Li 2 0 2 in Li-0 2 Batteries: Catalytic Enhancement of Electrochemical
oxidation and Thermophysical Transformations
ARCHAVES
by
MA SSACHUSETTS INSTf
E
OF TECHNOLOGY
Koffi Pierre Yao
B.S. Mechanical Engineering
JUN 2 5 2013
University of Delaware, 2010
LIBRARIES
SUBMITTED TO THE DEPARTMENT OF MECHANICAL ENGINEERING
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF
MASTER OF SCIENCE IN MECHANICAL ENGINEERING
AT THE
MASSACHUSETTS INSTITUTE OF TECHNOLOGY
June 2013
V 2013 Massachusetts Institute of Technology.
All rights reserved.
.
Signature of A uthor......................................
K-i-)
..
. ..................
...... ........
Department of Mechanical Engineering
>.
M
Mo
12012
Z
Certified by...................................................
........
/Yag
Gail E. Kend
A ccepted by.....................................................
Shao-Horn
rofes or of Mechanicl Engineering
A eT iSupervisor
-- ..................
David E. Hardt
Ralph E. and Eloise F. Cross Professor of Mechanical Engineering
Chairman, Department Graduate Committee
1
2
Li 2 0 2 in Li-0
2
Batteries: Catalytic Enhancement of Electrochemical oxidation and
Thermophysical Transformations
By
Koffi Pierre Yao
Submitted to the Department of Mechanical Engineering
on May 10, 2013 in Partial Fulfillment of the
Requirements for the Degree of Master of Science in
Mechanical Engineering
ABSTRACT
Electrification of transportation in the United States is of importance in reducing dependence on
foreign oil and curtailing global warming. However, optimal market penetration of electric
vehicles is confronted with the prohibitive cost and limited energy capacity of current state of the
art lithium-ion battery packs, factors which limit range below 300 miles. Lithium-air (Li-air or
Li-0 2 ) batteries could deliver more than three times the gravimetric energy of Li-ion batteries at
potentially reduced cost by replacing transition metal oxide cathode with formation of lithium
oxides (Li 2 0 2 and Li 20). Being in its infancy, the Li-0 2 technology faces multiple challenges
such as inadequate round trip efficiency (below 80%), low power capability, poor cycle life (less
than 100 cycles) and thermal safety concerns. This thesis is concerned with the poor oxidation
kinetics of the discharge product Li 2 0 2, root cause of poor round trip efficiency, and the thermal
stability of the candidate discharge products Li 20 2 and Li2 0.
Catalysis of the Li 20 2-oxidation by LaCrO 3, Bao. 5 Sro. 5 Coo.8Feo.20
3
(BSCF), LaNiO 3 , LaMnO 3+8,
and LaFeO 3 was systematically investigated. It was found that LaCrO3 , reported with the lowest
activity in aqueous OER, shows a threefold higher activity compared to BSCF, reported with two
orders of magnitude higher activity in aqueous OER. We postulate that efficient catalysts affect
the surface energy landscape of Li 2 0 2 at interfaces to result in larger proportions of low
oxidation-overpotential surface orientations and, therefore, enhanced Li 2 O2 -oxidation at lower
overpotentials. Regarding the thermal stability of Li 2 0 2 and Li 2 0, X-ray diffraction revealed
significant decrease in the c/a ratio of the lattice parameters of Li 2 0 2 from 280 'C to 700 'C,
which are attributed to the transformation of Li 2 0
2
to Li 2 02-6 . Upon further heating, a lithium-
deficient Li 2-6O phase appeared at 300 'C and gradually became stoichiometric upon further
heating to ~550 'C. XPS measurements showed growth of Li2 CO 3 on surfaces of Li2 0 2 and Li 2 0
at 250 'C attributable to chemical reactions between Li 2 0 2 /Li 2 O and carbon-containing species.
The origin of the high activity of LaCrO 3 and experimental understanding of the mechanism of
Li 20 2 oxidation are currently underway.
Thesis Supervisor: Yang Shao-Horn
Title: Gail E. Kendall Professor of Mechanical Engineering
3
4
5
Acknowledgements
I (Koffi Pierre Yao) am indebted to Professor Yang Shao-Horn for the opportunity to
undertake this intellectually challenging study in spite of my limited knowledge of the tools
involved at the time. This work has provided me with a greater understanding of XRD, TGA,
and XPS and also a working knowledge of DFT and its applications.
I would like to thank my father and mother for all the support. Acknowledgements in this
thesis will not do little justice to them, and as such, I elect to keep the extent of my thanks
implied. I am truly grateful to my "big brother" for the endless pep talks.
I am further grateful to Dr. Fanny Barde (Toyota Motor Europe, Belgium) for insightful
discussions and her attention to details. Her probing questions tremendously improved the rigor
of my investigation of the perovskites in Li-0 2 batteries.
I am surely grateful to the MIT electrochemical energy lab members for the tremendous
support, be it by providing scientific information or by inviting me to Friday night beers.
XPS data were obtained in collaboration with David Kwabi (under the funding of a
TOTAL Graduate Student Fellowship) and Yi-Chun Lu in the electrochemical energy laboratory.
DFT simulations were performed by Dr. Yueh-Lin Lee in the MIT electrochemical
energy laboratory.
Processing of XRD was done in collaboration with Dr. Alexis Grimaud in the MIT
electrochemical energy laboratory.
Investigation of the catalytic activity of perovskites in Li-0 2 batteries was supported by
Toyota Motor Europe at Toyota Motor Europe, Research & Development 3, Advanced
Technology 1, Hoge Wei 33 B, B-1930 Zaventem, Belgium.
Investigation of the thermal stability of Li 2 0 2 and Li 2 0 was supported by the Assistant
Secretary for Energy Efficiency and Renewable Energy, Office of FreedomCAR and Vehicle
Technologies of the DOE (DE-AC03-76SF00098 with LBNL), U.S. Department of Energy's
U.S.-China Clean Energy Research Center for Clean Vehicles (Grant DE-PI0000012), Office of
Naval Research (ONR) under contract numbers N00014-12-1-0096 (MIT) and N00014-12WX20818 (NSWCCD), and the Ford-MIT Alliance.
6
Table of Contents
16
1 Introduction and backgrounds........................
16
M otiv atio n .....................................................................................................................
1.1
16
Li-0 2 batteries: advantages and working principles .................................................
1.2
19
C hallenges in Li-0 2 batteries ......................................................................................
1.3
19
1.3.1 Inadequate pow er capability .................................................................................
20
............................................................................................
life
cycle
1.3.2 Inadequate
1.3.3 Inadequate round trip efficiencies: a combination of slow kinetics and parasitic
21
chem istries in Li-0 2 cells.................................................................................................
2 Scope of this thesis.............................................................
2.1
2.2
Thermal Stability Studies of Li 2O2 and Li 20.............................................................
Catalyzing the Li 2 0 2 -oxidation in Li-0 2 ............... ....... ...... . . . . . . . . . . . . . . . .
30
.. .
30
30
3 Catalyzing the Li 20 2-oxidation in Li-0 2 batteries using
33
ABO 3-type perovskites...........................................................
ABO 3 -Perovskites as a self-consistent platform for systematic study of Li 20 2 -oxidation
3.1
33
in L i-0 2 cells .............................................................................................................................
37
Exp erim ental.................................................................................................................
3 .2
38
3.2.1 Synthesis of the perovskite oxides........................................................................
40
3.2.2 Electrode ink preparation......................................................................................
41
3.2.3 E lectrode film casting ............................................................................................
42
3.2.4 Electrochem ical testing..........................................................................................
45
Results and Discussion ............................................................................................
3.3
................
Li202
commercial
3.3.1 Effective electro-oxidation of preloaded
.............. . . 45
45
3.3.2 Electrode background subtraction........................................................................
3.3.3
Catalytic performance of LaCrO 3, LaMnO 346 , LaNiO 3 , Bao.5 Sro. 5 Coo.8 Feo.2 0 3 , and
L aF eO 3 perovskites...............................................................................................................
3.3.4 Strongly diverging activity patterns of perovskites from H2 0-OER to Li 20 2 -OER
3.3.5 Proposed origin of observed divergence between H20-OER to Li 2 0 2 -OER .....
C on clu sion ....................................................................................................................
3.4
49
52
52
56
4 Thermal Stability Studies of Li 2 0 2 and Li 2 0: Combined In
Situ XRD, DFT, and XPS Studies upon Heating ................... 58
Crystal structure of Li 20 2 and Li 20 ..........................................................................
4.1
Previous studies of the thermal transformation of Li 2 02 . . . . . . . . . .
4.2
...........
Exp erim ental.................................................................................................................
4 .3
4.3.1 In situ X -ray diffraction (XRD )............................................................................
4.3.2 In situ X-ray photoelectron spectroscopy (XPS).................................................
4.3.3 Density functional theory (DFT) calculations .....................................................
4.3.4 Thermogravimetric analysis (TGA)......................................................................
Results and discussion ...............................................................................................
4.4
7
58
59
62
62
63
64
68
70
4.4.1 In situ X -ray D iffraction ........................................................................................
4.4.2 In situ XPS analysis of Li 2 0 2 and Li20 .................................................................
C on clu sion s...................................................................................................................
4 .5
5 Perpectives ........................................................................
8
70
78
90
92
List of figures
Figure 1: Estimated gravimetric and volumetric energy density of LiCoO 2 and 02 (Li 2 0 2 ) as the
positive electrode with carbon (C6) or Li as the negative electrode. The cell voltages used
for C-LiCoO 2 , C-Li 2 0 2 , Li-LiCo0 2 , and Li-Li 20
2
are 3.7 V, 2.45 V, 4.0 V, and 2.75 V,
respectively. Neither catalyst, carbon, nor electrolyte were included in the calculation for 02
cells. The positive electrode (LiCoO 2 or Li 2 0 2 ) is assumed to be lithiated and an additional
two times excess lithium is used as the lithium negative electrode. Figure and caption
reproduced from Ref. [6] with permission from The Royal Society of Chemistry........... 17
Figure 2: Ragone plot comparing gravimetric energy and power of Li-ion (LiCoO 2 , LiFePO 4, and
LiNio 5Mno.50 2 and Li-0 2 positive electrodes tested in non-carbonate electrolytes and
normalized to the positive active electrode weight only. Li-0 2 electrodes include Vulcan
Carbon (VC), carbon nanofibers (CNF), Super P carbon, freestanding hierarchically porous
carbon (FHPC graphene), and pristine Nao.44 MnO2 nanowires/Ketjen Black (P-ZMnO2/KB). The Li-0 2 values were normalized to the weight of the electrode in the
discharged state (C + Li 20 2 or C + catalyst + Li 2 0 2 , excluding binder) and were calculated
based on the reported average discharge voltage and total gravimetric capacity (for energy)
or current (power). The upper limit in the gravimetric energy of Li 2 0 2 was calculated
assuming a discharge voltage of 2.75 VU. Figure and caption reproduced from Ref. [6] with
18
permission from The Royal Society of Chemistry. ...........................................................
Figure 3: Schematic of the discharge (left) and charge (right) of a Li-0 2 electrochemical cell... 19
Figure 4: (a) 0 K-edge FY, (b) 0 K-edge TEY, and (c) C K-edge TEY/FY XANES spectra of
electrodes discharged to 1000 (at 250 mA- g~Icabon) and 4700 mAh-g-Icarbon (at 100 mA- g~
Carbon) on the 1st discharge. The reference spectra of commercial Li 2 0 2 (90%, Sigma
Aldrich), commercial Li 2 CO 3 (99%, Alfa Aesar), and pristine VACNTs (FY for C K-edge)
are included. (d) Schematic of discharge products formed at low and high capacity on
VACNTs on the 1st discharge. Figure and caption reproduced from figure 3 of reference
23
[21]; Copyright 2012 American Chemical Society. .........................................................
Figure 5: Composite electrodes (Super P/a-MnO 2/Kynar) that contain the discharge products
individually were subjected to charging in 1 M LiPF 6 in propylene carbonate under 02. (a)
FTIR spectra of the as-prepared electrodes and the charged electrodes for each of the
compounds, together with the spectrum of a pristine electrode. (b) The corresponding
charging curves at 70 mA g~carbon. Since the theoretical capacities of the different
compounds vary, to aid comparison the capacities are all normalized to unity (theoretical
capacities: Li propyl dicarbonate 1000 mAh-gcarbon, 2 e-/mol; Li 2 CO 3 1500 mAh-gcarbon,
2e-/mol; CH 3 CO 2Li 750 mAh-g-carbon, le /mol; HCO 2 Li 750 mA h/g, le/mol). (c-e) MS
gas analysis at the end of charging under 02 of CH 3 CO 2 Li (c), C 3 H 6 (OCO 2 Li)2 (d), and
HCO 2 Li (e). Note that unmarked peaks arise from fragments of C0 2 , H2 0, 02, and Ar. (f)
Gas evolution measured by DEMS on oxidation of a composite electrode containing Li2 CO 3
in response to a stepwise increased current under Ar. Figure and caption reproduced from
Figure 5 of reference [32]; Copyright 2011 American Chemical Society......................... 26
Figure 6: Schematic of molecular orbital splitting of the transition metal 3d band engendered by
hybridization of the metal 3d and oxygen 2p orbitals during adsorption. ........................ 33
Figure 7: The relation between the OER catalytic activity, defined by the overpotentials at 50 tA
cm 2 ox of OER current, and the occupancy of the eg-symmetry electron of the transition
9
metal (B in ABO 3). Data symbols vary with type of B ions (Cr, red; Mn, orange; Fe, beige;
Co, green; Ni, blue; mixed compounds, purple), where x = 0, 0.25, and 0.5 for Fe. Error bars
represent standard deviations of at least three independent measurements. The dashed
volcano lines are shown for guidance only. Figure and caption adapted from reference [53].
34
Reprinted with perm ission from AAAS. ...........................................................................
Figure 8: Variations of charge voltage and gas compositions (helium not included) during the
first charging process for the Li 2 0 2/Fe 3 O4/SP/PVDF electrode in a carbonate electrolyte.
Figure and caption adapted from reference [58], with permission from Elsevier. ........... 37
Figure 9: Phase purity of as-synthesized perovskites investigated by X-ray diffraction. Optimal
purity of each perovskite is observed. Minor impurity phases estimated to less than 1%
(peaks not very visible from scaling) were detected for LaCrO 3 and LaMnO3 ....... ...... . . . 40
Figure 10: Schematic of electrode ink preparation. Component mass ratios were set to
perovskite:vulcan carbon:Li 20 2 :lithiated nafion binder = 3:1:1:1 (carbon content =
41
0.21±0.05 mg cm-2 ......................................................
homogenized
the
to
right,
left
From
synthesis.
ink
following
of
electrodes
Figure 11: Fabrication
ink is drawn on an aluminum foil using a #50 Mayer-rod; upon evaporation of isopropanol,
half-inch diameter electrodes are punched and dried at 70 'C in a Buchi* B-585 oven.
Loading, drying and returning the electrodes to the glovebox using the Buchi* oven
42
prevents exposure of the electrodes to ambient atmosphere.............................................
4
at
and
post-charging
lines)
black
Figure 12: X-Ray diffractions of as-fabricated (Uncharged,
VLi (Charged, red lines) perovskite catalyzed Li 20 2-prefilled electrodes. For all five
perovskite-catalyzed electrodes, the strongest peak of Li 2 0 2 at ~34.97 from the (101)
crystallographic plane disappears after charging, indicating its effective oxidation. Keys: (*)
Li2 0 2 crystal Cu Ka peaks locations; (o) Corresponding perovskite Cu Ka peaks locations.
Residual LiClO 4 salt and aluminum substrate peaks are indicated on the figure. As intended,
47
no catalyst peaks are visible in the Vulcan carbon-only electrode. ..................................
Figure 13: Representative SEM images of pristine (Uncharged) versus post-charging at 4 VLi Of
Ba0 5Sro. 5Coo.8Feo. 2 0 3 and LaCrO3 catalyzed Li 20 2-prefilled electrodes. Key: (o)
corresponding catalyst particles locations; (o) Li 20 2 particles location. After potentiostatic
charging at 4 VLi, no trace of Li 2 O2 particles can be observed by SEM in the charged
electrodes. The same is observed for all other catalyzed electrodes studied within this report.
48
..........................................................
..
