Cation-Disordered Oxides for
Rechargeable Lithium Battery Cathodes
by
Jinhyuk Lee
B.S. Materials Science and Engineering
Seoul National University (2010)
Submitted to the Department of Materials Science and Engineering
in partial fulfillment of the requirements for the degree of
Doctor of Philosophy
in Materials Science and Engineering
at the
MASSACHUSETTS INSTITUTE OF TECHNOLOGY
September 2015
© Massachusetts Institute of Technology, 2015. All rights reserved.
Signature of author ....................................................................................................
Department of Materials Science and Engineering
July 21, 2015
Certified by ...............................................................................................................
Gerbrand Ceder
R. P. Simmons Professor of Materials Science and Engineering
Thesis Supervisor
Accepted by...............................................................................................................
Donald R. Sadoway
Chair, Departmental Committee on Graduate Students
2
Cation-Disordered Oxides for
Rechargeable Lithium Battery Cathodes
by
Jinhyuk Lee
Submitted to the Department of Materials Science and Engineering
On July 21, 2015, in partial fulfillment of the requirements for the degree of
Doctor of Philosophy in Materials Science and Engineering
Abstract
The demands for high-energy density cathode materials for rechargeable lithium batteries are
ever increasing. This is because such cathode materials will enable smaller and lighter
rechargeable lithium batteries for complex applications such as in electric vehicles or in grid
energy storage. Nearly all high-energy density cathode materials for rechargeable lithium
batteries have been sought from well-ordered close-packed oxides in which lithium and
transition metal ions occupy distinct sites. In contrast, cation-disordered oxides, whose cation
distribution is fully or partially random, have received only a limited attention, because lithium
diffusion tends to be limited by their cation-disordered structure.
In the first part of this thesis, from the study of Li1.211Mo0.467Cr0.3O2, it is demonstrated
that cation-disordered oxides can be promising cathode materials if they contain enough
lithium excess (x >1.09 in LixTM2-xO2). Li1.211Mo0.467Cr0.3O2 forms into a layered structure, but
transforms almost completely to a cation-disordered structure during cycling. While common
wisdoms would expect poor cyclability of this material due to limited lithium diffusion in its
structure, the reversible capacity of this material is remarkably high (~265 mAh/g), which
demonstrates that lithium diffusion can be facile in the cation-disordered structure. Using ab
initio calculations, we show that this counterintuitive behavior is due to percolation of a certain
type of active diffusion channels (0-TM channels) in disordered Li-excess materials, which
becomes more extensive as the lithium-excess level increases.
In the second part of this thesis, a new class of high capacity cation-disordered oxides
(Li-Ni-Ti-Mo oxides) is designed based on the 0-TM percolation theory. As the theory predicts,
the reversible capacity and rate capability in these materials considerably improve with lithium
excess. In particular, Li1.2Ni1/3Ti1/3Mo2/15O2 delivers high capacity and energy density up to
250 mAh/g and 750 Wh/kg at 10 mA/g. Combining in situ X-ray diffraction, electron energy
loss spectroscopy, and X-ray absorption near edge spectroscopy, we investigate the redox
mechanism of the new materials and discuss how oxygen loss with lattice densification can
affect lithium diffusion in the materials by decreasing the lithium-excess level. From these
understandings, strategies for further improvements are proposed, setting new guidelines for
the design of high performance cation-disordered oxides for rechargeable lithium batteries.
Thesis Supervisor: Gerbrand Ceder
Title: R. P. Simmons Professor of Materials Science and Engineering
3
Acknowledgements
First, I would like to thank my thesis advisor, Professor Gerbrand Ceder, for his intellectual and
mental supports. I was very excited and honored to work with him who is not only insightful
but also fearless in conducting research. Thanks to him, I could finally become an independent
researcher who can pursue ideas that can greatly contribute to the materials research society.
I would also like to thank my thesis committee members, Professor Donald Sadoway
and Professor Jeffrey Grossman, for their valuable comments and supports on my research.
Thanks to their generosity and supports, I could successfully complete my doctoral research.
I also thank my collaborators in the CEDER group (Dr. Alexander Urban, Dr. Xin Li,
and Dr. Dong-Hwa Seo) for their continuous support on my research. Without them, none of
my work could have been realized. The entire CEDER group members (Aziz Abdellahi, Rahul
Malik, Stephen Dacek, Sai Gautam, Eric Wang, Rui Wang, Lei Liu, Wenxuan Huang, Daniil
Kitchaev, Ian Matts, William Richards, Ziqin Rong, Wenhao Sun, Alexandra Toumar, Piere
Manuele Canepa, Shinyoung Kang, Jae Chul Kim, Dong-Hwa Seo, Nancy Twu, Alexander
Urban and Kathy Simons) are greatly acknowledged. The greatest thing about the CEDER
group is that you can work with the nicest and the smartest people in the world whom you can
be greatly inspired by. In particular, I thank the Korean members in the CEDER group (Dr. Jae
Chul Kim, Dr. Shinyoung Kang, Dr. Dong-Hwa Seo, and Prof. Byoungwoo Kang) for their
support. Thanks to them, my life in the CEDER group could be more lively and vivid. I always
felt joys when spending time with them. Moreover, Nancy Twu and Wenhao Sun are greatly
acknowledged. I had to go through many different moments during the graduate life, and I was
so happy to share happy and harsh moments with them as the same-year CEDER group
members.
