, K.; Wang, L.; Han, Y. N-Methyl-2-Pyrrolidone as
a Solvent for the Non-Aqueous Electrolyte of Rechargeable Li-Air
Batteries. J. Power Sources 2012, 219, 263−271.
(56) Nasybulin, E.; Xu, W.; Engelhard, M. H.; Nie, Z. M.; Burton, S.
D.; Cosimbescu, L.; Gross, M. E.; Zhang, J. G. Effects of Electrolyte
Salts on the Performance of Li-O‑2 Batteries. J. Phys. Chem. C 2013,
117, 2635−2645.
(57) Assary, R. S.; Lau, K. C.; Amine, K.; Sun, Y. K.; Curtiss, L. A.
Interactions of Dimethoxy Ethane with Li2o2 Clusters and Likely
Decomposition Mechanisms for Li-O‑2 Batteries. J. Phys. Chem. C
2013, 117, 8041−8049.
(58) Zheng, D.; Wang, Q.; Lee, H. S.; Yang, X. Q.; Qu, D. Catalytic
Disproportionation of the Superoxide Intermediate from Electrochemical O2 Reduction in Nonaqueous Electrolytes. Chem.Eur. J.
2013, 19, 8679−83.
22002
dx.doi.org/10.1021/jp5053584 | J. Phys. Chem. C 2014, 118, 21995−22002