(b)
and
BSCF-catalyzed
(a)
on
performed
subtraction
Figure 14: Examples of background
LaCrO 3 -catalyzed electrodes at 4.0 VLi. Little change is observed in the final current (Net
activity of electrode), which highlights the negligible magnitude of parasitic currents
compared to actual Li 20 2 -oxidation currents. Negligible and featureless current curves of
the electrode with no Li 2 0 2 compared to electrode with Li 2 0 2 confirms that the observed
performance of peroxide packed electrodes is due to effective oxidation of Li 2 02. ...... . . 49
Figure 15: Potentiostatic electrochemical performance of perovskite-catalyzed electrodes
(perovskite:Vulcan carbon:Li 20 2 :lithiated nafion = 3:1:1:1 mass ratio) at 4.0 VLi (carbon
content = 0.21±0.05 mg cm-2). (a, b) Mass-specific activity of electrodes with BSCF,
LaNiO 3, LaMnO 3+6 , LaCrO 3 and LaFeO 3 compared to an uncatalyzed Vulcan carbon
electrode (Vulcan carbon:Li 20 2 :lithiated nafion = 1:1:1 mass ratio). An increase in the
50
electrode current output is observed after addition of the perovskites. ...........................
nafion
Figure 16: (a) Catalyst area-specific activity of (Perovskite:Vulcan carbon:Li 20 2 :lithiated
= 3:1:1:1 mass ratio) versus filling of eg* antibonding orbital. (b) Area-specific activity of
10
(Perovskite:Vulcan carbon:Li 20 2 :lithiated nafion = 3:1:1:1 mass ratio) electrodes
normalized to the combined [Carbon+Catalyst]-total surface area. The average area-specific
activity of baseline carbon is added (dotted line). Activities of LaCrO 3 and
Bao 5Sro 5 Coo8Feo.20
3
are well above that of baseline carbon, which proves their activity
cannot be explained by mere surface area effects but rather actual catalysis. Catalytic effects
from LaNiO 3 , LaMnO 3+6 , LaFeO 3 , and La 0 .5Cao. 5FeO 3-6 are less unambiguous, and
considered comparable to carbon, considering experimental errors.................................. 50
Figure 17: (a) Mass-specific activity vs. potential for LaCrO 3, BSCF compared to Au/C, Pt/C,
Ru/C (Pt,Ru,Au:Vulcan carbon:Li 20 2 = 0.66:1:1; Perovskite:Vulcan carbon:Li202 = 3:1:1,
mass ratios) and VC-only reported previously 24 (b): Surface-area specific activity vs.
potential for the same electrodes. Considering that there is ~12 to 20 times more surface
area in the noble metal electrodes (~60 m 2 catalytic/gcarbon) compared to the perovskite
electrodes (3-5 m2catalytic/gcarbon), the present activity of LaCrO 3 and BSCF is close to that of
Pt/C and Ru/C on a surface area basis. Tafel slopes are ~250 mV/decades..................... 51
Figure 18: Linear sweep voltammetry following an hour long discharge at 2.25 V. Figure and
caption adapted from Figure 2e of reference [65]. Copyright 2013 WILEY-VCH Verlag
54
Gm bH & Co. K GaA , W einheim ........................................................................................
Figure 19: (a) Primitive hexagonal unit cell of Li 20 2 (space group P6 3 /mmc 6 6 ). (b) Primitive face
58
centered cubic unit cell of Li 20 (space group Fm3m 68). .....................................................
Figure 20: Atomic supercells of Li 20 2 used in DFT-simulation of Li2 0 2. Oxygen and lithium
ions are represented by red and green spheres respectively. The blue sphere depicts the
65
location of introduced oxygen vacancies during simulation.............................................
Figure 21: Atomic supercells of Li 2 0 used in DFT-simulation. Oxygen and lithium ions are
represented by red and green spheres respectively. The black sphere depicts the location of
66
introduced Li vacancies during sim ulation........................................................................
In
a)
XRD;
in
situ
by
investigated
Figure 22: Temperature dependent phase evolution of Li 20 2
situ scanning at 50'C steps from 25'C to 700 'C and back to 25'C (Coarse XRD); b) In situ
scanning at 10'C steps from 200 'C to 400 *C (Fine XRD). Crystallographic peaks are
70
coded by colors: blue: Li 2 0 2 , red: Li 20 and black: LiOH. ................................................
700*C
to
25"C
from
steps
at
50'C
Li
O
of
XRD
in
situ
coarse
20
range
to
60'
Figure 23: Full 30
2 2
71
an d b ack to 2 5'C ...................................................................................................................
Figure 24: Evolution of species lattice parameters during thermal decomposition of Li 2 0 2 . a) a
and c parameters of Li 2 0 2 and Li 20 during heating from room temperature to 700 'C.
Lattice parameters of Li 2 0 2 between 200 'C and 300 'C (square symbol) are collected in a
separate experiment at 10 'C temperature steps (fine XRD). Notice the lattice shrinkage of
Li 20 2 from its expanded state in the vicinity of the phase transformation temperature
(~280'C). b) Experimental c/a ratios in Li 2 0 2 during thermal treatment compared to DFTsimulated c/a ratios as a function of oxygen vacancies. Trend match in c/a ratios from
experiment and DFT suggests that thermal treatment of Li 2 0 2 results in formation of
oxygen vacancies between 280 'C and 310 *C prior to phase change. Comparison of DFT
and XRD-extracted lattice parameters are strictly interpreted in terms of trends and not
........ .................. 72
quantitatively. . . . .....................................................................................
Figure 25: (a) Calculated Li 2 0 2 total energy vs. unit cell volume with GGA-PBE in this work.
The corresponding bulk modulus and the first pressure derivative of the bulk modulus of
Li 2 0 2 at zero pressure fit with third-order Birch-Murnaghan equations of state are 71.12
and 4.30 GPa respectively. (b) Calculated volumetric thermal expansion coefficients
11
between T=27~327 *C at zero external pressure based on the quasi-harmonic Debye
model. 67 DFT data were obtained in collaboration with Dr. Yueh-Lin Lee in the
73
electrochem ical energy laboratory...................................................................................
Figure 26: Experimental lattice volume change in Li 2 0 2 (referenced to the expanded lattice
volume at 280 'C) during thermal treatment compared to DFT-simulated lattice volume
change as a function of oxygen vacancies. In the same fashion as the experimental and
calculated c/a ratios, a good trend match is found between the experimental and calculated
volume changes that confirm the formation of oxygen vacancies in Li 202 close to the
conversion temperature. DFT data were obtained in collaboration with Dr. Yueh-Lin Lee in
75
the electrochem ical energy laboratory ...............................................................................
Figure 27: Temperature dependent phase and lattice evolution in Li 2 0 investigated by in situ
XRD. a) Li 20 remains stable, experiencing a fairly linear thermal expansion (29.2- 10-6 K-)
from 25 'C to 700 'C (commercial Li 2 0). b) The lattice parameters of Li 20 from phase
transformation of Li 2 0 2 remained noticeably lower than those of commercial Li 20 between
300 0C and 550 'C suggesting a lithium-deficient Li 2-6 0 phase (see Figure 29 for DFT
calculated lattice parameters of Li-deficient Li 2 O) that progressively reaches Li 2 0stoichiometry with increasing temperature. Error bars are the reliability factors generated by
the Fullprof software reflecting uncertainty in the profile fitting...................................... 76
Figure 28: Full 30 to 700 20 range coarse in situ XRD scans of Li 20 at 50'C steps from 25'C to
77
700'C and back to 25'C ...................................................................................................
a
function
Figure 29: Dark blue: ratio of lattice parameters of Li 2 0 formed from Li 2 0 2 (adefect as
of temperature) decomposition to those of commercial Li 2 0 (astoich as a function of
temperature). Light blue: ratio of DFT-simulated lattice parameters of Li-defective LizO to
those of DFT-simulated stoichiometric Li 2 0. A good trend match these ratios is found that
supports the postulate of a lithium deficient phase of Li 20 being formed at initial stages of
decomposition of Li 2 0 2 to Li 20, which gradually becomes stoichiometric. DFT data were
obtained in collaboration with Dr. Yueh-Lin Lee in the electrochemical energy laboratory.
78
..........................................................
Figure 30: XPS spectra of lithium-oxygen compounds used in this study. The dotted red lines
mark the position of the main LiOH peak in Li Is and 0 Is according to literature
references. 6,97 The LiOH material is covered by Li 2CO 3 and therefore appears shifted in the
0 Is and Li Is photoemission regions compared to the literature values. XPS data were
obtained in collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT
80
electrochem ical energy laboratory....................................................................................
Figure 31: Temperature dependent in situ XPS spectra of Li 2 0 2 pellet in (a) C Is, (b) Ols, and
(c) Li Is regions. Atomic percentage of C, 0, and Li in Li 2 0 2 , Li 2 CO 3 and Li 20 obtained
from quantitative component analysis of (d) C Is, (e) Ols, and (f) Li Is spectra. XPS data
were obtained in collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT
81
electrochem ical energy laboratory...................................................................................
Figure 32: Discharge curve for Vulcan carbon electrode discharged at 100mA/gcarbon in 0.1 M
LiClO 4 in DME to -1600 mAh/gcarbon. XPS spectra of this electrode is presented in Figure
33. XRD profile of a similarly discharged vulcan carbon electrode is reported by Lu et al.5
and shows the presence of Li 2 0 2 . XRD does not probe the thin surface Li 2 CO 3 layer present
83
in this discharged electrode...............................................................................................
Figure 33: XPS spectra comparing electrochemically formed Li 2 0 2 (in a discharged carbon Li-0 2
electrode) to Li 2 0 2 heated to 350 *C. Carbonate formation via chemical in presence of
12
carbon-containing species likely contributes to the formation of Li 2 CO 3 . XPS data were
obtained in collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT
84
electrochem ical energy laboratory...................................................................................
(c)
and
Ols,
(b)
Is,
C
(a)
in
Figure 34: Temperature dependent in situ XPS spectra of Li 20 pellet
Li Is regions. Atomic percentage of C, 0, and Li in Li 2 0 2 , Li 2 CO 3 and Li 2 0 obtained from
quantitative component analysis of (d) C Is, (e) 01s, and (f) Li Is spectra. XPS data were
obtained in collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT
85
electrochem ical energy laboratory....................................................................................
290.0
at
Figure 35: Temperature dependent in situ XPS of Li 20 pellet calibrated to Li 2CO 3 peak
eV above 300 *C. XPS data were obtained in collaboration with David Kwabi and Dr. Yi87
Chun Lu in the MIT electrochemical energy laboratory. ..................................................
13
List of Tables
Table 1: Literature values for Li20 2 -oxidation activities under various cell conditions ........... 28
Table 2: SEM calculated particles size and surface areas of the perovskites investigated..... 41
Table 3: Examples of activity reversal patterns from Li 20 2-OER to H2 0-OER found in literature.
53
...............................................................................................................................................
Table 4: Theoretical overpotentials for Li 20 2-oxidation as a function of surface orientation.
54
Table data adapted from reference [64]............................................................................
Table 5: Literature data of Li 2 0 2 phase transformation under various heating conditions ..... 59
Table 6: Lattice vectors of the simulated Li 2 0 2 and Li 2 0 supercells. 51 , i,6, and jp
are the lattice vectors of the Li 20 2 8-atom (Li 2 0 12-atom) primitive cell
represented with the lattice parameters of a = 3.183 A and c = 7.726 A (a(Li2O)= 4.65 A)
-LiO0
-LiO0
i 0
and
cel
cll'
cell'
2
form.
vector
Cartesian
the
in
(a p2 b 2 , and j
2)
20
0
Li
SLi202
64-atom cell'
64-atom cell'
-b
2
64-atom cell
are the lattice vectors of the Li 2 0
-110
110
-D1,0
-D 0
2
32-atom and 64-atom
-D
0
2
the Li20
23
2
2
cell are
aeUeL
cell', 96-atom
b 96-atom
d6-atom cell' , b
c 48-atom
cell and
an 5'2
2
lattice vectors of the 48-atom and 96-atom supercells, respectively. All the supercells are
constructed by linear combination of the primitive cell lattice vectors............................ 67
and 5%'2
Supercells,
a482 atom
suerelsan
cell',
'%
4-ao
cell' ,
14
15
1 Introduction and backgrounds
1.1 Motivation
Curtailing global warming and achieving national energy security by lowering
dependence on foreign oil are pressing challenges for almost all nations today. The United States
import approximately 55% of its crude oil consumption. The transportation sector in the United
States accounts for ~70% of the nation's petroleum consumption and 30% of greenhouse gases
(GHG) emissions as of 2012.2
Therefore, introduction of C0 2-free fully-electric (EV) and
reduced emission hybrid (HEV) vehicles to replace hydrocarbon-burning internal combustion
engines would help reduce the nation dependence on oil import and lower atmospheric CO 2
down to its safe level of 350 ppm.3 However, market penetration of electric vehicles is
confronted with prohibitive costs (greater than $600 kWh' at the pack level) and low energy per
unit weight (~120-200 Wh kg') of current state-of-art lithium-ion batteries resulting in typical
ranges below 300 miles. 4
1.2 Li-0 2 batteries: advantages and working principles
Rechargeable lithium-oxygen (Li-0 2) batteries have the potential to provide gravimetric
energy three times or greater that of conventional Li-ion batteries 45 . The advantage of Li-0
2
systems stems from their anticipated lighter weight as the heavier transition metal intercalation
materials in Li-ion are replaced by a light porous cathode structure. As shown by Yi-Chun et al.6
(Figure 1), replacement of the typical LiCoO 2 in a Li-ion setting with Li 20 2 formation in a Li-0
2
setting could triple the expected gravimetric energy density of a cell using metallic lithium anode.
Although, more moderate gains are calculated for the volumetric energy density and both
gravimetric and volumetric densities using a LiC 6 cathode. In practice, laboratory Li-0 2 cathodes
16
have demonstrated gravimetric power and energy approximately threefold that of several typical
Li-ion battery materials (Figure
2)6,
albeit without important caveat discussed later. Altogether,
the Li-0 2 electrochemistry holds obvious advantages over Li-ion that justify the current scientific
interest.
2500
2500
2384..lo
f2000 -2000
2040
2002
1500
1500
LL
1000 >
99
1000
L
LL
500o
600
E
E
a
0
0
>LICoO. UL20 2
LICoO. LIA
4
2
Figure 1: Estimated gravimetric and volumetric energy density of LiCoO 2 and 02 (Li 2O2 ) as the
positive electrode with carbon (C6) or Li as the negative electrode. The cell voltages used for CLiCoO 2 , C-Li 2 0 2 , Li-LiCoO 2 , and Li-Li2 0 2 are 3.7 V, 2.45 V, 4.0 V, and 2.75 V, respectively.
Neither catalyst, carbon, nor electrolyte were included in the calculation for 02 cells. The
positive electrode (LiCo0 2 or Li 2 0 2 ) is assumed to be lithiated and an additional two times
excess lithium is used as the lithium negative electrode. Figure and caption reproduced from Ref.
[6] with permission from The Royal Society of Chemistry.
The structure of Li-0
2
cells is fundamentally described as a stack of metallic lithium
anode, an electrolyte layer, and an oxygen and electrolyte permeable porous cathode (typically
carbon although a nanoporous structure of gold has been demonstrated 7). The discharge reaction
in an Li-0 2 battery is the reduction of 02 with lithium ions to form an oxide of lithium. (Figure 3,
left): (1) 2Li
+ 02 <-
Li 2 0 2 at 2.96 V8'9 ; or lithium oxide (2) 4Li +
Only Li 2 0 2 is confirmed upon discharge of Li-0
2
cells1 0
12 . Recharge
02
++2Li 2 0 at 2.91 V8 ' 9 .
of the battery results in the
decomposition of the Li 2O2 deposit regenerating the lithium metal anode and clearing the
17
cathode structure for subsequent discharge (Figure 3, right). Mechanistic studies have shown that
013
formation of Li 2 0 2 proceeds through the sequence shown schematically below.1 '
b 02
02+e-
I
+Li+
+Li02
LiO 2
r
UNi Mn.0
WI
Disproportionation
2Li02
FHPC
Graphene [17]:
2
CNF (16]
S103
\
LiFePOe
VC[
102
Super P
[14) P-ZMnOJKB
(13]
10
103o
Gravimetric Energy (WhIkg .,)
102
4
Figure 2: Ragone plot comparing gravimetric energy and power of Li-ion (LiCoO2 , LiFePO4,-and
LiNio5 Mno ,O 2 and Li-0 2 positive electrodes tested in non-carbonate electrolytes and normalized
to the positive active electrode weight only. Li-0 2 electrodes include Vulcan Carbon (VC),
carbon nanofibers (CNF), Super P carbon, freestanding hierarchically porous carbon (FHPC
graphene), and pristine NaoMnO 2 nanowires/Ketjen Black (P-Z-MnO 2/KB). The Li-0 2 values
were normalized to the weight of the electrode in the discharged state (C + Li 20 2 or C + catalyst
+ Li 20 , excluding binder) and were calculated based on the reported average discharge voltage
and total gravimetric capacity (for energy) or current (power). The upper limit in the gravimetric
energy of Li 20 2 was calculated assuming a discharge voltage of 2.75 VU. Figure and caption
reproduced from Ref. [6] with permission from The Royal Society of Chemistry.
2
Little experimental evidence has been found for the strict reversal of the above sequence during
recharge and a direct two-electron decomposition of Li 2 O2 is postulated to explain experimental
observations.' 0 Yi-Chun et al.' 4 proposed that initial stage of recharge proceeds though
delithiation of the Li 2 0 2 surface to arrive at LiO 2 which subsequently disproportionate to oxygen.