4
Samsung Scholarship is greatly acknowledged for supporting me both financially and
mentally. I will not forget happy moments that I shared with the Samsung Scholarship
community. I also thank the Robert Bosch Corporation and the Umicore Specialty Oxides and
Chemicals for their financial support on my research.
All the Korean graduated students in the DMSE are greatly acknowledged. They have
been a social family to me, guiding me toward a right direction for my graduate life. In
particular, I thank all the Korean students who joined DMSE MIT in the same year of 2010
(Donghun Kim, Mansoo Park, Jiyoun Chang, Gyehyun Kim, and Sangtae Kim). I spent more
time with them than any other people in MIT, and shared all the fun and happy moments with
them during the graduate life. I hope to keep in touch with them and wish their future stays as
bright as now.
I cannot thank enough to my parents and my brother. I could not have grown up strong
and confident without their endless support and love. Finally, I thank my beloved wife, Lina
Dahye Song, whose patience, support and love made this possible.
5
6
Contents
List of Figures and Tables ................................................................................................ 11
Part I. Introduction .......................................................................................................20
1.1 Motivation.....................................................................................................................21
1.2 Overview of rechargeable lithium batteries .........................................................22
1.3 Layered lithium transition metal oxides for Li-ion battery cathodes ........... 26
1.4 Ordering paradigm for the oxide cathodes ...........................................................30
1.5 Overview of the thesis ...............................................................................................35
Part II. Unlocking the potential of cation-disordered oxides for
rechargeable lithium batteries ................................................................................36
2.1 Introduction ..................................................................................................................37
2.2 Methodology ................................................................................................................39
2.2.1 Experimental methodology ........................................................................................ 39
2.2.1.1 Synthesis .............................................................................................................39
2.2.1.2 Electrochemistry..................................................................................................45
2.2.1.3 Characterization ..................................................................................................45
2.2.2 Computational methodology...................................................................................... 46
2.2.2.1 Density-functional theory calculations ............................................................... 46
2.2.2.2 Structural model ..................................................................................................47
2.2.2.3 Percolation simulation ........................................................................................ 48
7
2.3 Experimental results ...................................................................................................49
2.3.1 Characterization of the as-synthesized Li1.211Mo0.467Cr0.3O2 .....................................49
2.3.2 The electrochemical properties of Li1.211Mo0.467Cr0.3O2 ............................................51
2.3.3 The structural change in Li1.211Mo0.467Cr0.3O2 during cycling ...................................55
2.3.4 Redox mechanism of Li1.211Mo0.467Cr0.3O2 ................................................................ 59
2.4 Computational results ................................................................................................61
2.4.1 Li diffusion channels in cation-disordered oxides ..................................................... 61
2.4.2 Li diffusion barriers in cation-disordered oxides: active 0-TM channels ..................62
2.4.3 Percolation of 0-TM diffusion channels ....................................................................65
2.5 Discussions ...................................................................................................................67
2.6 Conclusion ....................................................................................................................69
Part III. A new class of high capacity cation-disordered oxides for
rechargeable lithium batteries: Li-Ni-Ti-Mo oxides .................................71
3.1 Introduction ..................................................................................................................72
3.2 Methodology ................................................................................................................74
3.2.1 Experimental methodology ........................................................................................ 74
3.2.1.1 Synthesis .............................................................................................................74
3.2.1.2 Electrochemistry .................................................................................................74
3.2.1.3 X-ray diffraction (XRD), scanning electron microscopy (SEM), and
electron energy loss spectroscopy (EELS) .......................................................... 75
3.2.1.4 In situ X-ray diffraction ...................................................................................... 76
3.2.1.5 Ex situ X-ray near edge spectroscopy (XANES) ................................................76
8
3.2.2 Computational methodology...................................................................................... 77
3.3 Experimental results ..................................................................................................79
3.3.1 Design strategy form percolation theory....................................................................79
3.3.2 Characterization of Li-Ni-Ti-Mo oxides ....................................................................81
3.3.3 The electrochemical properties of Li-Ni-Ti-Mo oxides .............................................86
3.3.4 The structural evolution of Li1.2Ni1/3Ti1/3Mo2/15O2 during cycling ............................ 92
3.3.5 Investigation on the redox mechanism ......................................................................93
3.4 Computational results ................................................................................................99
3.4.1 Oxygen loss mechanisms ........................................................................................... 99
3.4.2 Calculated voltage profiles from possible redox mechanisms .................................101
3.4.3 Reduction of Mo and Ti after oxygen loss ............................................................... 103
3.5 Discussions .................................................................................................................105
3.5.1 Redox mechanism ....................................................................................................105
3.5.2 Electrochemical performance .................................................................................. 107
3.5.3 Strategies for improvements .................................................................................... 109
3.6 Conclusion .................................................................................................................. 111
Part IV. Conclusions .................................................................................................... 112
References .......................................................................................................................... 117
9
10
List of Figures and Tables
Figure 1-1 Gravimetric power and energy densities for different rechargeable batteries, which
are currently being investigated for grid storage applications.
Figure 1-2 Schematic description of Li-ion batteries: reversible shuttling of Li+ ions between
the cathode and the anode enables Li-ion batteries.
Figure 1-3 Illustrations of the crystal structure of (a) a layered Li-TM oxide, (b) a spinel Li-TM
oxide, and an olivine LiMPO4 (M = Fe, Mn etc.)