On the other hand, the bulk of Li 2O2 appears to go undergo a two-phase (solid Li 20 2 directly to
02
gas) decomposition characterized by a plateau during galvanostatic charging.
18
e
e
Negative Electrode Electrolyte Positive Electrode
Negative Electrode Electrolyte Positive Electrode
2Li'+2e-+0
2 -
Li2 0
Li20
2
2
-2Li'
+2e- +02
Figure 3: Schematic of the discharge (left) and charge (right) of a Li-0 2 electrochemical cell.
Promising gravimetric energy densities notwithstanding, the Li-0
2
system face severe
obstacles to becoming industrially viable. Extensive research efforts have been devoted to
improving the performance of Li-0 2 batteries including round-trip efficiency 5 , power capability5
and cycle life7'1'17". An additional road block is the instability of common organic carbonate and
ethers used as electrolyte solvents most laboratory cells. The round trip efficiency of Li-0
2
batteries is of particular interest in the present thesis.
1.3 Challenges in Li-0 2 batteries
1.3.1 Inadequate power capability
The geometric power capability of Li-0 2 batteries is a major disadvantage with regard to
their application in designing EVs. In fact, while Li-ion cells can easily deliver 30 mW cm-2 ,
current laboratory Li-0 2 batteries will only offer -0.3-3 mW cm-2 with severe decay in cell
19
voltage as current is increased.1 8 Cell voltage was found to drop from 2.7 V to 2.4 V when
geometric current is increased from 0.05 mA cm
2
to ~1 mA cm-2 on Vulcan carbon
(uncatalyzed) cathodes. 5 The observed low power of Li-0
2
batteries has been attributed to
electron transfer limitations stemming from the highly insulating discharge product Li 20 2.
Measurements performed by Viswanathan et al.19 suggest a conductivity on the order of 10-12
S-cm
1
for the electrochemically formed Li 20 2.6 The picoSiemens range conductivity results in
~0.1 V drop in voltage for every nanometer of Li 20 2 formed and corroborate the rapid drop in
discharge plateau as rate is increased.5 Resistive Li 2 0 2 is not the only limitation to higher rates,
however. The kinetics of Li 20
2
formation are also sluggish and contribute significantly to
increased overpotential at moderate rates.
1.3.2 Inadequate cycle life
Laboratory Li-0 2 cells have not been cycled beyond 100 cycles and the moderate cycling
20
(10-50 cycles) reported resulted in substantial decline in cell capacities.16, Although recent
work by Peng and coworkers 7 demonstrates cycling up to 100 cycles, said Li-0 2 cells required a
cathode of noble nanoporous gold which sacrifices storage density and large scale economic
feasibility. The short cycling stability of Li-0
2
cathodes which are mainly based on carbon
matrices is shown to result from permanent deposit of parasitic carbonates from carbon corrosion
and electrolyte decomposition. 2 ' Viable Li-0 2 technology is thereby far from the 1000-cycles
target for 10 years of service-life in electric vehicles set by the U.S Advanced Battery
Consortium.
Nonetheless, cycle life is secondary to efficiency, power, and storage capability;
an optimally operating cell needs be developed prior to resolving cycling.
20
1.3.3 Inadequate round trip efficiencies: a combination of slow kinetics and
parasitic chemistries in Li-0 2 cells
1.3.3.1 Sluggish electrochemical kinetics
The discharge reaction proceeds with a Tafel slopes of 150 mV decades-
23
while
recharge displays -250 mV-decades-' with and without catalysts. 2 4 The Tafel slope is extracted
from the following equation (Eq 1) characterizing the kinetics of electrochemical reactions on
surfaces at sufficiently high overpotentials:
a77
Blog(i)
(Eq 1)
2.3-RT
a-n-F
where il, i, R, T, a, n, and F are, in that order, the overpotential, the electrical current, the ideal
gas constant, the process temperature, the transfer coefficient of the reaction, the number of
electrons involved in the rate limiting step, and Faraday's constant. This formula suggests a Tafel
slope of -60 mV-decades-1 for one-electron reactions and -30 mV-decades-I for 2-electrons
reactions at room temperature and assuming a maximum transfer coefficient of one.2 5 The Tafel
slope values reported in Li-0
2
cells are much larger than values predicted from the above
formula given the fact the maximum number of electrons involved in the formation of Li 2 0 2 or
Li 20 would be two. These large Tafel slopes point to either very low transfer coefficients (in the
order of 0.2-0.4 on discharge and 0.1-0.3 on recharge) or more complex coupled electron
transfers and chemical reactions that deviate from the Tafel description. We note that the small
transfer coefficients could be explained by the large changes in material phase from gas to solid
during discharge and vice versa during charge in the framework of the Marcus theory. The
apparent insensitivity of the large Tafel slopes to catalysis23 2 4 suggests that gains in efficiency
would have to be achieved through increase of surface exchange current densities, denoted by i0
21
in the Butler-Volmer equation (Eq 2) below. This observation, in turn points to catalysis, the
traditional means of enhancing exchanging currents in electrochemistry.
nF
i = i0 [e-aIi
_
-
nF1a
1]2
25
(Eq 2)
Sluggish kinetics are certainly major contributors to the large voltage hysteresis between
discharge and charge at rather low applied rates. However, it is worth noting that the sluggish
kinetics cannot be entirely decoupled from the formation of parasitic products during operation
of Li-0 2 cells.
1.3.3.2 Carbonate formation due to the organic solvent of the electrolyte
Owing to their close similarity with lithium ion batteries (Li-ion), Li-0
2
research
employed much of the same carbonate-based organic aprotic electrolytes. However, in contrast
to Li-ion batteries which are based on the intercalation/deintercalation of lithium cations within
accommodating material structures, discharge of Li-0
2
cells proceeds through the cathodic
reduction of molecular oxygen by electrons from a lithium anode to superoxide 02; The strong
nucleophile superoxide26,27 attacks positive positively charged centers of the aprotic solvents in
the absence of protons, resulting in the poor stability of carbonate solvent2 8 otherwise stable in
the superoxide-free Li-ion batteries. In fact, the superoxide formed during the first reaction step
is highly reactive towards most organic carbonates such as propylene carbonate (PC), ethylene
carbonate (EC), Dimethyl carbonate (DMC), and ethyl-methyl carbonate (EMC)29-32 and also
33
ethers such as tetraglyme (TEGDME), 1,3-dioxylane, 2-methyltetrahydrofuran . Chemical
reduction of the electrolyte solvent by superoxide anions in presence of lithium cations results in
excessive formation of Li 2 CO 3 , C 3 H6(OCO 2 Li) 2 , HCO 2Li, CH 3CO 2Li, esters, C0 2 , and H2 0 in
lieu of the desired Li 20232, 33 . Although ethers such as glymes and dimethoxyethane are found
more stable towards reduction by the superoxide 34' 35, steady buildup of unwanted solvent
22
degradation products occurs during cycling of Li-0 2 cells (Figure 4).2'3 The presence of
unwanted discharge products from degradation of the electrolyte solvent impedes fundamental
investigation of the Li 20 2 decomposition during recharge.
*1
I
4
0
z
30
Photon Energy
540
535
545
55
Photon Energy (eV)
(eV)
L i
O
a
"LoW" capacity
(-1000 mAhIg)
'igh- capacity
(-4700 mAWge)
0
282
294
Photon Energy
Morphology on First Discharge
(V)
Figure 4: (a) 0 K-edge FY, (b) 0 K-edge TEY, and (c) C K-edge TEY/FY XANES spectra of
electrodes discharged to 1000 (at 250 mA- g1carbon) and 4700 mAh-g Carbon (at 100 mA - g1Carbon)
on the 1st discharge. The reference spectra of commercial Li2 0 2 (90%, Sigma Aldrich),
commercial Li 2 CO 3 (99%, Alfa Aesar), and pristine VACNTs (FY for C K-edge) are included.
(d) Schematic of discharge products formed at low and high capacity on VACNTs on the 1st
discharge. Figure and caption reproduced from figure 3 of reference [21]; Copyright 2012
American Chemical Society.
23
1.3.3.3 Carbonate formation due to the carbon support used in most cathode
structures
Formation of lithium-carbonate compounds in the Li-0
2
is not limited to the
decomposition products of carbonate and ether electrolyte solvents (designation: electrolyteLi 2 CO 3 ). Using X-ray photoelectron 3 ' and X-ray absorption near edge spectroscopy 2', it has been
reported that reaction of Li 2 0 2 (or LiO 2 ) with the carbon matrix of the cathode is responsible for
the formation of about a monolayer of Li 2 CO 3 (designation: C-Li 2CO 3) at the immediate
interface of the solid discharge deposit. The presence of this monolayer during subsequent
recharge is possibly responsible for a 10-100 fold decrease in the exchange current at the
carbonjLi 2 0 2 interface.3 1 This Li 2CO 3-layer fundamentally changes the surface of the cathode
structure, impacting cycle life and performance.
1.3.3.4 Formation of parasitic Li-salt reaction products
Thus far, it has been made clear that formation of Li 2 0 2 during discharge of Li-0
2
batteries is undesirably accompanied , often superseded, by formation of lithium carbonates and
alkyl-carbonates which negatively affect reaction kinetics, increasing overpotentials and
reducing round trip efficiencies and cell cyclability. It is also the case that Li-0
2
batteries suffer
a third parasitic reaction pathway due to the nucleophilic nature of the superoxide anions formed
on discharge. Common lithium battery electrolyte salts such as lithium bis(oxalato)borate,
lithium
perchlorate,
lithium
bis(trifluoromethanesulfonyl)imide,
and
lithium
hexafluorophosphate are found to decompose to form parasitic deposits (Li 2 C2 0 4 , LiB 3 0 5 , LiCl,
LiF)36,37
24
1.3.3.5 Low round trip efficiencies
Altogether,
of Li-0
discharge
using
batteries
2
currently available
electrolytes
compositions results in the formation of Li/solvent, Li/carbon and Li/salt parasitic species arising
mostly from the reactive nature of the superoxide anion inevitably formed during discharge.
These unwanted discharge products adversely affect the kinetics of discharge/charge of Li-0
2
batteries and reduce cyclability. For qualitative estimation of the effects of the parasitic reaction
products on the kinetics of Li-0
2
batteries, the Gibbs free energies of formation of Li 20 2 , Li 20,
3 8 39
Li 2 CO 3 and LiCH 3CO 2 at standard temperature and pressure are provided below. '
2Li + 02
2Li + 02
2Li + C +
(Reaction 1)
L i 2 02 AGf(Li 2 02) = -571.91 kJ/mol
_
Li 2 0 2 AGf(Li 2 0 2 )
#
2
02
=
-561.91 kJ/mol
- L i2 CO3 AGf(Li 2 CO3 )
Li + 2C + 3H + 02 - LiCH3 C0
2
=
(Reaction 2)
-1132.1 kJ/mol
AGy(LiCH 3 CO 2 )
=
-663.6 kj/mol
(Reaction 3)
(Reaction 4)
Thermodynamically, these larger free energies of the parasitic discharge products suggest that
the presence of Li 2 CO 3 and LiCH 3 CO 2 in the electrode will increase the required overpotential
for recharge (products decomposition) of the Li-0
2
cells. This assessment is substantiated
experimentally by Freunberger and coworkers 32 by charging of the above-mentioned four main
carbonate parasitic species (Figure 5b). By preloading a super P/a-MnO 2/Kynar cathode with
commercially obtained Li 2 CO 3 , C 3 H6 (OCO 2 Li)2 , HCO 2Li, or CH 3CO 2 Li, and charging thus
assembled Li-0
2
cells at 70 mA-g1 Carbon in IM LiPF 6/PC, the authors report charging potentials
of ~3.8-4.2 V for Li2 CO 3 , ~3.75 V for HCO 2 Li, -3.8V-4.0 V for CH 3CO 2Li and -3.5-3.8 V for
C 3 H 6(OCO 2 Li) 2 . Notably, the charging of those compounds releases copious amounts of CO 2 , H2
and H2 0 (Figures 5c-f). Those decomposition voltages are much greater than the 3.5-3.65 V
range found for oxidation of Li 2 0 2 in the same type of cathodes.4 0
25
a
b
A
U
C 10,
45
4
........
105:...
a te
3
C 1
...
SC
H2
....
2.521
0
10
0.5
01
H2O ..
I
0
10 20 30 40 50
CO2..20
Mass (m/z)
(/)
Normalized capacity
Uformated1
010
-
n0
r
10
~ co2
2
0
H2
Qu2CJ3
10-10L io
U
0
10-10
102030
405
0 10
0
1------------Mass (emz)
---.
Mass (mz) 5
a
---
--
20304050
iupropyl
di
carbonate
0c
pristine electrode
0
2000
500
1000
1500
Wavenumber (cm- )
)time
u
0
0
250
200
400
600
800
(mn)
1000
1200 1400
Figure 5: Composite electrodes (Super P/a-MnO2 /Kynar) that contain the discharge products
individually were subjected to charging in 1 M LiPF 6 in propylene carbonate under02. (a) FTIR
spectra of the as-prepared electrodes and the charged electrodes for each of the compounds,
together with the spectrum of a pristine electrode. (b) The corresponding charging curves at 70
rnAg1Cabon. Since the theoretical capacities of the different compounds vary, to aid comparison
the capacities are all normalized to unity (theoretical capacities: Li propyl dicarbonate 1000
remol;
CH 3 C 2 Li 750 mAhgCbonI e
0AgCasbon 2 elmol; Li 2 C 3 1500 mAgs'Cbod 2
/mol; HC0 2 Li 750 mA h/g, le-/mol). (c-e) MS gas analysis at the end of charging under 02 Of
CH 3 CO 2 Li (C), C 3 H6 (0
fragments
(d), and HC0 2 Li (e). Note that unmarked peaks arise from
and Ar. (D)Gas evolution measured by DBMS on oxidation of a
2 Li)2
Of C0 2 , H 2 0, 02,
composite electrode containing Li 2 CO 3 in response to a stepwise increased current under Ar.
Figure and caption reproduced from Figure 5 of reference [32]; Copyright 2011 American
Chemical Society.
McCloskey et al.
31find
from first approximation electrochemical and charge transfer modeling
that C-Li2 CO 3 would decrease the exchange current density during charge by a factor of 10 to
100 just as the electrolyte-Li2CO3 would also drive up required charging overpotentials. A
26
similar finding appears in the work of Albertus et al. 4 ' who find an exponential increase in
cathode resistance with increasing deposit of materials likely of carbonate nature.
Low round trip efficiencies in Li-0
2
cells is a cumulative consequence the low power
capability from resistive Li 20 2 , slow kinetics of reduction and oxidation, and the pervasive
formation of parasitic products. Current efficiencies reported in the literature have not exceeded
80%, even under catalysis from noble metals to transition metal oxide. 16 ,24 ,40 ,42 ,43 Most energetic
losses occur on charge. Power (energy) loss on discharge is ~12% of the stored energy (most
discharge occur around 2.6 V at rates of ~70-100 mA-g'ICarbon (0.1 pA-cm- 2carbon) with little
variation from catalysts4 3 and cell OCV is 2.95 V for Li 20
2
formation) compared to an average
of 30% excess energy input required on charge (charging typically requires ~4 V and above
using uncatalyzed porous carbon cathodes).
E icicencydischarge
=
Thermodynamic stored energy - Dischargeenergy
Thermodynamic stored energy
77discharge
=
Eff icienc
arge
=
dreg
.
100% - 12%
(Eq 3)
Chargeenergy - Thermodynamic stored energy 100%
Thermodynamic stored energy
=
lcharge-
100%
= 30%
(Eq 4)
Erey
It is thus logical that short term efficiency improvement efforts be directed towards catalyzing
the oxidation of Li 2O 2 (charging) which displays much worse kinetics compared to discharge.
Significant resources have been devoted to identifying such reaction promoters for the recharge
of Li-0
2
battery as summarized in Table 1.
27
Catalyst
Electrolyte used
Rate
(mA g~
Charging
voltage
Carbon)
(V)
-3.6-3.7
Pt/NC
Ru/NC
Au/NC
-3.6-3.7
0.1 M LiClO4 1,2
Dimethoxyethane
-4.2
Vulcan Carbon (VC)
-4.1
SP/a-MnO 2 NW
SP/ a-MnO 2 bulk
SP/EMD
-3.55
-3.7
-3.7
-3.85
-4.0
-4.12
-4.15
SP/Co3 0 4
SP/CuO
SP/NiO
SP/a-Fe 2 0 3
Super P (SP)
SS/Fe 30 4
1 M LiPF 6
Propylene
carbonate
70
-4.25
-3.8
~4.25
-4.35
-4.25
-4.75
-4.0
1 M LiPF 6
Propylene
carbonate
Li-0 2 cell
(Carbon/Catalyst 4=3 4 4
95/2.5 molar ratio) '
Li-0 2 cell
(Carbon/Catalyst
95/2.5 molar ratio)45
-3.83
44 MnO 2
-4.0
KB/Pristine Nao. 4 4 MnO2
Ketjen black (KB)
KB/Lead ruthenate
KB/Bismuth ruthenate
Ketjen Carbon (KB)
Li 2 0 2-prefilled
(SP/Catalyst/Li202 =
40
1/1.7/1 mass ratio)
-4.0
KB/Acid leached
Nao
Li 20 2-prefilled
(VC/Catalyst/Li202=
1/0.66/1 mass ratio) 24
-4.3
SS/Co 3 04
SS/CuO
SS/CoFe 2 04
SS/EMD
Super S (SS)
SP/a-MnO 2 NW
SP/NiFe 20 4 (183m 2/g)
Cathode structure
-4.15
-4.0
-4.0
-4.2
1 M LiPF 6
TEGDME
KB/
20.1
La 1 .7Cao. 3Nio. 75Cuo.2504
-3.62
Li-0 2 cell (KB/Catalyst
= 1/0.4)46
Li-0
2
cell (KB/Catalyst
=
1/1)47,48
Li 20 2-prefilled
(KB/Catalyst/Li 20 2 =
1/0.3/0.3)49
Table 1: Literature values for Li20 2-oxidation activities under various cell conditions
28
29
2 Scope of this thesis
2.1 Thermal Stability Studies of Li 2 0 2 and Li 2 O
The above introduction listed representative findings of a great many investigations of the
Li-0 2 electrochemistry. The safety aspect of the rechargeable Li-0 2 batteries, which is one of the
most important considerations in practical applications, has yet to be considered to date. In
general, inorganic peroxides are highly reactive and may undergo violent decomposition
reactions due to their weak 0-0 bonds50 . The activation energy for decomposition of the Li-0
2
discharge product, Li 2 0 2 , is reported at ~50 kcal/mole5 1 , in good agreement with reported
decomposition enthalpies of most known, highly reactive peroxides such as K20250.
Consequently, the thermal stability of the Li-0 2 reaction products, Li 2 0 2 and possibly Li 20, is of
critical importance toward the realization of the practical Li-0
2
batteries. In the present thesis,
we investigate the structural and chemical changes occurring in Li 2 0 2 and Li 2 0 at elevated
temperatures. This contribution is valuable to gaining basic initial insights into thermal behavior
of Li-0
2
independently of the specific cathode structure chosen. Foreseeable extensions to this
work are discussed.
2.2 Catalyzing the Li 2 O2-oxidation in Li-0
2
The round trip efficiency and rate capability of Li-0 2 batteries is undeniably a major and
most noticeably considered obstacle to their practical application. The need to improve surface
exchange current densities to boost charging rates invokes the use of catalysis. Efforts to enhance
the Li 20 2 -oxidation have demonstrated that charging voltage can be lowered (as compared to
pure carbon cathodes) using traditional noble metal catalysts as well as inexpensive metal oxides.
Harding et al.2 4 demonstrated, using the same method utilized herein, that a two order of
30
magnitude increase in electrode activity could be achieved on charge using the noble metal
catalysts platinum and ruthenium. More recently non-noble transition metal oxides have been
used to influence both the charging of Li-0
2
cells 4 3 ,4 6 ,4 7 ,52 and Li 20 2-prefilled 4 0' 4 9 electrodes
(Table 1). For example, MnO 246 (70 mA g-carbon, 1 M LiPF 6 tetraglyme, 3.8-4.0 VU charging),
pyrochlores 47' 48 (70 mA g~1carbon,
1 M LiPF 6 tetraglyme , 3.9-4.0 VU charging), and
La 1 .7 Cao. 3 Nio. 7 5Cuo.25 0 4 layered perovskite 4 9 (20 mA gIcarbon, 1 M LiTFSI tetraglyme, 3.62 VU
charging) significantly decreased the oxidation potential of Li 2 0 2 . Nonetheless, the variety of
testing conditions used (applied rates, cut-off voltages, the amount of Li 20
2
present in the
electrodes prior to charging, electrolyte employed etc...) impedes potential understanding of the
mechanics of Li2 0 2 -oxidation based on trends, a technique ubiquitously used in the study of
aqueous oxygen reduction and evolution. 53-5 6 In this thesis, we employ the versatility of
perovskite oxides and their advantage as a systematic catalyst system to investigate the Li 20 2 oxidation reaction with the goal of improving recharge rates and gaining mechanistic insights
into oxygen evolution involved in the oxidation reaction.
We begin with a report on results obtained on the catalysis of Li 2 0 2 oxidation by ABO 3-type
perovskites. Later, we detail findings on the thermal transformation of Li 2 0 2 and Li 20 at
elevated temperatures.
31
32
3
Catalyzing the Li 2 O2-oxidation in Li-0
2
batteries using
ABO 3-type perovskites
3.1 ABO 3-Perovskites as a self-consistent platform for systematic
study of Li 2 O2-oxidation in Li-0 2 cells
Investigation of the H2 0-oxidation in aqueous 0.1 M KOH by Suntivich et al. 53 provides
an immediate set of candidate catalysts for the systematic study of the Li 2 0 2-oxidation reaction
involved in the charging of Li-0
2
batteries. Such systematic study has the potential of not only
revealing fundamental processes involved Li 2 O2-decomposition but also material descriptors for
guided search of future high-activity catalysts.
Antibonding states
e
//
Transition
Metal 3d
t2g*
t
OER
/
Oxygen 2p
eg
Bonding states
Figure 6: Schematic of molecular orbital splitting of the transition metal 3d band engendered by
hybridization of the metal 3d and oxygen 2p orbitals during adsorption.
In their investigation, Suntivich et al. E take a molecular orbital approach to describe the
oxygen-evolution (OER) activity of the perovskite oxides in aqueous 0.1 M KOH. More
specifically, the electron distribution of the surface transition metal 3d into eg* antibonding
orbitals (Figure 6) is found to describe the interaction strength of the perovskite with the
33
adsorbed oxygenated species during OER. This hypothesis led to a strong volcano correlation
between the antibonding eg* occupancy of perovskites catalysts and their aqueous OER activity
(Figure 7). The volcano correlation is proof that the antibonding eg* filling of the perovskite is a
proxy to the strength of interaction of oxygenated adsorbates involved in the rate determining
step (as defined in the Sabastier principle: neither too weak nor too strong binding of adsorbates
for optimal heterogeneous catalysis).
1.4
,
.
.
.
,
.
,
@ /= 50 ptA cm
LaNIO3
15
1.6
LaCOO3
LaoCa 5MnO
-
LaMnu
LaMnOW
1.7
BaSrCoWe2O3
LaCa.CoOH
LaMnNi.O,
La CaFeO
LaMnOa
LaCrO3
1.8
0.0
0.5
1.0
1.5
2.0
2.5
e, electron
Figure 7: The relation between the OER catalytic activity, defined by the overpotentials at 50 RA
cm 2,x of OER current, and the occupancy of the eg-symmetry electron of the transition metal (B
in ABO 3). Data symbols vary with type of B ions (Cr, red; Mn, orange; Fe, beige; Co, green; Ni,
blue; mixed compounds, purple), where x = 0, 0.25, and 0.5 for Fe. Error bars represent standard
deviations of at least three independent measurements. The dashed volcano lines are shown for
guidance only. Figure and caption adapted from reference [53]. Reprinted with permission from
AAAS.
As seen in Figure 7, Bao. 5Sro.5Coo. 8Feo.20 3 (BSCF) was shown to have high activity toward H20oxidation in 0.1 M KOH aqueous media as a result of its optimal eg*-occupancy (eg ~ 1). The
remaining perovskites with eg*-occupancy much less or much greater than one, display
diminished activity compared to BSCF. Here, we systematically examine whether such a
correlation exists for Li20 2-oxidation reaction in the relatively stable 0.1 M LiClO 4 DME
34
electrolyte using electrodes pre-packed with commercial Li 202.28, 5 7 This approach has many
advantages as detailed in the experimental section.
35
36
3.2 Experimental
It is clear that discharge and subsequent recharge of Li-0
2
batteries in available
electrolytes would result in biased findings regarding the intrinsic electrochemistry of Li 20 2 oxidation. In the present work, we are concerned with probing the catalytic oxidation of Li 202
using perovskite catalysts. Consequently, we bypass the oxygen reduction step involved in the
discharge (unequivocally identified as the primary source of non-Li 2O 2 byproducts) by
preloading ex-situ chemically synthesized Li 20 2 . In so doing, we gain control over the amount of
Li 2 0 2 loaded within each cathode structure (defined discharged state) under study.
0.3
-
5
*
-%CO2
* 0.1
,a
0
4
S
12
16
Time (hours)
20
24
Figure 8: Variations of charge voltage and gas compositions (helium not included) during the
first charging process for the Li 20 2 /Fe30 4 /SP/PVDF electrode in a carbonate electrolyte. Figure
and caption adapted from reference [58], with permission from Elsevier.
Furthermore, the chemical nature (stoichiometry of the Li2 0 2 , eventual contaminants...) is
maintained constant across the various cathodes compositions under investigation. We highlight
here that oxidation of ex-situ synthesized Li2 0 2 manually loaded in a cathode structure has been
shown to proceed with minimal side reactions, releasing molecular oxygen and benign amounts
of CO 2 (only toward the end of recharge) 58 as shown with differential electrochemical mass
37
spectrometry (DEMS, Figure 8). The methodology allows a systematic performance comparison
of various perovskite catalysts during the oxidation of Li 20 2. We hope that such a systematic
study will inform later design of Li-0
2
batteries as progress is made towards discovery of more
stable electrolytes.
Vulcan carbon (VC-only) and five perovskites including LaCrO 3 , Bao. 5 Sro. 5 Co 0 .8Feo.2 0 3
(BSCF), LaNiO 3, LaMnO 3+6 , and LaFeO 3 (selected to span the meaningful range of eg-filling
from eg
=
0 to eg
2.0, Figure 7) were investigated. All perovskites and Li 20 2 (Alfa Aesar,
Purity: > 90%, ball milled to -345 nm) powders were ball-milled separately using a planetary
ball mill (Pulverisette 6, Fritsch Inc., sealed argon-filled zirconia crucible) at 500 rpm. Ballmilling of Li 2 O2 was performed with the crucible sealed in a "heat-seal" bag filled with argon.
This measure is necessary to avoid reaction of the Li 2 0 2 with ambient moisture and CO 2 which
would result in excessive formation of undesired surface Li 2 CO 3 and LiOH. Electrodes were
synthesized by preloading the ball-milled commercial lithium peroxide (Li 20 2 , Alfa Aesar,
Purity: > 90%) in the perovskite-catalyzed matrix of Vulcan carbon (Vulcan XC72, Premetek
Inc., 100 m2.
Carbon).
3.2.1 Synthesis of the perovskite oxides
Two methods were used to obtain the perovskites investigated in the present thesis: coprecipitation and nitrates combustion. All methods have been reported previously, therefore, are
described briefly below.
Co-precipitation":
This method was used for the synthesis of LaCrO 3, LaNiO 3 and LaMnO 3+6 . Nitrates of
lanthanum and the transition metal (99.98%, Alfa Aesar) were mixed in de-ionized water (Milli-
Q water, 18 MQ-cm) at metal molar ratio of 1:1 and total concentration of 0.2 M. The solution
38
was subsequently titrated using an aqueous 1.2 M solution of tetramethylammonium hydroxide
(100%, Alfa Aesar) resulting in precipitation. The precipitate was then filtered and collected to
dry. Finally, the precipitate powder is heat treated in a tube oven at -1000 'C under dry air for
approximately 10 hours.
Nitrate combustion5:
This method was used for the synthesis of Bao.5 Sro.5 Coo. 8Feo.2 0 3 and LaFeO 3 . Nitrates of the rare
earth and transition metal cations (Sigma Aldrich, >99.99%) were mixed in a 2000 mL beaker at
the required molar ratios of cations and total metal concentration of 0.2 M. Approximately, 0.1
M glycine was added to the mixture and homogenized using a magnetic stir plate. The mixture
was heated until full evaporation of the water, followed by combustion of the solid deposit
within the beaker on the heating plate. The powder was collected and heat treated under dry air at
~1000 'C for 24 hours in a tube furnace.
Investi2atingz phase purity of the perovskite by X-ray diffraction
Purity of the synthesized perovskites was investigated using a PANanalytical X'Pert
ProTM X-ray diffractometer with copper Ka wavelength (k = 1.5418). All obtained materials
were confirmed to be optimally pure (Figure 9). Some minor impurities estimated below 1% of
the total perovskite phase were observed for LaCrO 3 and LaMnO 3 and are not expected to
influence the subsequent electrochemical studies.
39
-
-J
-
A
A
20
40
30
LaFeO 3
50
20 CuKa
A
BSCF
.......
--k LaM nO3
A
/-'LaCrO 3
-- 1
80
60
70
Figure 9: Phase purity of as-synthesized perovskites investigated by X-ray diffraction. Optimal
purity of each perovskite is observed. Minor impurity phases estimated to less than 1% (peaks
not very visible from scaling) were detected for LaCrO 3 and LaMnO 3.
3.2.2 Electrode ink preparation
Carbon, perovskite oxide catalyst, Li 20 2 , and lithium-exchanged nafion binder (Ion Power
USA, LITHion TM , 7.2 wt% binder) were homogenized in isopropanol (IPA) by probe pulse
sonication at 40 W for an hour (Figure 10) inside an argon-filled glovebox (MBraun Inc., H2 0
0.1 ppm and
02
<
<
1 ppm). Component mass ratios were set to perovskite:vulcan
carbon:Li 20 2 :lithiated nafion binder = 3:1:1:1 (carbon content = 0.21±0.05 mg cm-2). The high
mass loading of the catalysts was chosen to compensate for the low surface areas of available
perovskites (Table 2) in order to resolve the effect of the perovskites on the Li 20 2-oxidation. In
the case of Vulcan-only electrodes, mass ratios were set by simply omitting the perovskite from
40
the electrode to have Vulcan carbon:Li 20 2 :lithiated nafion binder
=
1:1:1. The viscous slurry
thus obtained was then casted on an aluminum foil.
Carbon C
1X C§
O
0 0---Tip
Oxide Cat.
3X
sonication
L
1X
O0
L
Li2 0 2
L
LiNafion
iX
,-
O
40 W pulses
L.
1hour
L
&
Mixing in ~2 mL
isopropanol
Figure 10: Schematic of electrode ink preparation. Component mass ratios were set to
perovskite:vulcan carbon:Li 20 2 :lithiated nafion binder = 3:1:1:1 (carbon content = 0.21 0.05 mg
cm-2).
Oxide
BSCF64.2L9
LaMnO
dvia (nm)
Surface Area (m /g )
574.53
1.821
LaNiO 3
206.114.8
LaFeO
454.05
1.981
LaCrO,'
951.08
G.,94.6
Table 2: SEM calculated particles size and surface areas of the perovskites investigated.
3.2.3 Electrode film casting
Upon synthesis of the inks, electrodes were fabricated by liquid film coating on an
aluminum foil using a #50 Mayer rod, still within the same argon filled glovebox. Half-inch
diameter electrodes (area = 1.27 cm2) were punched after complete evaporation of the IPA. The
drying glass tube of a Buchi* B-585 vacuum oven was inserted into the glovebox chamber,
loaded with the freshly fabricated electrodes, sealed full of argon, and brought to 70 'C for at
41
least 12 hours (Figure 11). All fabrication tools were dried in a convection oven at 70 *C
overnight prior to use. No instance of exposure to atmospheric moisture occurred from
fabrication of the electrodes to drying to cell construction; the chemical integrity of the preloaded
Li 2 0
2
is maximized (as discussed in section 2, atmospheric exposure of Li2 O2 can significantly
compromises its surface).
Drying in
Buchi* Oven
Mayer-Rod #50
Al. foil
Punching of 1/2
inch diameter
electrodes
Ink coating on
Aluminum foil
Figure 11: Fabrication of electrodes following ink synthesis. From left to right, the homogenized
ink is drawn on an aluminum foil using a #50 Mayer-rod; upon evaporation of isopropanol, halfinch diameter electrodes are punched and dried at 70 'C in a Buchi* B-585 oven. Loading,
drying and returning the electrodes to the glovebox using the Buchi* oven prevents exposure of
the electrodes to ambient atmosphere.
3.2.4 Electrochemical testing
The above described electrodes were tested for their electrochemical performance within 2electrodes cells (TJ-AK; Tomcell Japan Inc.) as described elsewhere. 24 Electrochemical cells
made of a lithium anode, 150 tL of 0.1 M LiClO 4 1,2 dimethoxyethane electrolyte (Novolyte
USA, H2 0 < 20 ppm) on two Celgard* C2500 separators, and the fabricated cathode were
assembled in an argon-filled (H2 0 and 02 contents below 0.1 ppm) glovebox and
electrochemical data were collected using a Solartron 1470 (Solartron Analytical, UK). The
42
catalytic performance of each electrode was defined as the net current output after "background
substraction".