Figure 1-4 The first-cycle voltage profile of bare and coated LiCoO2 when cycled between
2.75−4.4 V at a rate of 70 mA/g.
Figure 1-5 The first-cycle voltage profile of Li/LiNi1/3Mn1/3Co1/3O2 with a current density of
0.2 mA/cm2.
Figure 1-6 The first-cycle voltage profiles of 0.3Li2MnO3·0.7LiMn0.5Ni0.5O2.
Figure 1-7 The discharge profile at various C rates for the ion-exchanged (IE)-LiNi0.5Mn0.5O2
(above) and sold state (SS)-LiNi0.5Mn0.5O2 (below). In the test, the cell was charged at C/20 to
4.6 V, then held at 4.6 V for 5 hours and discharged at different rates. 1C corresponds to
280 mA/g.
Figure 1-8 (a) The octahedral to tetrahedral to octahedral Li diffusion in general rocksalt-type
Li-TM oxides (b) The 1-TM channel in stoichiometric Li-TM oxides
Figure 1-9 (a) The decrease in the Li slab distance (tetrahedron height) by cation mixing: the
layered structure has much larger Li slab spacing than the cation-disordered structure. (b) The
calculated activation barrier for Li migration in LiNiO2 with cation mixing (stars) and without
cation mixing (squares). The Li slab distances with and without cation mixing are indicated
11
with thick and thin dashes at different Li contents, respectively. The shift on the left is from
fully lithiated state and the one on the right is from partially delithiated state.
Figure 1-10 (a) The X-ray powder diffraction patterns of LiCoO2 powders prepared at high
temperature after subsequent exposure to different milling times. (b) Voltage profiles of
cathodes prepared from LiCoO2 samples exposed to different milling times.
Figure 1-11 (a) Charge-discharge curves of high-temperature α-LiFeO2 and low-temperature ɑLiFeO2 (b) Capacity retention of low-temperature α-LiFeO2.
Figure 2-1 Illustration of the layered rocksalt structure and disordered rocksalt structure: cation
mixing between Li and TM layers lead to a disordered rocksalt structure, which tends to result
in poor cyclability of Li-TM oxides by slowing down Li diffusion.
Figure 2-2 The voltage profile of Li2VO3.
Figure 2-3 High resolution transmission electrons microscopy images in two directions of the
carbon-coated Li1.211Mo0.467Cr0.3O2 particle (top), and uncoated Li1.211Mo0.467Cr0.3O2 particle
(bottom)
Figure 2-4 The two upper left images show a Li metal foil and a separator from a
Li1.211Mo0.467Cr0.3O2 half-cell that was charged/discharged for 100 cycles between 1.5−4.3 V at
164 mA/g at room temperature. The upper right SEM image shows spherical particles on the
colored area on the separator (dotted circle). The lower image shows an energy dispersive Xray spectroscopy (EDS) spectrum of a particle attached to the separator, indicating that the
particles from the colored area are composed of Mo and Cr (Mo:Cr = 8.875:1).
Figure 2-5 Li1.211Mo0.467Cr0.3O2 half-cells were disassembled at points A, B, and C on the first
charge voltage profile of Li1.211Mo0.467Cr0.3O2 at 3.2 mA/g. The three inset images with red
boundaries show Li metal anode foils (attached to a current collector rod) from
Li1.211Mo0.467Cr0.3O2 half cells that were disassembled after four days of resting at the first
charge state of A) 200 mAh/g, B) 250 mAh/g, and C) 300 mAh/g. The inset image with a blue
boundary shows a Li metal anode foil that was disassembled from a carbon-coated
12
Li1.211Mo0.467Cr0.3O2 half-cell after four days of resting at the same first charge state of
300 mAh/g as point C.
Figure 2-6 The discharge-capacity retention of carbon-coated and uncoated
Li1.211Mo0.467Cr0.3O2 during cycling between 1.5−4.3 V at 16.4 mA/g at room temperature. The
two inset images show Li metal anode foils from uncoated and carbon-coated
Li1.211Mo0.467Cr0.3O2 half-cells after ten cycles. Carbon coating retards transition metal
dissolution, improving the capacity retention of carbon-coated Li1.211Mo0.467Cr0.3O2.
Figure 2-7 The first-cycle voltage profile of the carbon-coated and uncoated
Li1.211Mo0.467Cr0.3O2 when cycled between 1.5−4.3 V at 32.7 mA/g at room temperature.
Figure 2-8 Structural model of disordered-Li2MoO3: The simulation cell contains 12 formula
units (72 atoms). In the figure, large red balls represent oxygen atoms, and small green and
violet balls represent the octahedral sites of lithium and molybdenum, respectively.
Figure 2-9 0-TM wrapping probability (a) and accessible lithium fraction (b) as a function of
the Li-excess level for different FCC supercell sizes. Results for fully disordered structures are
shown. The legends refer to supercells of the primitive FCC unit cell.
Figure 2-10 Rietveld refinement on the XRD pattern of pristine Li1.211Mo0.467Cr0.3O2 powder.
The atomic ratio between elements was first set to the ICP result, and then Mo and Cr
occupancy in Li and TM layers were refined. The refinement suggests no transition metals in
Li layers at all. Li occupancy was refined after Mo and Cr occupancy had been determined.