43
44
3.3 Results and Discussion
3.3.1 Effective electro-oxidation of preloaded commercial Li 2 O 2
X-Ray diffraction (XRD) patterns of pristine (not yet electrochemically charged) electrodes
show peaks assigned to the crystalline perovskites catalysts used. In addition to the perovskite
peaks, a strong peak at ~34.97 assigned to the (101) crystallographic plane of preloaded Li 2 0 2 is
observed in all electrodes (Figure 12). Secondary peaks are also observed which unambiguously
confirm the presence of crystalline Li 2 0 2 within all electrodes. After potentiostatic charging at 4
disappears from all electrodes along with all secondary peaks
VLi, the primary peak at -34.97
due to Li 2 02 (Figure 12). The preloaded Li 2 0 2 is effectively oxidized, and the observed electrode
current can safely be attributed to the intended Li 20 2-oxidation. Figure 13 provides scanning
electron microscopy (SEM) images of as-prepared (uncharged) versus charged electrodes for
representative LaCrO 3 and BSCF catalyzed electrodes. Upon charging, all electrode structures
feature empty spaces with size matching the observed -350 nm diameter of preloaded Li 2 02
particles. The presence of Li 2 0
2
in the uncharged electrode and their subsequent disappearance
after charging indicate effective removal of Li 2 0 2 by electrochemical oxidation.
3.3.2 Electrode background subtraction
An important concern during the oxidative electrochemistry in this study is the eventual
significant parasitic oxidation of the cathode structure (carbon + perovskite catalyst) in lieu of
the desired Li 2 O2-oxidation. XRD and SEM data presented above provided proof of reasonably
complete removal of Li 2 0 2 . Nonetheless, we ensure minimal parasitic oxidation of the cathode
structure by studying the electrochemical response of so-called "background electrodes".
Electrodes fabricated without the reactant Li 2 0 2 (perovskite:Vulcan carbon:lithiated nafion
45
binder = 3:1:1, designated "background electrode") are polarized at 4.0 VU. As shown in Figure
14, little current is observed from the background electrodes. This result further substantiates that
the current measured in the Li 20 2-packed electrodes is due to Li202-oxidation. Net mass-specific
current (normalized to carbon mass, imass) of the background electrode is subtracted from that of
the Li 20 2-prefilled electrodes in the time domain to arrive at "net currents" (Eq 5, Figure 14).
__
imass
-
Iobserved
-electrode
inet ~~ mass
Q
-15
[mA - gcarbon]
Carbon mass
-background
~ mass
(Eq 5)
Sm'carbon ]
(Eq
Eq66)
= ftjme=o inet -adt [mAh gcrbon]
iarea = 1net '
cataostmass
(Eq 7)
Catalyst Area
[MA
-m 2 ]
(Eq 8)
Cell capacity is calculated by integrating the "net current" in time (Eq 7). Area-specific currents
(normalized to the true area of the perovskite catalysts reported in Table 2) are also calculated
from the "net current" (Eq 8). Catalyst activity towards Li2 0 2-oxidation is quantified using both
the net mass-specific current and the net area-specific current seen from each Li 20 2-prefilled
electrode. In the case of surface-area specific activity, the mathematical average of current is
considered in the range of 10 to 60% of the total charge Li 20 2-charge loaded in each electrode.
Our electrodes were prefilled with ~1168 mAh-gcarbon equivalent to a mass ratio Vulcan
carbon:Li 20 2 equal 1:1.
46
C+U 20 2 (No perovskite) electrodes
C+LaNIO 3+L2 02 electrodes
Charged
Charged JL-J
Uncharged
Unchargedq
z
40
30
20
50
60
20
40
30
20 CuKa
I0
50
60
20 CuKa
C+LaFeO3 +U2 0 2 electrodes
C+LaMnO 3 +6 +U
2 0 2 electrodes
0
0
0
Charged
Uncharged
30
20
40
50
60
io
o
40
50O
6D
20 CuKa
C+LaCrO 3+U20 2 electrodes
. 50CO.8F 0 .20 3+U20 2
C+8e 0 5 Sr
U.
electrodes
0
0
I
20
30
I
40
I
50
I
60
20 CuKa
II
20
30
40
50
60
20 CuKa
Figure 12: X-Ray diffractions of as-fabricated (Uncharged, black lines) and post-charging at 4
VL (Charged, red lines) perovskite catalyzed Li2 0 2 -prefilled electrodes. For all five perovskitecatalyzed electrodes, the strongest peak of Li20 2 at ~34.970 from the (101) crystallographic plane
disappears after charging, indicating its effective oxidation. Keys: (*) Li2 0 2 crystal Cu Ka peaks
locations; (o) Corresponding perovskite Cu Ka peaks locations. Residual LiClO 4 salt and
aluminum substrate peaks are indicated on the figure. As intended, no catalyst peaks are visible
in the Vulcan carbon-only electrode.
47
Uncharged electrode
1 im
Charged electrode
0
4-,
(N
_j
0.
U
U
-o
Figure 13: Representative SEM images of pristine (Uncharged) versus post-charging at 4 VU of
Ba0 .5 Sro.5 Coo.8 Feo.2 0 3 and LaCrO3 catalyzed Li 20 2-prefilled electrodes. Key: (o) corresponding
catalyst particles locations; (o) Li 2 0 2 particles location. After potentiostatic charging at 4 Vi, no
trace of Li2 O2 particles can be observed by SEM in the charged electrodes. The same is observed
for all other catalyzed electrodes studied within this report.
48
100
Electrode w/ Li20 2100
et activity of electrode
Net activity of electrode
electroderode
10
10
Background electrode: Backgroun
NO Li 2O
0
20
40
60
80
0
10
Time (hrs)
20
30
40
Time (hrs)
Figure 14: Examples of background subtraction performed on (a) BSCF-catalyzed and (b)
LaCrO 3 -catalyzed electrodes at 4.0 VU. Little change is observed in the final current (Net
activity of electrode), which highlights the negligible magnitude of parasitic currents compared
to actual Li20 2-oxidation currents. Negligible and featureless current curves of the electrode with
no Li 20 2 compared to electrode with Li 20 2 confirms that the observed performance of peroxide
packed electrodes is due to effective oxidation of Li 2 0 2.
3.3.3 Catalytic
performance
of
LaCrO3,
LaMnO3+8,
LaNiO 3,
Ba0 . 5Sr 0. 5Co0 . 8Fe 0.2 0 3 , and LaFeO 3 perovskites
Net mass-specific activity under perovskite catalysis at 4.0 VU are presented in Figures
15a-b. Increase in gravimetric current over that of the baseline uncatalyzed Vulcan carbon is
observed for oxidation of the preloaded Li 20 2 . All tested electrodes deliver between ~800 to
~1000 mAh-gcarbon net charge out of the estimated 1168 mAh-gCarbon equivalent of Li 2 0 2 -
preloaded. Therefore, approximately 70 to 80% of the expected Li 2 O2 is oxidized by all
electrodes at the applied voltage of 4 VU, which is in agreement with the absence of Li 20 2 signal
from the XRD and SEM presented. Although observed net mass-specific activities from all
perovskite-catalyzed electrodes is greater than that of uncatalyzed carbon (-30-125 mA-g-Carbon
for the perovskite-electrodes versus -20 mA-gcarbon for uncatalyzed carbon), the possibility of
surface area effect is investigated using area-specific activities (Figure 16).
49
a)
BSCF
a 4.0 VU
0
LaCrO
b)
4.0 VU
0
LaNiO L~O.
100
Carbon
Carbon
0
0
400
200
600
800
1200
1000
Net charge (mAh/g ,b)
Net charge (mAh/gc,w)
Figure 15: Potentiostatic electrochemical performance of perovskite-catalyzed electrodes
(perovskite:Vulcan carbon:Li 2 0 2 :lithiated nafion = 3:1:1:1 mass ratio) at 4.0 VU (carbon content
= 0.21±0.05 mg cm-2). (a, b) Mass-specific activity of electrodes with BSCF, LaNiO 3,
LaMnO 3+8,LaCrO 3 and LaFeO 3 compared to an uncatalyzed Vulcan carbon electrode (Vulcan
carbon:Li 2 0 2 :lithiated nafion = 1:1:1 mass ratio). An increase in the electrode current output is
observed after addition of the perovskites.
10a)
@4.0 Ve
b)
@4.0 Vu
0.1
00
09
90
-_
0.1 0.0
0.4
0.8
1.2
1.6
2.0
0.01
0
0
-~ --
o46
a
00
e ectrorn
Figure 16: (a) Catalyst area-specific activity of (Perovskite:Vulcan carbon:Li 2O2 :lithiated nafion
= 3:1:1:1 mass ratio) versus filling of eg* antibonding orbital. (b) Area-specific activity of
(Perovskite:Vulcan carbon:Li 20 2 :lithiated nafion = 3:1:1:1 mass ratio) electrodes normalized to
the combined [Carbon+Catalyst]-total surface area. The average area-specific activity of baseline
carbon is added (dotted line). Activities of LaCrO3 and Bao.5 Sro.5 Coo.8 Feo. 2O 3 are well above that
of baseline carbon, which proves their activity cannot be explained by mere surface area effects
but rather actual catalysis. Catalytic effects from LaNiO 3 , LaMnO 3 +8 , LaFeO3 , and
Lao. 5Cao.5FeO 3.6 are less unambiguous, and considered comparable to carbon, considering
experimental errors.
50
From Figures 16a-b, we observe that the surface area activities of BSCF and LaCrO 3 catalyzed
electrodes are well above that of the carbon-only baseline while activities from LaNiO 3,
LaMnO 3+6 , and LaFeO 3 are considered similar to that of uncatalyzed carbon (accounting for
experimental errors).
Motivated by their more discernible catalytic effect on Li 2 O2 -oxidation, we proceed to
investigate the performance of BSCF and LaCrO 3 as a function of applied potential from 3.8 VLi
to 4.1 VLi (Figure 17). The superior Li20 2 -oxidation activity of BSCF and LaCrO 3 compared to
uncatalyzed Vulcan carbon and Au/C is notable at all applied potentials, as quantified both on
mass and surface area bases. Strikingly, the area-specific activity of LaCrO 3 rivals that of highly
active and nanodispersed noble metal Pt/C and Ru/C reported previously (Figure 17b). 24
b)
a)
1000
AO10
10
.1
I1
3.4
3.6
3.8
4.0
4.2
4.4
4.6
IE-3
3.4
3:
3.6
3:
3.8
4.
4.0
4.
4.2
44
4.4
4.
4.6
Voltage (V)
Voltage (V)
Figure 17: (a) Mass-specific activity vs. potential for LaCrO 3, BSCF compared to Au/C, Pt/C,
Ru/C (Pt,Ru,Au:Vulcan carbon:Li 20 2 = 0.66:1:1; Perovskite:Vulcan carbon:Li 20 2 = 3:1:1, mass
ratios) and VC-only reported previously 24 (b): Surface-area specific activity vs. potential for the
same electrodes. Considering that there is ~12 to 20 times more surface area in the noble metal
electrodes (-60 m 2 catalytic/gcarbon) compared to the perovskite electrodes (3-5 m 2 catalytic/gcbon), the
present activity of LaCrO 3 and BSCF is close to that of Pt/C and Ru/C on a surface area basis.
Tafel slopes are ~250 mV/decades.
The mass-specific activities of BSCF and LaCrO3 being intermediary to that of uncatalyzed
2
carbon and the noble metals can be attributed to their micron-sized particle (0.5-1pm, ~1 m .g
for the perovskites) reported in Table 2 compared to nanometer-size particles for the noble
51
metals (5-10 nm, ~100 m2 -g for the noble metals). The Tafel slopes on BSCF and LaCrO 3 ,
estimated to ~ 250 mV/decades, are in good agreement with that observed on for the noble
metals and uncatalyzed carbon electrodes. These similar Tafel slopes indicate similar ratedetermining steps on all surfaces utilized for the oxygen evolution reaction (OER) involved in
Li 20 2-oxidation.
3.3.4 Strongly diverging activity patterns of perovskites from H 2 0-OER to
Li2 O2 -OER
The area-specific activities of the five perovskites studied are plotted against their eg*
orbital filling in Figure 16a. Area-specific activities are ordered as follows: LaCrO 3 >> BSCF >
LaNiO 3
LaMnO 3+8~ LaFeO 3. No volcano-shaped trend vs. eg-electron count was observed for
the activity of the perovskites during the OER from Li2 0 2 -oxidation. This is in contrast to the
clear volcano trend identified by Suntivich et al. 53 during the OER from H20-oxidation in 0.1 M
KOH (Figure 7). Furthermore, LaCrO 3 (the lowest activity perovskite in aqueous OER) is found
to have -4 times higher area-specific activity than BSCF (the best aqueous OER catalyst) during
non-aqueous OER from Li 20 2 -oxidation. Again, this result contrasts starkly against more than
two orders of magnitude higher activities found for BSCF compared to LaCrO 3 in aqueous OER.
3.3.5 Proposed origin of observed divergence between H 2 0-OER to Li 2O 2 -
OER
Researching the literature on aqueous H20-oxidation and Li 2 0 2 -oxidation, additional
reversals in activities activity trends are identified and summarized in Table 3. Consequently, the
case of BSCF vs. LaCrO 3 is not isolated and points to significantly different reaction mechanics
from H20-oxidation to Li 20 2 -oxidation. We note that the conventional electrocatalysis for H2 0-
52
oxidation mainly concerns the liquid H2 Olsolid (electrode) interactions, while the Li 2 0 2 oxidation is strongly influenced by the interactions between both solid Li 2 0 2 and solid
(electrode) and solid (Li 20 2 ) and liquid (electrolyte). This difference is expected to contribute to
the differences in OER activity trend observed in aqueous and nonaqueous environments.
H 20-OER
Li 2 O2 -OER
LaCiO>B>
La iO0 (bt
MnO 2 > Co 304
MnO2 < C0 3 0 4
Table 3: Examples of activity reversal patterns from Li20 2-OER to H20-OER found in literature.
We hypothesize that the observed Li 20 2-oxidation activity is governed by the ability of the
catalyst to promote and/or stabilize oxygen rich Li 2-x02 species. LiO 2-like species observed by
Raman spectroscopy and SQUID display a low charging potential below 3.5 VLu at 62.5 mA g~
Carbon-60
We note in that regards that C0 3 0 4 , MnO 2 and Cr0 2 (surfaces of LaCrO 3 are likely
covered with oxides of chromium parentage as is the case for most chromium oxide surfaces 61)
surfaces have comparable calculated oxygen affinity in the range of 1.64-3.08 eVs4, 62 which may
aid in stabilizing such an electrically conductive 63 oxygen rich Li2 -x02 layer at their interfaces
with bulk Li 20 2. We reasonably speculate that interfacing of the metal oxide with the Li 2 0 2
particles results in modification of facets (surface orientations listed in Table 4) distributions on
the Li2 0 2 Wulff shape. 64 In effect, the surface energy landscape of Li 2 0 2 would be influenced by
the presence of the metal oxide, eventually biasing the surface to favorable terminations such as
(1 1 01).64 Such mechanism rooted in enhancing the oxidation of Li 2 0 2 via preferential
stabilization of low overpotential surface orientations would explain the catalyst-invariant 3.2
VU onset potential for Li 20 2 -oxidation (Figure 18).
65
This onset potential is in agreement with
the smallest required overpotential of 0.2 V for Li2 0 2 -OER calculated on (1T 01) surface facets.
53
UhAA
Figure 18: Linear sweep voltammetry following an hour long discharge at 2.25 V. Figure and
caption adapted from Figure 2e of reference [65]. Copyright 2013 WILEY-VCH Verlag GmbH
& Co. KGaA, Weinheim.
Within the scope of the present postulate (1T 01) all terminations would be available regardless
of the catalyst employed. The catalyst would only dictate the relative proportion of each
termination so as to define attainable rates of oxidation at the various potentials (without
changing the onset of oxidation). This observation is somewhat in line with the work of Black et
al.6 s who suggested transport
rather generation of Li 2 xO2 to be responsible for observed
enhancements. Further studies are ongoing to understand the origin of the differences in the OER
mechanism in aqueous and nonaqueous media.
Surface orientation
(IT 01)
(1T 00)
(1121)
(0001)
(1120)
Overpotential1 (V)
0.2
;;27
3
A
6
Table 4: Theoretical overpotentials for Li 20 2 -oxidation as a function of surface orientation.
Table data adapted from reference [64].
54
55
3.4 Conclusion
To summarize, we carried a systematic study of the catalytic activity of LaCrO 3 ,
LaMnO 3+6 , LaNiO 3 , Bao. 5 Sro.5 Co 0 8 FeO.20 3 , and LaFeO 3 towards recharge of Li-0
using electrodes preloaded with Li 20
oxidation of the preloaded Li 20
2
2
2
batteries in
to circumvent parasitic discharge products. Effective
was observed redundantly by XRD, SEM and electrolysis.
Compared to uncatalyzed carbon and the other perovskites, LaCrO 3 and BaO. 5 SrO.5 Coo.8 Feo.2 0
3
display much enhanced oxidation of Li 2O2 . In particular, the area-specific activity of LaCrO 3
equaled that of highly active nanodispersed noble metal catalysts Pt/C and Ru/C reported. Pt/C
and Ru/C were reported with hundredfold gravimetric currents compared to uncatalyzed Vulcan
carbon. Interestingly, LaCrO 3 having two orders of magnitude lower activity in H2 0-oxidation
showed approximately 3-4 times higher activity towards Li 20 2-oxidation compared to BSCF,
one the highest activity perovskite in H2 0-oxidation. This strong inversion in activity trend
underscores major divergences in the OER process from H20-oxidation to Li20 2-oxidation. We
postulate that catalysis during Li 20 2 -oxidation relies on the modification of surface chemical
stoichiometry of the reactant at the interface of the catalyst with Li 20 2 ; oxygen rich surface
orientations of Li 2.mO2 are shown to have a lower oxidation potential from theoretical studies.
Available cyclic voltammetry data indicate little change in onset potential regardless of the
catalyst used but enhanced oxidation beyond the onset. We speculate that changing relative
distributions of surface orientations on the Wulff shape of the Li 20
2
particles in presence of the
so-called catalyst would explain the invariant onset followed by enhanced oxidation at higher
potentials. In the following section, we present a preliminary study on the thermophysical
properties of Li 2 0 2 and Li 20, the two possible products in a Li-0 2 cell to address thermal safety.
56
57
4 Thermal Stability Studies of Li 2 O 2 and Li 2 0: Combined
In Situ XRD, DFT, and XPS Studies upon Heating
4.1 Crystal structure of Li 20 2 and Li 20
Li 20 2 crystallizes with a hexagonal symmetry in the space group P6 3 /mmc. 66 Along the caxis, the unit cell can be described by stacks of lithium planes followed by oxygen planes. The
unit cell can be described as a A(Li)B(0)C(Li)B(O)A(Li)C(O)B(Li)C(O)A(Li) stack (Figure 19
a).67 Li20 adheres to a cubic antifluorite (face centered cubic) lattice, with an O-Li motif (Figure
19b). The crystal lattice is a stacking of A(O) planes followed by B(Li). The unit cell is a simple
A(O)B(Li)A(O)B(Li)A(O) stack. The thermophysical properties of Li 2 0 2 , which undergoes
significant structural and chemical changes at elevated temperatures, is of scientific interest.
(a)
(b)
Figure 19: (a) Primitive hexagonal unit cell of Li 20 2 (space group P6 3/mmc 66). (b) Primitive face
centered cubic unit cell of Li 2 0 (space group Fm3m 68)
58
4.2 Previous studies of the thermal transformation of Li 2 02
The thermal stability and structural evolutions of Li 2 0 2 at elevated temperatures have
been investigated by ex situ X-ray diffraction (XRD) and various thermal analyses.si,6 9 7 2 The
thermal decomposition temperatures obtained from these studies are summarized in Table 5.
Environment
Author
et O
Tsentsiper
al. 5 '
Tnenov
Li 2O2 Decomp. Temp. (*C)
xygen/Nitrogen
Investigation method
>270
Thermostatic heating,
Differential
manometer, KMnO 4
titration
~297
Adiabatic scanning
calorimetry
/Vacuum
a.
Tanifuji et al.7
Vacuum/Argon
Table 5: Literature data of Li 2 0 2 phase transformation under various heating conditions
Rode et al.6 9 reported that the thermal decomposition of Li 20 2 occurs around 315 'C under dry
air (283 'C under vacuum) via progressive loss of oxygen to reach Li2 Oi.0 5 at 375 'C as
determined by chemical titration. The thermal decomposition temperature of Li 20
2
is consistent
across a number of studies; Li 2 0 2 is found to decompose to Li 2 0 above 300 'C using
thermogravimetric analysis (TGA) and differential thermal analysis in air, vacuum, and inert
- . In addition, Tsentsiper et al.5 1 suggested that the decomposition of Li 2O2 at
atmospheres70 72
fixed temperatures between 270 *C and 320 'C under nitrogen, oxygen, or vacuum proceeds via a
59
solid solution of Li 2 0 2 -Li 2 O (by ex situ XRD and KMnO 4 titration) that shifts in favor of Li 20
only after
~50%
of the Li 20 2 has decomposed.
While the thermal decomposition
at elevated temperatures5,69-72 and surface reactivity 2 1,3 1 of Li 20
2
and Li 2 0
have been
investigated by ex situ characterizations, little is known about the real time changes in the bulk
structure and surface chemistries of Li 20 2 and Li 20 as a function of temperature. Such a study is
needed to understand the stability of Li 20
2
and Li 20 and their decomposition mechanisms in
order to devise strategies to enhance the safety characteristics of rechargeable Li-0 2 batteries.
In this thesis, we report the crystallographic and surface chemistry evolution of Li 2 0 2 and
Li 2 0 at elevated temperatures using in situ techniques including Xray diffraction (XRD),
thermogravimetric analysis (TGA), X-ray photoelectron spectroscopy (XPS), and density
functional theory (DFT). A significant decrease in the lattice parameters of Li 2 0 2 between 280
'C and 310 'C was noted and attributed to temperature-induced oxygen defects formation within
Li 2 0 2 prior to its conversion to Li 2 0. The formation of oxygen defects in Li 20 2 upon heating is
further supported by DFT calculation. In situ XPS studies of Li 2 0 2 as a function of temperature
revealed a change in surface chemistry to Li 2 0 beginning at 250 'C along with the appearance of
Li 2 CO 3 , which is attributed to a reaction between surface Li 20 2 and hydrocarbon species. The
implication of the evident chemical reactivity between Li 2 0 2 and carbon will be discussed in the
context of carbon-based Li-0
between Li 20
2
2
electrodes. The implication of the evident chemical reactivity
and carbon revealed here for Li-0
discussed.
60
2
batteries using carbon-based electrodes is
61
4.3 Experimental
4.3.1 In situ X-ray diffraction (XRD)
Lithium peroxide (Li 20 2 , Alfa Aesar, purity greater than 90%), and lithium oxide (Li 2 0,
Alfa Aesar, purity: 99.5% metal basis) were used as received. All powders were stored in an
argon glovebox with moisture and oxygen content below 0.1 ppm until used. The starting Li 2 02
powder sample contains LiOH and Li 2 CO 3 impurities estimated to ~13% and -1%, respectively,
from room temperature XRD data. The thermal stability of Li 20
using a PANanalytical X'Pert Pro
2
and Li 2 0 were investigated
X-ray diffractometer, in Bragg-Brentano geometry,
equipped with an Anton Paar* HTK-1200N environmental and temperature control stage.
Samples were mounted in a corundum (A12 0 3 ) sample crucible under ~10-3 mbar vacuum. X-ray
profiles were collected in two separate experiments for 37 minutes each: the first, from 25 C to
700 C (sample temperature) at 50 C steps (designated "coarse XRD") and the second, from 200
C to 400 'C at 10 C steps (designated "fine XRD"). The time-averaged heating rates were 1
C/min for the coarse in situ XRD and 0.5 C/min for the fine XRD. The diffractometer was
configured to 45 kV and 40 mA using copper wavelength (Ka
= 1.5405980
patterns were recorded in continuous scanning mode with a step size of 0.0167
0
A). Diffraction
from 15 to 72
in 20. Exposure of samples to air prior to reaching 1 mbar vacuum was below 5 minutes. The
main impurity phase, LiOH, persisted up 300 *C but became barely detectable beyond 350 'C.
These impurity phases are not expected to influence the crystallographic changes in the main
Li 20
2
phase under investigation. The lattice parameters of Li 2 0 2 (space group P6 3 /mmc) were
extracted by profile fitting of the twelve reflections available in the scanned range with emphasis
on the goodness of fit of the (002), (100), (101), (102), (004), (103), (110), (104), (102), (200),
62
(201) and (105) reflections. Those of Li 2 0 (space group Fm3m) were extracted from (111),
(200), (220), (311) and (222) reflections. Lattice parameters extraction was performed using the
FullProf software, 73 and confirmed using Highscore Plus X-ray analysis software.
4.3.2 In situ X-ray photoelectron spectroscopy (XPS)
XPS spectra of Li 20 2 , Li 20, LiOH (99.95%, monohydrate, Aldrich), Li 2CO 3 (90%, Alfa
Aesar), and discharged Li-0 2 electrodes were analyzed using a Physical Electronics model 5400
X-ray photoelectron spectrometer. Data were collected using a non-monochromatic Al K,
(1486.6 eV) X-ray source operating at 400 W (15 kV and 27 mA) and ~5- 10-8 mTorr vacuum.
The X-ray source was located at 54.7 relative to the analyzer axis, and samples were analyzed at
an electron takeoff angle of 45'. For Li2 0 2 , Li 2 0, and Li 2 CO 3 , about 300 mg of powder was
loaded into a die in an argon-filled glovebox, and then removed from the glovebox and pressed
in a dry room (<0.05% relative humidity) to form a pellet. The pellet was then transferred from
the dry room to the introduction chamber of the XPS without further exposure to ambient
conditions. LiOH was spread on electrically conductive carbon tape in the argon glovebox.
Multiplex spectra of various photoemission lines were collected at high resolution using
an analyzer pass energy of 22.35 eV, an increment of 0.1 eV/step, and an integration interval of
50 ms/step. To compensate for sample charging effects, all spectra of LiOH, Li 2 CO 3 , Li 2 0 2 and
Li 2 0 were calibrated with the C Is photoemission peak for adventitious hydrocarbons at 285.0
eV,74 while spectra for the discharged VC electrode were calibrated to 284.6 eV 75' 76 considering
contributions from adventitious hydrocarbon at 285.0 eV and carbon black at 284.4 eV.7 5 XPS
data were collected from Li 20
2
and Li 20 at room temperature and during heating in the XPS
chamber at 150, 250, 300, 350, 400, 450 and 500 C. XPS data were collected after sample
outgassing was minimized and the analysis chamber pressure stabilized. Curve fitting and atomic
63
percentage analysis for chemical components in the Li Is, C Is and 0 Is regions were performed
with CasaXPS analysis software, using a Shirley background and Gaussian-Lorentzian curve to
fit each line-shape. Full widths at half maxima were constrained to be less than 2 eV for all
components. Relative sensitivity factors used for Li Is, C Is and 0 Is were 0.028, 0.314 and
0.733 respectively. The standard deviation in atomic percentage assignments was set to 5% for
Li 20 2 , and calculated according to the standard formula for Li 2 0, since two heating experiments
were performed.
4.3.3 Density functional theory (DFT) calculations
Spin polarized calculations were performed with the Vienna Ab-initio Simulation
Package (VASP)
77' 78
using Density Functional Theory (DFT) and the Projector-Augmented
plane-Wave (PAW) method 79' 80. Exchange-correlation was treated in the Perdew-BurkeErnzerhof (PBE)8 1 generalized gradient approximation (GGA) with electronic configurations of
2s and 2s2p for valence states of Li and 0, respectively. An energy cutoff of 450 eV was chosen
for Li 20
2
and Li 2 0 bulk geometric optimizations, and full relaxation (both lattice vectors and ion
coordinates) was performed with a force tolerance of 0.02 eV/A or less.
For Li 20
2
bulk calculations, the Brillouin zone was sampled by Monkhorst-Pack k-point
meshes of 5x5x2, 3x3x3, and 2x2x 1 for the 8-atom hexagonal unit cell 8 2, 32-atom, and 64-atom
supercells, respectively. Relaxed lattice constants of the 8-atom primitive unit cell are in
agreement with Cota et al. (a=3.183 A and c =7.726 A).66 The simulated supercells are created
from linear combination of the 8-atom primitive unit cell lattice vectors as listed in Table 6 and
depicted in Figure 20. Oxygen vacancy calculations were performed by removing an 0 atom
from the perfect bulk in 32-atom, and 64-atom supercells, which corresponds to defect
concentrations of 1/16, and 1/32, respectively.
64
8-atom unit cell
MA
aL
C0
32-atom supercell
64-atom supercell
Figure 20: Atomic supercells of Li 20 2 used in DFT-simulation of Li 2 0 2. Oxygen and lithium
ions are represented by red and green spheres respectively. The blue sphere depicts the location
of introduced oxygen vacancies during simulation.
For Li 2O bulk calculations the Brillouin zone was sampled by Monkhorst-Pack k-point
meshes of 4x4x4, 3x3x2, and 2x2x2 for the 12-atom Fm3munit cell, 48-atom, and 96-atom
supercells, respectively. The relaxed lattice constant of the 12-atom Li 2O primitive unit cell is in
agreement with Islam et al.'s DFT study (a(Li 20)=4.65
A
in this work vs. a(Li 2O)=4.64
A in
Islam et al.'s work), where the calculated DFT lattice constant of Li 2O is slightly larger than the
experimental values (4.61 A at room temperature and 4.57
A
extrapolated to 0 K).
Lithium
vacancy calculations were performed by removing single or multiple Li atoms from the perfect
65
94
Allk
7 ow*000
0
O0
D
S D~00
4
12 atom unit cell (Li 20)
(Li vac. conc. = 0)
000
12 atom unit cell
(Li vac. conc. = 2/8 = 1/4)
%000
ob.4
10gr 000 00 000
0%
000
0o0000
II
II
48 atom unit cell
48 atom unit cell
(Li vac. conc. = 4/32 = 1/8)
(Li vac. conc. = 2/32 = 1/16)
IL 'I
F
96 atom unit cell
96 atom unit cell
(Li vac. conc. = 2/64 = 1/32)
(Li vac. conc. = 1/64)
Figure 21: Atomic supercells of Li 2 0 used in DFT-simulation. Oxygen and lithium ions are
represented by red and green spheres respectively. The black sphere depicts the location of
introduced Li vacancies during simulation.
bulk, 12 atom, 48-atom, and 96-atom supercells with Li-vacancy concentration of 1/64, 1/32,
1/16, 1/8, and 1/4, for obtaining the relationship of lattice constant change vs. Li vacancy
concentration. Graphical representations of Li 2 0 supercells used in our simulation are presented
in Figure 21.
66
Lattice vectors
in the Cartesianvectorform
Li 2 O 2
Li 2 O
8-atom unit cell
12-atom unit cell
d5 = [ a,
0,
0]
dp2
[ 0, a(Li2)
=-a
ao63S2
bp=[2 ,2a,0p=
2)0
jp 2 =
c]
0,
SP= [ 0,
=2aa
p1
Li2O2
32-atom cell
pl
i202
_
dLi2O
48-atom cell
Li2O
4 8-atom
64-atom supercell
p2
b p2
p2
p2
__
cell
__
-LiO
c 2
48-atom cell
Li2O2
C
=C
pl
32-atom cell
Supercell lattice vectors
[ 0, 0, a(Li 2 0)]
48-atom supercell
32-atom supercell
32-atom cell
[ a(Li 20), 0, 0]
=
p2
96-atom supercell
-Li2O2
0i2
64-atom cell
Li202
64-atom cell
p
p
p
64-atom cell
Li 0
96-atom cell
p2
p2
96-atom cell
p
Table 6: Lattice vectors of the simulated Li 2O2 and Li 20
2
p2
Li2O
LiO2
(ip
96-atom cell
supercells. d , b,, and JP1
, bp 2 , and Jp2) are the lattice vectors of the Li 2 0 2 8-atom (Li 2 0 12-atom) primitive cell
represented with the lattice parameters of a = 3.183 A and c = 7.726 A (a(Li 20) = 4.65
Cartesian vector form.
32-atom cell'
0
, 2 2
3-ao cell' I
32-atom cell
and
5
,
64-atom cell'I
b 64-atom cell'I
2
64-atom cell
D2O cell'
2
lattice vectors of the Li 2 0 2 32-atom and 64-atom supercells, and d 48-atom
celI 48-atom
A) in the
are the
- 2O
cell
C48-atom
are the Li 20 lattice vectors of the 48-atom and 96-atom
C2
,b
and 5
a da96-atom cell' b 6ao cell'I 96-atom cell
supercells, respectively. All the supercells are constructed by linear combination of the primitive
cell lattice vectors.
67
4.3.4 Thermogravimetric analysis (TGA)
Thermogravimetric analysis was performed on an 8.9 mg sample of Li 2O2 at a rate of 1
'C/min under inert high purity helium (Airgas, UHP300, 99.999%) balance gas (10 mL/min) and
nitrogen (Airgas, 99.999%) sample gas (90 mL/min) using a TA Instruments* Q50
thermogravimetric analyzer. An Li 20
2
sample was mounted in a platinum pan and the analyzer
was allowed to equilibrate for 20 minutes with gas flow prior to data collection. Exposure of the
sample to air prior to insertion in the inert atmosphere was kept below 5 minutes.