Figure 2-11 (a) The (003) peak in the XRD pattern on Li1.211Mo0.467Cr0.3O2 [black] and that of
LiCrO2 [blue] (b) The electron diffraction pattern along ZA[010] of Li1.211Mo0.467Cr0.3O2.
Figure 2-12 The broader XRD peaks of the carbon-coated Li1.211Mo0.467Cr0.3O2 powder (lower
left) compared to the sharper XRD peaks of pristine Li1.211Mo0.467Cr0.3O2 powder (upper left)
indicate that the carbon-coated Li1.211Mo0.467Cr0.3O2 particles are smaller than the uncoated
particles. Furthermore, SEM images (right) confirm that the particles of the carbon-coated
Li1.211Mo0.467Cr0.3O2 are smaller than those of uncoated one. During the carbon coating process,
13
sucrose (carbon precursor) and the Li1.211Mo0.467Cr0.3O2 powder were mixed in a planetary ballmill for six hours at 300 rpm, resulting in smaller particles after carbon coating.
Figure 2-13 The voltage profile of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled
between 1.5−4.3 V at the rate of 16.4 mA/g at room temperature
Figure 2-14 The 10-cycle voltage profiles of the carbon-coated Li1.211Mo0.467Cr0.3O2 when
cycled between 1.5−4.3 V at the rate of (a) 4.1 mA/g, (b) 8.2 mA/g, and (c) 16.4 mA/g at room
temperature (d) Corresponding specific capacity vs. cycle number plot
Figure 2-15 The voltage profile of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled
between 1.5−4.3 V at a rate of C/10, C/2, 1C, 2C, 4C and 10C (1C = 327 mA/g)
Figure 2-16 Capacity retention of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled between
1.5-4.3 V at a rate of C/10, C/2, 1C, 2C, 4C and 10C (1C = 327 mA/g)
Figure 2-17 The ex situ XRD pattern of the carbon-coated Li1.211Mo0.467Cr0.3O2 electrode
before cycle, after one cycle, after two cycles, and after 10 cycles between 1.5−4.3V at
32.7 mA/g.
Figure 2-18 Rietveld refinement on the XRD pattern of a carbon-coated Li1.211Mo0.467Cr0.3O2
electrode after ten cycles between 1.5−4.3 V at 32.7 mA/g. Kapton film was applied to prevent
oxidation, resulting in humps below 30°. From XRD refinement, it is difficult to precisely
quantify and distinguish the level of Mo migration and Cr migration to Li layers. Therefore, we
refined the XRD pattern with two different methods: (a) Refine Mo occupancy first then Cr
occupancy. Here, refinement suggests that only Mo migrates (0.2634 Mo in Li layers). (b)
Assume all Cr are in Li layers then refine Mo occupancy. Refinement is much improved by
assuming 0.097988 Mo in addition to 0.3 Cr in Li layers. In both methods, Li occupancy was
later refined by distributing 0.92 Li (based on the cumulative charge-discharge capacity after
ten cycles) to Li and TM layers. However, the values are more subject to error. The two ways
of refining give two extremes of the transition metal (TM) migration level. Refinement based
14
on the method A suggests a TM migration level of ~34.3% (i.e. ~34.3% of overall TM ions are
in Li layers), and refinement based on the method B suggests a TM migration level of ~51.88%.
From the two values, we can expect that at least ~34% to ~52% of TM ions have migrated to
Li layers after ten cycles, indicating substantial cation mixing in Li1.211Mo0.467Cr0.3O2. Note that
the TM migration level can slightly vary with the rate of cycling which changes the delithiation
level at the end of charge. To determine the average tetrahedron height of the disorderedLi1.211Mo0.467Cr0.3O2, eight tetrahedron heights (four from the tetrahedron in Li layers, and four
from that in TM layers) were averaged out, as Li1.211Mo0.467Cr0.3O2 is not yet completely
disordered. Method A yields 2.39701 Å (× 3), 2.38362 Å (× 3), 2.41831 Å (× 1), and
2.37791 Å (× 1), resulting in the average tetrahedron height of 2.39226 Å . Method B yields
2.39816 Å (× 3), 2.38775 Å (× 3), 2.40875 Å (× 1), and 2.37756 Å (× 1), resulting in the
average tetrahedron height of 2.393 Å . Therefore, both methods suggest the average
tetrahedron height of disordered- Li1.211Mo0.467Cr0.3O2 to be ~2.39 Å .
Figure 2-19 Left: the scanning transmission electron microscopy images along the [010] zone
axis in a carbon-coated Li1.211Mo0.467Cr0.3O2 particle before cycling and after 1 and 10 cycles
between 1.5−4.3 V at 16.4 mA/g. Right: Corresponding line profiles of the Z-contrast
information with the measured spacing of Li-Mo-Cr layers.
Figure 2-20 (a) The cyclic voltammetry profile of the carbon-coated Li1.211Mo0.467Cr0.3O2, (b)
Cr EELS L3/L2 ratio in Li1.211Mo0.467Cr0.3O2 before cycling, after the first charge, and after the
first cycle. (c) Integrated spin density of Mo and Cr in Li1.211Mo0.467Cr0.3O2 as a function of the
delithiation level, obtained by DFT calculations.