68
69
4.4 Results and discussion
4.4.1 In situ X-ray Diffraction
XRD data collected in situ during heating of Li 20 2 are shown in Figure 22 (32' to 360 20range), and profiles in the full 20 range from 30' to 600 are included in Figure 23. At room
temperature, all expected peaks of Li 2O2 are observed in addition to peaks attributable to LiOH.
The LiOH phase is estimated to ~13% of the total material and is not expected to significantly
influence the structural changes reported for the main Li2 O2 phase under study. The crystal
structure of Li 2O2 can be described by a hexagonal unit cell with space group P6 3/mmc, which
has lattice parameters of a = 3.142
A and
c = 7.650 A at room temperature. No significant
changes were observed in the XRD peaks of Li 2 0 2 and LiOH with heating up to 250 *C.
(b)
(a)
03
4000
25C
.=
8
01
5000C
~70000,
R
2800C
3300"C
C250"C
2200C
1
000
32
33
34
20 [Cu Ka]
200"C
2CCD
17
35
36
32
33
35
34
20 [Cu Ka]
36
Figure 22: Temperature dependent phase evolution of Li 2 0 2 investigated by in situ XRD; a) In
situ scanning at 50*C steps from 25*C to 700 *C and back to 25'C (Coarse XRD); b) In situ
scanning at 10'C steps from 200 *C to 400 'C (Fine XRD). Crystallographic peaks are coded by
colors: blue: Li 2 0 2 , red: Li 20 and black: LiOH.
70
0
0
0
x = Li2CO3 or Al 2O3
(
CC4
CN~
JC4i
C4 0
3350
0[u
20
0C'%1
i
455
I
55
0
60
I -
|
-0
LO)
CN'
30
35
40
45
50
55
60
20 [Cu Ka]
Figure 23: Full 30 to 60' 20 range coarse in situ XRD of Li 2O2 at 50'C steps from 25'C to 700*C
and back to 25*C.
All peaks were found to shift to lower angles with increasing temperature as a result of thermal
expansion. The absence of any changes besides the expected shift to lower angles of the peaks
does not support a phase transformation from a-Li 20 2 to P-Li 2 O2 reported previously. 69 Profile
fitting of XRD peak positions of Li 20 2 allowed for the determination of the Li 20 2 lattice
parameters as a function of temperature (Figure 24). Volumetric and lattice parameters thermal
expansion were calculated based on equations (Eq 9) and (Eq 10). The volumetric thermal
expansion coefficient of Li 2 0 2 was found to be -73
11-10-6 K', and the thermal expansion
coefficients along the a and c axes were found to be 35.0±1.9 -10~6 K- (R2 = 0.98544) and
71
29.9±3.6-10-6 K-1 (R2 = 0.93042), respectively. The volumetric thermal expansion coefficient of
Li 2 0 2 was found to be -73±11-10-6 K-1, and the thermal expansion coefficients along the a and c
axes were found to be 35.0±1.9 -10-6 K-' (R2
0.98544) and 29.9±3.6-10~6 K-1 (R2
=
=
0.93042),
respectively. It is interesting to note that the volumetric thermal expansion coefficient found in
this study is considerably larger than the values of 40 to 50- 106 K- obtained from previous abinitio calculations. 67
av
1
dV
= V(RT ) dT
ca =
(Eq. 9)
1
da
a(RT) dT
(Eq. 10)
S in Li20 2 -
0.0 0.1 0.2
7.75 (a)
CU2
7.65
. 4.70
a1 Q
.4.65E
2.42
7.60
3.16.
2.43 (
4.6
4.6
100 200 300 400 500 600 700 4.55
Temp. [*C]
L2 2..e-experim.
c/a vs. temp.
0 DFT c/a vs.0-vacancy
2.41180 200 220 240 260 280 300 320
Temp. [OC]
Figure 24: Evolution of species lattice parameters during thermal decomposition of Li 20 2. a) a
and c parameters of Li 2O2 and Li 2 0 during heating from room temperature to 700 "C. Lattice
parameters of Li 20 2 between 200 "C and 300 *C (square symbol) are collected in a separate
experiment at 10 'C temperature steps (fine XRD). Notice the lattice shrinkage of Li 2O2 from its
expanded state in the vicinity of the phase transformation temperature (-280*C). b) Experimental
c/a ratios in Li 2O2 during thermal treatment compared to DFT-simulated c/a ratios as a function
of oxygen vacancies. Trend match in c/a ratios from experiment and DFT suggests that thermal
treatment of Li 2O2 results in formation of oxygen vacancies between 280 *C and 310 'C prior to
phase change. Comparison of DFT andXRD-extracted lattice parametersare strictly interpreted
in terms of trends and not quantitatively.
72
The volumetric thermal expansion coefficient found in this study is in good agreement with
values (69- 10~6
80- 10-6 K- between 27 ~ 327 *C, Figure 25b) calculated from the quasi-
harmonic Debye model using the GGA-PBE total energy versus volume relationship fit with the
Birch-Murnaghan equation of state. In contrast, coefficients reported using the local density
approximation (LDA) are underestimated67 . Overall, the fact that the experimental values are
within the GGA and LDA predictions is consistent with general observation that LDA tends to
overestimate bond energies in solids (shorter bond distance and too large bulk moduli), while
84 8 5
GGA provides proper corrections. '
-1.36
80 (b)
78
(a)
-1.38
2
,76
o74
2 -1.40
uw 0)0
9-1.42
-1.44
300
e 72
70
68
500
400
Volume (Bohr3 )
600
0
100
200
Temp. [oC]
400
300
Figure 25: (a) Calculated Li 2 0 2 total energy vs. unit cell volume with GGA-PBE in this work.
The corresponding bulk modulus and the first pressure derivative of the bulk modulus of Li 20 2 at
zero pressure fit with third-order Birch-Murnaghan equations of state are 71.12 and 4.30 GPa
respectively. (b) Calculated volumetric thermal expansion coefficients between T=27-327 *C at
zero external pressure based on the quasi-harmonic Debye model.6 7 DFT data were obtained in
collaboration with Dr. Yueh-Lin Lee in the electrochemical energy laboratory.
Upon heating to 300 'C and higher, the XRD patterns changed significantly and new
peaks appeared and grew in intensity at the expense of Li 20 2 peaks with increasing temperature
(Figure 22 and Figure 23). These new peaks can be indexed to cubic Li 20 with the strongest
(111) peak at ~340,86 indicating a phase transformation from Li 2 0
02).
2
to Li 2 0 (Li 2 0 2
->
Li 2 0
+
2
In situ XRD data collected at finer temperature increments (fine XRD) show that the onset
temperature of Li 2 0 2 decomposition is in the vicinity of ~280 *C (Figure 22b), which is in
73
agreement with previous work by Rode et al.69 and Tanifuji et al. 72 . The small differences in the
decomposition temperatures of Li 2 0 2 to Li 20 can be attributed to dissimilar heating rates. 72 The
conversion of Li 2O 2 to Li 20 was found complete beyond 350 C (no XRD-detectable peaks of
Li 20 2 ) and Li 20 remained stable until 700 'C. It is interesting to note that the onset and complete
decomposition temperatures of Li 20 2 in vacuum are considerably lower than those in air reported
previously (having an onset of~315 C by Rode et al.6 9 and ~325 C by Vol'nov 70 , Table 6).
The lattice parameters of hexagonal Li 20 2 were found to decrease prior to decomposition
to Li 20 from 250 'C to 300 'C (Figure 24), which can be attributed to the loss of oxygen from
Li 2 0 2 resulting in lattice shrinkage. To provide further insights into the physical origin of the
lattice contraction, the ratio of the c to a lattice parameter is shown as a function of temperature
in Figure 24b. This c/a ratio and the unit cell volume change are compared with weight loss from
a Li 2 0 2 sample in a thermogravimetric measurement (Figure 26). The sharp reduction in the c to
a ratio at 280 C coincides with the onset of weight loss, which suggests a change in the
stoichiometry of Li 20
2
associated with the release of oxygen from Li 2 0 2.
022-
peroxide groups
are arranged along the c-direction in the Li 2O2 structure (Figure 19a). Oxygen loss would lead to
the formation of oxygen 02- species with ionic radius (-1.38 A in a 4-fold coordination 8 7)
smaller than the 0-0 bond distance reported for lithium peroxide (do=o ~ 1.55 A);66 this situation
would explain the contraction along the c-direction of the peroxide. This hypothesis is supported
by our DFT studies of the equilibrium crystal structures of oxygen-deficient lithium peroxide
(Li 2 02-6 ), where increasing oxygen deficiency (increasing 6 in Li2O2-s) reduces the c/a ratio, as
shown in Figure 24b. A similar trend match is found for the experimental lattice volume
contraction compared to the computed volume (Figure 26).
74
0.00
8 in Li2 O 2
0.05
0.10
0.15
0.8
0
0.6 o
1Q
0.4g
-o- Expejimental AVN
-2 vs. Temperature
o DFT IVN vs 8
-3
0.2
I-
0.0
270 280 290 300 310 320
Temp. [OC]
Figure 26: Experimental lattice volume change in Li 2 0
2
(referenced to the expanded lattice
volume at 280 'C) during thermal treatment compared to DFT-simulated lattice volume change
as a function of oxygen vacancies. In the same fashion as the experimental and calculated c/a
ratios, a good trend match is found between the experimental and calculated volume changes that
confirm the formation of oxygen vacancies in Li2 0 2 close to the conversion temperature. DFT
data were obtained in collaboration with Dr. Yueh-Lin Lee in the electrochemical energy
laboratory.
In situ XRD profiles and the lattice parameters of commercial Li 20 (space group Fm3m)
as a function of temperature are presented in Figure 27a. The full pattern in the 20 range of 300 to
700 is displayed in Figure 28. Again, very minor peaks attributable to LiOH are observed. The
presence of LiOH on all oxides of lithium studied here is a consequence of their reactivity with
trace moisture that could not be avoided during sample loading in the X-ray spectrometer. No
crystallographic change was observed for Li 2 0 upon heating to 700 'C. A linear thermal
expansion coefficient of -29.2±0.8- 10-6 K- was found for Li2 0 between room temperature and
700 0C, in reasonable agreement with the findings of Hull 8 3 and Kurasawa 8 (33.6+0.8-10-6 K-).
75
25*C
500 OC
70000
200"C
60*C
I25*C
32
4.72
33
55
34
28 [Cu Ka]
56
57
(b)
4.70
E 4.68
Commercial
Li 0 heated
a- 4.66
S4.64.
S4Li20
4.60 -
from phase
transf. of Li 0
2
2
- - - --_
0 100 200 300 400 500 600 700
Temp. [*C]
Figure 27: Temperature dependent phase and lattice evolution in Li 20 investigated by in situ
XRD. a) Li2 0 remains stable, experiencing a fairly linear thermal expansion (29.2-10- K-) from
25 0C to 700 *C (commercial Li 2 0). b) The lattice parameters of Li 20 from phase transformation
of Li 2 0 2 remained noticeably lower than those of commercial Li 20 between 300 0C and 550 'C
suggesting a lithium-deficient Li 2.0 phase (see Figure 29 for DFT calculated lattice parameters
of Li-deficient Li 2 O) that progressively reaches Li2 0-stoichiometry with increasing temperature.
Error bars are the reliability factors generated by the Fullprof software reflecting uncertainty in
the profile fitting.
76
0
00
NC
2
[CuKa
0
A
0
LO
30
35
40
45
50
55
20 [Cu Ka]
60
65
70
Figure 28: Full 30 to 700 20 range coarse in situ XRD scans of Li 2 0 at 50*C steps from 250C to
700*C and back to 25'C.
Interestingly, the lattice parameters of Li 20 obtained from decomposition of Li 20 2 were
found to be consistently lower than those of commercial Li 20 between 300 0C and 500 'C
(Figure 27b). This difference gradually disappeared at 500 C and higher. As reported in
previous theoretical and experimental works8 9 ~91, the predominant defects in Li 20 are cation
point defects (either Li vacancy or Li Frenkel-pair), which indicates that Li 20 transformed
initially from Li 20 2 might be lithium deficient relative to Li 20 and gradually approaches the
equilibrium Li 2 O stoichiometry and structure upon further heating. This postulate is substantiated
by DFT simulation of a lithium-defective Li 2 0 phase presented in Figure 29. The behavior of the
77
lattice parameter of the nascent Li 20 formed from Li 20 2 at ~300 *C and above is found to follow
that of the calculated Li 2 -60 where decreasing lithium defects relative to oxygen results in the a
parameter slowly increasing towards its stoichiometric value. The disappearance of the lattice
parameter difference between the commercial Li 20 and Li 2O from conversion of Li 20 2 above
500 'C suggests annihilation of Li vacancies upon heating through further oxygen loss in such a
temperature range. The synergistic use of in situ XRD, DFT and TGA presented here provides
structural information for the Li 2 0-Li 2O 2 solid solution suggested previously.51
Li-defect in Li2 0 (6)from DFT
0.6
0.4
0.2
0.0
1.000
Li20 from phase
0
transf. of Li20
0.996
OO6
Q
0.992
o 0.988
0.984
250 300 350 400 450 500 550
Temp. [*C]
Figure 29: Dark blue: ratio of lattice parameters of Li 20 formed from Li 20 2 (adefect as a function
of temperature) decomposition to those of commercial Li 20 (astoich as a function of temperature).
Light blue: ratio of DFT-simulated lattice parameters of Li-defective Li 20 to those of DFTsimulated stoichiometric Li 2 0. A good trend match these ratios is found that supports the
postulate of a lithium deficient phase of Li 20 being formed at initial stages of decomposition of
Li 2O2 to Li 20, which gradually becomes stoichiometric. DFT data were obtained in collaboration
with Dr. Yueh-Lin Lee in the electrochemical energy laboratory.
4.4.2 In situ XPS analysis of Li 2 0 2 and Li 2 O
4.4.2.1 XPS analysis of Li 20 2 , Li2 0, LiOH and Li2CO 3 at Room-Temperature
The surface chemistries of lithium-oxygen reference compounds were examined using
XPS at room temperature. Figure 30 shows 0 is, Li Is and C Is XPS spectra of reference Li 2 0 2 ,
78
Li 2 0, LiOH, and Li 2 CO 3 at room temperature. The main 0 Is photoemission peak from the
Li 2 0 2 and Li 2 CO 3 samples are located at 531.2 eV and 532.1 eV, respectively, which is in
reasonable agreement with reported binding energies for Li 20 2 (531.5 eV) 9 2 and Li 2 CO 3 (532.2
eV).7 The Li Is peak of Li 2 02at 54.5 eV and Li 2 CO 3 at 55.4 eV are in agreement with reported
values for Li 20 2 (~54.7 eV) 93 and Li 2 CO 3 (55.4 eV). 75 The main 0 1s peak from Li 2 0 is centered
at 531.1 eV. This binding energy is not in agreement with values reported for Li 20, neither at
low temperatures (between 530.0 and 530.5 eV 94 '95 at -248
temperatures (528.5 eV at 600
oC
96
).
C and -238 *C ) nor at high
Tanaka et al. 96 have reported surface coverage of Li 2 0 by
LiOH in the event of moisture exposure with XPS peak at 531.0 eV. This peak assignment is
further supported by the presence of saddle peaks at 20.48' and 32.58' on the XRD profiles of
Li 20 at room temperature (Figure 28). Interestingly, the binding energy of the main 0 Is
photoemission peak for the LiOH at 532.0 eV is similar to that of Li 2 CO 3 but significantly
deviates from previously reported values for LiOH (531.096 - 531.5 eV9 7). It is likely that the
surface of the LiOH sample is covered by Li 2CO 3 given the similarities between photoemission
peaks for Li 2 CO 3 and LiOH. Peaks attributable to Li 2 CO 3 are also observed during XRD of
LiOH at room temperature. This hypothesis is further supported by the Li Is spectra, where the
binding energy of the LiOH at 55.4 eV agrees well with that of Li 2 CO 3. 75
79
CO
3
C-C
C 1s
LiOH
01S
LiOH
Li is
LiOH
LiOH
LiOH
L 2O
L 2 CO 3
L1iCO 3
3
U1202
292 290 288 286 284 282 280
Binding Energy (eV]
i2 2
U_____________L202
Li 2
536
528
534 532 530
Binding Energy (eV]
526
58 57 56 55 54 53 52 51
50
Binding Energy [eV]
Figure 30: XPS spectra of lithium-oxygen compounds used in this study. The dotted red lines
96 97
mark the position of the main LiOH peak in Li Is and 0 Is according to literature references. ,
The LiOH material is covered by Li 2 CO 3 and therefore appears shifted in the 0 1s and Li Is
photoemission regions compared to the literature values. XPS data were obtained in collaboration
with David Kwabi and Dr. Yi-Chun Lu in the MIT electrochemical energy laboratory.
4.4.2.2 In situ XPS analysis of Li 2 0
2
at Elevated Temperatures
The changes in the surface chemistry of the Li 20 2 powder at elevated temperatures were
examined via in situ XPS as a function of temperature. Formation of Li 20 and Li 2 CO 3 on the
surface of Li 20 2 was clearly noted upon heating the Li 20 2 pellet to 300 'C and higher. Figures
31 a-c show the C Is, 0 1s and Li Is spectra of the Li2 0 2 pellet upon heating in the XPS chamber
from room temperature to 500 'C, and Figures 3 1d-f show atomic percentages of components
within each region as a function of temperature.
80
CO
(a)
Li 2O2/LiOH
Li2C0 31
2
(b)
i,
0
\%-40
LiO2 LiOH
LiCO,| LiO
Li2O O s
Li Is
(c)
*C
350 *C
300 *C
250 *C
150 *C
RT
292 290 288 286 284 282 280
Binding Energy [eV]
8.
-2
-6
4'L
Li2
50 (e)
C is
10 (d)
,40
c-c
/LiOH
Li20/LiOH
(f)
O 1s
Li Is
22
240.
30
O2
Li2CO
2CO320
Li2O
.0 20'
10 Lii2C
2CO,
2Li2
0
0 r0.
0
100 200 300 400 500
Temperature rC]
0
100 200 3 400 500
Temperature [*C]
0
100 200 300 400 500
Temperature [*C]
Figure 31: Temperature dependent in situ XPS spectra of Li 2 0 2 pellet in (a) C 1s, (b) Ols, and
(c) Li is regions. Atomic percentage of C, 0, and Li in Li 2 0 2 , Li 2CO 3 and Li 20 obtained from
quantitative component analysis of (d) C Is, (e) Ols, and (f) Li Is spectra. XPS data were
obtained in collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT electrochemical
energy laboratory.
4.4.2.2.1 C is region
The spectra show two different components. A peak of aliphatic chains at 285.0 eV,
which can be assigned to the hydrocarbon on the surface of Li 2 0 2 , was found to decrease with
increasing temperature (Figure 3 1a and Figure 3 1d). The intensity reduction in the hydrocarbon
peak is accompanied by the growth of a peak at 290.0 eV as temperature is increased. This new
81
peak at 290.0 eV can be attributed to the CO 3 2- groups in Li 2 CO 3. 75 The simultaneous reduction
of hydrocarbons and increase in Li 2 CO 3 can be explained by chemical reactions of adventitious
carbon species with surface Li 2 0 2 below 250 'C and with the Li 2 0 2 -Li 2 O solid solution formed
at higher temperatures (Figures 3 1e-f).
4.4.2.2.2 0 is region
The main photoemission peak at room temperature is at 531.2 eV, which is consistent
with the presence of surface Li 2 0 2 , surface LiOH, or a mixture of Li2 0 2 and LiOH on the
surface. 92 ,94 ,96 ,97 Upon heating to 250 C and higher, the 0 is signal broadened and split into two
new peaks centered at 532.0 and 528.6 eV, which can be attributed to the formation and growth
of Li 2 CO 3 75 and Li 2 0,
96
respectively. The onset of phase transformation to Li 2 0 observed at 250
C on the surface of the Li 2 0 2 is shown clearly by the component analysis in Figure 3 1e, which
is in good agreement with the decomposition temperature from in situ heating XRD (Figure 22b).
At 350 C, only ~13% Li 2 0 2 remains from both the Li Is and 0 Is spectra, indicating almost
complete
decomposition of Li 2 0 2 to Li 2 0.
This is in reasonable agreement with the
disappearance of Li2 0 2 XRD-peaks past 350 C (Figure 23).
4.4.2.2.3 Li is region
As Li Is binding energies of LiOH and Li 2 0
2
overlap, 96 -98 the room temperature Li Is
peak at 54.5 eV may correspond to surface Li 2 0 2 , surface LiOH, or a mixture of Li 2 0 2 and LiOH
on the surface. Upon heating to 250 C, the Li Is signal was found to considerably broaden and
split into two peaks centered at 53.6 eV and 55.4 eV at 500 C, in agreement with the 0 Is
spectra. The lower binding energy peak at 53.6 eV can be assigned to Li 2 0 68 while the higher
binding energy component corresponds to Li 2 CO 3 .75 Onset of transformation to Li 2 0 is observed
as early as 250 'C on the surface of Li 2 0 2 as shown by component quantification (Figure 31 e).
82
This early onset coincides with saddle decay in the lattice parameters extracted from fine in situ
XRD (Figure 24a, fine XRD). The formation of carbonate-type byproducts such as alkyl
carbonates and Li 2 CO 3 in Li-0
batteries with carbon-based oxygen electrodes has been well
2
documented. 2 1,3 1,3 2 ,9 9 In particular, McCloskey et al. 3 1 showed, using isotope labeling in a Li-0
2
cell, that a monolayer of Li 2 CO 3 forms at the interface of the cathode carbon matrix and the
electrochemically deposited Li 20 2 . Such reaction is predictable from the thermodynamically
31 38
favorable chemical reactions displayed below. '
Li 2 0 2 + C + 02 # Li 2 CO 3
AG = -542.4 k]/mol
(Reaction 5)
AG = -166.5 kj/mol
(Reaction 6)
Li 2 02 + CO 2
LizCO3 + 2 Oz
Li 2 02 + CO
Li 2 CO 3
AG = -423.7 kj/mol
(Reaction 7)
2Li 2 0 2 + C
Li 2 0 + Li 2 CO 3
AG = -533.6 kj/mol
(Reaction 8)
3.0
2.8-~2.6
'Z
2.4 0.
2.22.00
400
800
1200
1600
Charge [mAh gCarbon
Figure 32: Discharge curve for Vulcan carbon electrode discharged at 10 mA/gabon in 0.1 M
LiClO 4 in DME to ~1600 mAh/garbon. XPS spectra of this electrode is presented in Figure 33.
XRD profile of a similarly discharged vulcan carbon electrode is reported by Lu et al.5 and
shows the presence of Li 20 2. XRD does not probe the thin surface Li 2 CO 3 layer present in this
discharged electrode.
83
C is
LUIs
O2
2ULi
at 350*C
Li202 at 350 *C
Li20 at 350 *C
292 290 288 286 284 282 280
Binding Energy [eV]
536
7
534
532 530
528
Binding Energy [eVJ
526
59 5857 5 5554535251 50
Binding Energy [eV]
Figure 33: XPS spectra comparing electrochemically formed Li202 (in a discharged carbon Li-0 2
electrode) to Li 2 0 2 heated to 350 C. Carbonate formation via chemical in presence of carboncontaining species likely contributes to the formation of Li 2CO 3 . XPS data were obtained in
collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT electrochemical energy
laboratory.
Our in situ heating XPS results lends further support to the hypothesis that Li 2CO 3 detected in the
discharged oxygen electrode (Figure 32 and Figure 33) may in part come from the strong
chemical reactivity between Li 20 2 and carbon in Li-0 2 cells.'2 1
31
84
4.4.2.3 In situ XPS analysis of Li 2 0 at elevated temperatures
The C Is, 0 Is, and Li Is XPS spectra of the Li 20 upon heating to 500 0C are shown in
Figures 34a-c and the corresponding atomic percentages of individual components as a function
of temperature are shown in Figures 34d-f.
LiOH
LiCO
' 3,
LiOH
(a)
C is
C03
(b)
Li2CO,,
' '
Ol
0 1S
Li,CO,
500 *C
.450 0C
400 *C
350 *C
300 *C
250 *C
1500C
RT
292 290 288 286 284 282 280
58 57 56 55 54 53 52 51 50
Binding Energy [eV]
Binding Energy [eV]
---
5
*-
(d)
C-C
40
4
3
Li2CO3
2
O1s
50 (e)
C1s
U
LiOH
L
0
30
OA
E 20 Li2CO
0
10 :2:
1
100 200 300 400 500
Temperature rC]
0
0
100 200 300 400 500
Temperature [*C]
Temperature [C]
Figure 34: Temperature dependent in situ XPS spectra of Li 20 pellet in (a) C Is, (b) Ols, and (c)
Li Is regions. Atomic percentage of C, 0, and Li in Li 20 2, Li 2CO 3 and Li 2 O obtained from
quantitative component analysis of (d) C Is, (e) Ols, and (f) Li Is spectra. XPS data were
obtained in collaboration with David Kwabi and Dr. Yi-Chun Lu in the MIT electrochemical
energy laboratory.
85
4.4.2.3.1
C Is region
Similar to Li 2 0 2 , the spectra show two different components over the entire temperature
range (Figure 34a). The peak at 285.0 eV corresponding to aliphatic hydrocarbon species on the
surface of the Li 20 was found to decrease with increasing temperature, reaching ~1.3% of total
atomic composition at 500 'C (Figure 34d). In a similar fashion to Li2 0 2 , chemical reactions
between adventitious carbon species and surface Li 20 or LiOH upon heating results in reduction
of hydrocarbon species and simultaneous growth of Li 2 CO 3 (increase in the intensity of the C Is
peak at 290.0 eV). Several pathways exist for Li 2 0 to form Li 2 CO 3 by reaction with carbon
38 00
species, three of which are listed below. "1
Li 2 0 + CO2 # Li2 CO 3
Li 2 0 + CO +
102
Li 2 0 + C + 02
AG = -176.5 kJ/mol
# Li 2 CO 3
AG = -433.7 kj/mol
Li 2 CO 3 AG = -570.9 kJ/mol
(Reaction 9)
(Reaction 10)
(Reaction 11)
4.4.2.3.2 0 is region
The 0 Is spectrum (Figure 34b) shows one component centered at 531.2 eV at room
temperature, which is consistent with LiOH, as discussed previously. Upon heating above 150 'C
or starting at 250 *C , this peak was found to broaden, and split into two peaks at 528.6 and 532
eV, corresponding to Li 20 9 6 and Li 2 CO 3 , 5 respectively. At 300 'C and higher, a systematic shift
to higher binding energies was observed and this shift was also found in the Li Is region, and
discussed in detail below. Furthermore, less than 5% LiOH is detected in both Li Is and 0 Is
spectra, indicating nearly complete decomposition of the LiOH surface layer to reveal Li20.
4.4.2.3.3 Li Is region
At 250 'C, the peak originally centered at -54.5 eV corresponding to LiOH present on
the surface of the Li 2 0 was found to shift to -53.6 eV. This shift can be attributed to the removal
86
of surface LiOH and appearance of Li 2 0 on the surface. Upon heating to 300 C and higher, a
systematic shift of this component to higher binding energies was observed, resulting in a
spectrum at 500 C with the main component at ~55 eV (0.9 eV higher than at 250 0C). At 500
0C,
the peak positions of the components originally assigned to Li 2 O and Li 2 CO 3 cannot be
assigned to any of the three main compounds under investigation by XPS and XRD, or literature
values for other lithium-oxygen compounds. Since no structural changes were detected by XRD
upon heating of Li 2 0 (Figure 27a), we hypothesize that these shifts arise due to differential
charging as more resistive Li2 0 begins to replace surface hydrocarbons and dominate the surface
above 300 *C. This hypothesis is supported by the fact that systematic shifts in the binding
energy of Li2 0 become negligible when the Li 2 CO 3 peak at 290.0 eV (XPS spectra collected
from the Li 2 0 at 300 C) were used for binding energy calibration for spectra at temperatures
greater than 300 0C instead of aliphatic carbon at 285.0 eV (Figure 35).
LIOH
58
57
56
55
54
53
52
Binding Energy [eV]
51
50
292
290
LIOH
Li2CO
2
3
CI
CO
LiIs
LICO
LiO
Li 3g:
U2?
288
286
2482280
Binding Energy [eV]
536
534
532
Li0
530
01
528
526
Binding Energy [eVA
Figure 35: Temperature dependent in situ XPS of Li2O pellet calibrated to Li2CO3 peak at 290.0
eV above 300 C. XPS data were obtained in collaboration with David Kwabi and Dr. Yi-Chun
Lu in the MIT electrochemical energy laboratory.
87
The rate of growth of Li 2 CO 3 is more pronounced on the surfaces of Li 20 2 compared to
Li 2 0 surfaces although the predicted carbonate forming reactions are thermodynamically favored
on Li 2 0. The increased surface reactivity of Li 2 O2 compared to Li 2 0 can be explained by the
formation of non-stoichiometric and likely more unstable phase (more positive Gibbs free
energies) such as oxygen deficient Li 2 02-6 and lithium deficient Li 2. 6O during thermal
decomposition. Li 2O2 is utilized as CO 2 scrubber in air purification systems in sealed
environments because it is known to readily reacts with ambient CO 2 around 200 C
101,102.
The
extremely low oxygen partial pressure in the XPS chamber may have been a factor in the
apparent resistance of Li 20 to carbonate formation since a large fraction of possible pathways
requires 02 as reactant (Reactions 10 and 11). In the case of Li 2 0 2 , molecular oxygen is required
only in reacting with pure carbon (Reaction 5). It is not clear whether the apparent stability of
Li 20 toward Li 2 CO 3 formation can be realized in Li-0
2
cells. Li-0
2
batteries typically operate at
1 atm oxygen partial pressure; Li 2 CO 3 will likely still form in presence of the carbon matrix
regardless of the discharge product being Li 20.
88
89
4.5 Conclusions
In this study, we examine structural and surface compositional changes of Li 2 0 2 and Li 20
upon heating via in situ XRD and XPS techniques. In situ XRD reveals that the c/a ratio of the
hexagonal unit cell of Li 20 2 decreases significantly in the temperature range from 280 to 310 'C.
The decreased c/a ratio can be attributed to oxygen loss from Li 20
2
forming oxygen-deficient
Li 202 6 , which is supported by DFT calculations. In addition, a lithium-deficient Li 20 phase
appears at 300 'C from thermal decomposition of Li 2 0 2 , which gradually approaches Li 2 0
stoichiometry upon further heating. In contrast, the structure of Li 2 0 was shown to remain stable
upon heating to 700 0C, confirming that the bulk structure of Li 20 is much more stable than that
of Li 2 0 2. In situ XPS results show that Li 20 and Li 2 CO 3 appear and grow on the surface of the
Li 2O2 starting from 250 'C.
The formation of Li 2 CO 3 may result from chemical reactions
between hydrocarbons (in the XPS chamber and absorbed on the Li 2O2/Li 2 0 surfaces) and the
surfaces of Li 20 2/Li 2 O starting at 150 'C. Our study identifies strong surface chemical reactivity
between Li 2 0 2 /Li 2 O and carbon-containing species, which highlights the challenge and
importance of developing stable carbon-based electrodes for rechargeable Li-0
2
batteries. The
decomposition temperatures of Li 2 0 2 and Li 2 0 electrodes with carbon and electrolyte in Li-0
cells might be different from those of Li 20
2
2
and Li 20 shown in this study, which should be
examined in future studies. Our findings on the thermal stability of Li 20
2
and Li 20 provide the
fundamental insights required to make informed decisions in future designs of practical Li-0
2
batteries.
90
91
5 Perpectives
The systematic study of catalysis of Li 20 2 -oxidation oxidation using ABO 3 -type perovskites
has uncovered an intriguing aspect of Li 20
2
decomposition electrochemistry and promotion of
reactions. Indications that "electrocatalysis"
of Li20 2-oxidation may proceed not in the
traditional sense of absorption and desorption of solution-labile intermediates but rather by
structural and/or crystallographic modifications of the reactant provides an opportunity for
important new scientific approaches to boost the recharge kinetics of not only Li-0
2
batteries.
Such approach may later be applied broadly to other metal air and metal sulfur batteries with
crystalline products. More specific to this work, further experimental understanding of the origin
of the unexpected high activity of LaCrO 3 is critical and is underway. It is also imperative to
obtain conclusive experimental and/or computational evidence (and rebuttal thereof) for the
many mechanisms proposed in the literature regarding the role of catalysts in Li-0
2
batteries. We
believe the present work provide yet another critical piece to the puzzle of the Li-0
2
electrochemistry and merits future consideration.
This work has also provided pseudo real time insights into the mechanics of decomposition
of Li 2 0 2 that will potentially serve to predict the behavior of mixtures of Li 2 0
compounds involved in the Li-0
2
2
and other
cell based on chemical reaction principles. However, it is
necessary to extend these experiments to the empirical study of mixtures encountered in the Li02
cell. In particular, Li-0 2 cells currently tested in laboratories are comprised of nanopowders
of carbon, metal or metal oxide catalysts and organic solvents. In light of the reactivity of
peroxides reported by Clark5 0 especially in contact with organics, it will become critical to
ensure the thermal stability of Li-0
2
battery assemblies prior to any industrialization. To that
goal, we propose systematic and exhaustive studies of mixtures Li 2 0 2 and any potential
92
candidate
material
(cathode
packaging...) for use in Li-0
2
structure
material,
catalyst,
electrolyte
salt and solvent,
batteries. Critical extensions to this work would be investigations
of the heat generation and other thermal hazards of mixtures that may appear in Li-0
2
batteries
using tools such as differential thermal scanning calorimetry and the in-operando tools reported
in this thesis.
93
94
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