Figure 2-21 (a) o-t-o diffusion: Two tetrahedral paths connect each pair of neighboring
octahedral sites (b-d) The activated Li+ ion can face-share with no octahedral transition metals
(0-TM channel) (b), one transition metal (1-TM channel) (c), or two transition metals (2-TM
channel) (d): note that the 2-TM channels do not support divacancy diffusion mechanism, and
thus diffusion barrier through the 2-TM channels are extremely high.
Figure 2-22 Calculated Li migration barriers along 1-TM (Mo4+) channels (red squares), 1-TM
(Cr3+) channels (blue triangles), and 0-TM (Li+) channels (black circles) as a function of the
15
average tetrahedron height of model disordered structure (disordered Li2MoO3, disordered
LiCrO2). Error bars denotes standard deviation. The shaded area highlights typical tetrahedron
heights of disordered materials
Figure 2-23 (a) Computed probability to find a percolating network of 0-TM channels (color)
vs. Li content (x in LixTM2-xO2) and cation mixing (TMLi layers/TMTM layers ×100%) (b) The
accessible Li contents by a percolating 0-TM network (color) vs. Li content and cation mixing.
In the simulation, cations were randomly distributed at each cation mixing level
Table 2-1 The average tetrahedron height of disordered rocksalt-type Li-TM-oxides, deduced
from the literatures. The tetrahedron height of a disordered Li-TM-oxide is equivalent to its
(111) plane distance, which can be derived by dividing the a-lattice parameter of each material
by √3.
Figure 2-24 (a) The XRD patterns of the carbon-coated Li1.211Mo0.467Cr0.3O2 after ten cycles,
after ten cycles then charged to Li0.6165Mo0.467Cr0.3O2, and after ten cycles then charged to
Li0.3082Mo0.467Cr0.3O2 when cycled between 1.5−4.3 V at 32.7 mA/g. (b) The c- and a-lattice
parameter in disordered Li1.211Mo0.467Cr0.3O2 (based on the space group of R-3m) at different
delithiation states (x = 0.291, 0.5945, 0.90275 in Li1.211-xMo0.467Cr0.3O2). The change in the
lattice parameters is very small, leading to negligible volume change (<~0.12 %) upon
delithiation
Figure 3-1 (a) The crystal structure of disordered rocksalt-type solid solution compounds
between LiNi0.5Ti0.5O2 and Li1+xM1-xO2 (M = Ti4+, Nb5+, Mo6+) (b-d) Theoretical capacities are
given when M is (b) Ti4+ (c) Nb5+ and (d) Mo6+. Each figure plots three different capacities: Li
capacity that assumes full extraction of available Li-ions, Ni2+/Ni4+ redox capacity, and 0-TM
capacity that is the Li capacity accessible by the percolating 0-TM network
Figure 3-2 The X-ray diffraction patterns of LiNi0.5Ti0.5O2 (LNTO), Li1.05Ni11/24Ti11/24Mo1/30O2
(LNTMO5), Li1.1Ni5/12Ti5/12Mo1/15O2 (LNTMO10), Li1.15Ni3/8Ti3/8Mo1/10O2 (LNTMO15), and
Li1.2Ni1/3Ti1/3Mo2/15O2 (LNTMO20): insets are the a-lattice parameters of the compounds
Table 3-2 Target vs. actual Li:Ni:Ti:Mo atomic ratio as determined by direct current plasma
16
emission spectroscopy.
Figure 3-3 Rietveld refinements on the XRD patterns of (a) LiNi0.5Ti0.5O2 [LNTO], (b)
Li1.05Ni11/24Ti11/24Mo1/30O2 [LNTMO5] (c) Li1.1Ni5/12Ti5/12Mo1/15O2 [LNTMO10], (d)
Li1.15Ni3/8Ti3/8Mo1/10O2 [LNTMO15], and (d) Li1.2Ni1/3Ti1/3Mo2/15O2 [LNTMO20]: structural
parameters from the refinements are listed in Table 3-2.
Table 3-2 Structural parameters from the Rietveld refinements in Figure 3-3: crystallographic
information file of Fm-3m LiFeO2 was first used as an input file. The atomic occupancies were
initially set to the atomic ratio obtained from the elemental analysis by direct current plasma
emission spectroscopy, based on which the lattice parameters were initially refined. Then, we
further refined the lattice parameters and the atomic occupancies simultaneously: TM
occupancies were first refined, and then Li occupancy was refined. O occupancy did not
change after the refinement.
Figure 3-4 Scanning electron microscope (SEM) images of LNTO, LNTMO5, LNTMO10,
LNTMO15, LNTMO20, and high-energy ball-milled LNTMO20 (HB-LNTMO20)
Figure 3-5 (a) The first-cycle voltage profiles of LNTO, LNTMO5, LNTMO10, LNTMO15
and LNTMO20 [1.5−4.5 V, 20 mA/g], and (b) 20-cycle capacity retention of the compounds.
Figure 3-6 The voltage profiles for 10 cycles of (a) LNTO, and (b) LNTMO20 [1.5−4.5 V,
20 mA/g, room temperature].
Figure 3-7 The voltage profiles of (a) LNTO and (b) LNTMO20 when charged and discharged
once at 10 mA/g, and then at 20, 40, 100, 200, and 400 mA/g for the subsequent cycles.
Figure 3-8 (a) The first discharge voltage profile of LNTMO20 from a galvanostatic
intermittent titration test after charging to 270 mAh/g: the inset zooms in the time range
between 240 h and 270 h. (b) The voltage profile of LNTMO20 when charged-discharged three
times between 2.0−4.1 V (black) then between 1.5−4.5 V (blue) at 20 mA/g.
17
Figure 3-9 (a) The In situ XRD patterns of LNTMO20 upon two galvanostatic chargedischarge cycles between 1.5−4.8 V at 10 mA/g (b) The corresponding voltage profile (c) and
the a-lattice parameter from single phase XRD refinements are shown. (d) The (002) peak is
zoomed in for the intensity comparison.
Figure 3-10 The X-ray absorption near-edge structures of (a) Ni, (b) Ti, and (c) Mo in
LNTMO20 before cycle [black], after the first charge to 4.8 V [blue, ~300 mAh/g charged],
and after the first discharge to 1.5 V [red, ~250 mAh/g discharged] at 20 mA/g.
Figure 3-11 Electron energy loss spectra of Ti L-edge and O K-edge in LNTMO20 before
cycling [black] and after 20 cycles between 1.5−4.5 V [red] at 20 mA/g: the inset focuses on
the Ti L-edges.
Figure 3-12 The first-cycle cyclic voltammetry profiles of LNTMO20 when voltage-swept
between 1.5−4.5 V [black solid] and 1.5−4.1 V [red dash] at 0.1 mV/s.
Figure 3-13 Back-of-the-envelope calculation: Average Mo oxidation state expected after the
first charge to 300 mAh/g then discharge to 250 mAh/g as a function of the oxygen-loss
capacity during the first charge. It was assumed that Ni, Ti, and O stay as Ni2+, Ti4+, and O2after the first discharge, and that there is no loss in the TM content during the first cycle.
Figure 3-14 The calculated energies of densified Li0.52Ni0.37Ti0.37Mo0.15O2 structures and
Li0.47Ni0.33Ti0.33Mo0.13O1.8 structures with oxygen vacancies: From the plot, it is seen that the
energies of the densified Li0.52Ni0.37Ti0.37Mo0.15O2 structures are lower than those of
Li0.47Ni0.33Ti0.33Mo0.13O1.8 structures with oxygen vacancies. This indicates that oxygen loss
with densification is thermodynamically more favorable than that with oxygen vacancies in the
lattice.
Figure 3-15 (a) The voltage profiles of Li1.2-xNi0.33Ti0.33Mo0.13O2: The black curve is an
experimental profile from the galvanostatic intermittent titration test (GITT) during the first
charge, and the red line is calculated voltage profile of Li1.2-xNi0.33Ti0.33Mo0.13O2 that assumes
18
no oxygen loss during the first charge: the dotted red arrows specify the region of Ni oxidation
and O oxidation. Finally, the dashed blue line indicates the oxygen-loss potential of Li1.2xNi0.33Ti0.33Mo0.13O2-y.
(b) The calculated contribution of Ni [black] and oxygen [red] oxidation
upon delithiation in Li1.2-xNi0.33Ti0.33Mo0.13O2 when no oxygen loss is assumed (the red curve in
Figure 3-15a)
Figure 3-16 (a) The voltage profiles of Li1.11-xNi0.37Ti0.37Mo0.15O2. (b) The average net moments
of Ni, Ti, Mo and oxygen ions in Li1.11-xNi0.37Ti0.37Mo0.15O2 (x = 0, 0.074, 0.148, 0.222, 0.296,
0.370, 0.444, 0.519, 0.593, 0.667, 0.741, 0.815, 0.889, 0.963 and 1.037).
Figure 3-17 Proposed first charge mechanism of LNTMO20: Ni2+/Ni~3+ oxiation, oxygen loss,
and oxygen oxidation largely in sequence. The voltage profile of LNTMO20 from the in situ
XRD test and the corresponding a-lattice parameters from single-phase XRD refinements are
overlayed.
Figure 3-18 Illustrations of a LNTMO20 particle before and after oxygen loss with
densification near the surface: oxygen loss with densification reduces the Li-excess level, and
thus decreases the 0-TM capacity.
19
-------------Part I-------------
Introduction
20
1.1 Motivation
Sustainable growth of our modern society requires clean energy production, storage, and
transport. For the energy storage, rechargeable lithium batteries (i.e. Li-ion batteries) have
played an irreplaceable role as they outperform the other types of technologies in terms of
energy densities and power densities (Fig. 1-1).1,2 Due to their great performance, the Li-ion
batteries have enabled many of the convenience of modern life, powering increasingly capable
electronics. However, as the batteries are requested for more complex technologies where
higher power consumptions are required, such as for electric vehicles and grid energy storage,
the demand for higher performance Li-ion batteries has been ever increasing.2–6
Figure 1-1 Gravimetric power and energy densities for different rechargeable batteries, which
are currently being investigated for grid storage applications.2
The development of better Li-ion batteries requires improvements in the electrode
materials.1,5 In particular, current cathode materials do not satisfy all desired characteristics for
21
high-performance Li-ion batteries, including energy density, rate capability, safety, cycle life,
cost, and toxicity.5 Hence, the development of novel cathode materials is critical for the design
of high-performance Li-ion batteries for future applications such as in grid storage and longrange electric vehicles, thereby our society can grow sustainably and continuously.2,3,6
1.2 Overview of rechargeable lithium batteries
Figure 1-2 Schematic description of Li-ion batteries: reversible shuttling of Li+ ions between
the cathode and the anode enables Li-ion batteries.
The Li-ion batteries are electrochemical energy storage devices that operate by shuttling Li+
ions and electrons between the cathode and the anode both of which serve as a reservoir for Li
(Fig. 1-2).5,7–9 They are composed of three main components: the cathode, the anode, and the
electrolyte. The cathode and anode are the positive and negative electrodes, respectively, where
oxidation and reduction occur during charge and discharge.5,7 The electrolyte is an ionic
conductor but electronic insulator, allowing Li+ ions to migrate between the cathode and the
22
anode. Upon charge, an external power supply is applied to convert electrical energy into
chemical energy, transferring Li+ ions (through the electrolyte) and electrons (through an
external circuit) from the cathode to the anode, for example, with the following reactions:4,5,9
(1) Cathode:
(2) Anode:
LixMOy (cathode)  σ Li+ (electrolyte) + σ e- + Lix-σMOy (cathode)
σ Li+ (electrolyte) + σ e- + σ C6 (anode)  σ LiC6 (anode)
During charge, oxidation and reduction occurs at the cathode and anode, respectively. Upon
discharge, the reverse reactions occur, thus Li+ ions and electrons migrate from the anode to the
cathode. Concurrently, reduction and oxidation occurs at the cathode and anode, respectively,
releasing chemical energy into electrical energy.
In commercial Li-ion batteries, graphite is typically used as the anode, which allows
intercalation of Li+ ions between its sheets to form LiC6.4 The cathode usually is composed of a
cathode material from inorganic compounds with transition metal ions as redox active species,
carbon black to enhance electronic conductivity, and a polymeric binder such as poly
tetrafluorethylene (PTFE). The electrolyte typically consists of Li salts, such as LiPF6,
dissolved in organic solvents, such as ethylene carbonate (EC) with dimethyl carbonate
(DMC).7,10,11
One of the greatest aims in the Li-ion battery research is to develop very high energy
density Li-ion batteries. The energy density of a battery is defined as the energy per unit weight
or volume, which is equal to the product of the cell voltage and specific capacity.5,12 Voltage of
a Li-ion battery whose cathode and anode are both intercalation compounds for Li is
determined by the difference in the chemical potential of Li in the cathode and the anode:5,12
V=−
µLi(𝑐𝑎𝑡ℎ𝑜𝑑𝑒) − µLi(𝑎𝑛𝑜𝑑𝑒)
𝐹
In most commercial Li-ion batteries, LiCoO2 and graphite are used as the cathode and the
anode, respectively, which gives an average voltage of ~3.6 V.1,8
23
The specific capacity of Li-ion batteries is determined by the cyclable Li content per
unit weight or volume.5 As both cathode and anode are intercalation compounds for Li ions, the
specific capacity of a battery depends on those of both electrodes. The specific capacity of an
intercalation electrode can be calculated using the following equation:
Specific capacity =
𝛥𝑥·𝐹
𝑀
×
1000
3600
[mAh/g]
where Δx is the number of moles of Li that can reversibly inserted/extracted in and out of the
electrode structure within a designated voltage window. M is the molar mass of the electrode
compound. Finally, F is the Faraday constant. As can be inferred from the above formula, a
high specific capacity of an electrode requires a light electrode material (small M) that can
reversibly store and release a large amount of Li+ ions (high Δx).
High power density (or high rate capability) is also a highly desired property of Li-ion
batteries, particularly for high-power applications such as for electric vehicles or power
tools.13,14 The power density of a battery is defined as the battery power per unit weight, which
is the product of the current and the cell voltage.13 The power density is determined by the cell
resistances from various factors such as mass transfer resistance of electrodes (related to Li
diffusion), charge-transfer resistance at the electrode/electrolyte interface (related to solvation
and desolvation of Li salts), and electric resistances (related to ionic conduction) in the
electrolyte.8,13,14 A battery with large resistances will have a poor rate capability, and a large
portion of the chemically stored energy will dissipate into heat upon discharge.
Cycle life, cost, and safety are the other desired properties of Li-ion battery.8 Cycle life
is defined as the number of charge-discharge cycles that can be performed before the specific
capacity decreases below a certain cut-off limit. Using Li-ion batteries for a while, many kinds
of side reactions can occur, which can build up internal resistances on the batteries. Also, the
crystal structures of electrode materials can change during cycling, adding resistances as well.
24
Such increase in the resistances can lead to poor cycle life.8 Low cost is also very critical for
successful Li-ion batteries, because the cost determines the economic viability of a battery.8
The deployment of electric vehicles has been slow because of the high cost of Li-ion batteries,
resulting in no advantages of buying the electric vehicles instead of conventional gasoline
vehicles.15 The cost of a battery will decrease as processing cost and materials’ cost decrease.
The performance of the Li-ion batteries is largely determined by the properties of
electrode materials. Thus, the search for better cathode/anode materials has been intense.4,5,8
Electrode materials for high performance Li-ion batteries should have the following
characteristics. First, the materials should have a stable and sturdy crystal structure with many
available Li sites to reversibly store and release a large amount of Li+ ions. Second, the
material’s (molar) mass should be light such that the charge-storage capacity per unit weight
can be high. Third, the Li-intercalation potential of the electrode material should be within a
voltage window of an electrolyte for safety.5,8 Within an acceptable voltage window, high (low)
intercalation potential is preferred for a cathode (anode) material, because it will result in high
voltage thus high energy densities. Fourth, Li diffusion should be facile in the crystal structure,
allowing for fast Li (de)intercalation reaction with negligible resistances.13,14 Fast Li diffusion
is critical to achieve high energy densities and high rate capability from an electrode material
and thus from a battery. Finally, the electrode materials should be thermally safe (stable) such
that they do not release oxygen and heat at a high temperature. Oxygen evolution is an
exothermic reaction, and thus can substantially increase temperature inside a battery. With
oxygen gas, high temperature, and a flammable organic electrolyte, a Li-ion battery can catch
fire, which must be prevented for safety.16,17
25
1.3 Layered lithium transition metal oxides for
Li-ion battery cathodes
While all components are important in Li-ion batteries, the cathode has been the limiting
component on the current stage of Li-ion battery technology. This is because the cathode is the
major factor that determines the energy density, rate capability, safety, and cost of Li-ion
batteries.5,8 As the cathode materials, intercalation compounds have been most widely used due
to their peculiar ability to accommodate Li+ ions over large concentration intervals.
Figure 1-3 Illustrations of the crystal structure of (a) a layered Li-TM oxide, (b) a spinel Li-TM
oxide, and an olivine LiMPO4 (M = Fe, Mn etc.)
Today, three types of intercalation cathodes are being intensively studied, so called
layered (rocksalt-type) lithium transition metal oxides (Li-TM oxides),18–24 spinel-type Li-TM
oxides,25–30 and olivine LiMPO4 (M = Fe, Mn etc.) (Fig. 1-3).14,31–37 While each class has its
own strength, none of them are perfect. For example, Li-excess layered oxides offer the highest
capacity (> 220 mAh/g) and energy density (~900 Wh/kg), but are less safe than other types of
cathode materials.38–45 On the other hand, spinel-type oxides have excellent rate capability but
26
with low specific capacity (~100 mAh/g), low energy density, and poor cycle life at high
temperature.25,27,30,46–49 Finally, olivines such as LiFePO4 are cheap, safe and have excellent
cycle life and rate capability, but with low gravimetric and volumetric energy densities
(590 Wh/kg, 2000 Wh/l) due to their low density (~3.4 kg/l) and low average voltage (~3.5 V
vs. Li/Li+).14,16,32,33
Among the cathode materials, the layered Li-TM oxides have been the most successful
class, delivering high capacities, high volumetric, and gravimetric energy densities.40,41,50 The
layered Li-TM oxide consists of close-packed oxygen planes that are stacked with the ABC
stacking. Between the oxygen planes, layers of TM ions are alternating with layers of Li
ions.13,51 TM ions in the TM layers can release or take electrons by redox reactions. Li layers
are where Li ions are stored and Li diffusion takes place. Many commercial Li-ion batteries
adopt cathode materials from the layered oxides.
Figure 1-4 The first-cycle voltage profile of bare and coated LiCoO2 when cycled between
2.75−4.4 V at a rate of 70 mA/g.52
27
LiCoO2 and LiNiO2 were the first to be investigated as cathode materials.19,21,52–57 In
particular, LiCoO2 was the cathode material in one of the first commercially successful Li-ion
batteries, and has been used most widely until now. LiCoO2 delivers a high gravimetric
(~600 Wh/kg) and volumetric (~3000 Wh/l) energy density due to its high voltage (~3.9 V) and
high density (5 kg/l) (Fig. 1-4). Nevertheless, high cost of cobalt, a limited specific capacity
from LiCoO2 (~160 mAh/g) and poor thermal safety have encouraged researchers to search for
the other types of layered Li-TM oxides to replace LiCoO2.20,43
To develop layered cathode materials that outperform LiCoO2 both in performance and
cost, Li(Ni/Mn/Co/Al)O2 have been intensively studied, in which TM layers are occupied
mixed TM ions of Mn, Ni, Co and Al. In particular, LiNi1/3Mn1/3Co1/3O2, LiNi0.5Mn0.5O2, and
LiNi0.8Co0.15Al0.05O2 have been most intensively studied (Fig. 1-5).13,22,58 They deliver high
energy densities (~700 Wh/kg) and specific capacities (~180 mAh/g). Furthermore, they are
cheaper than LiCoO2 as well, because they contain less cobalt which is expensive. Thus, the
second generation Li-ion batteries use one of the materials as the cathode material.
Figure 1-5 The first-cycle voltage profile of LiNi1/3Mn1/3Co1/3O2 with a current density of
0.2 mA/cm2.22
28
While most of the traditional layered materials are stoichiometric materials, there have
been great efforts to achieve very high capacities and energy densities from Li-excess layered
materials.40,45,59,60 Here, Li-excess refers to a composition in which the number of Li sites is
greater than that of TM sites (x > 1 in LixTM2-xO2). In particular, many high capacity materials
have been designed out of solid-solution compounds (or nano-composite compounds) between
Li(Li1/3Mn2/3)O2 and Li(Ni/Mn/Co)O2.40,41,59–61 These materials can cycle more than
250 mAh/g at an average voltage of ~3.6 V, delivering a high energy density (~ 900 Wh/kg)
(Fig. 1-6). Nevertheless, their crystal structures evolve during cycling, leading to voltage
fading.41,62–64 Such evolution of the voltage profile is undesirable for Li-ion batteries, because
the state of charge cannot be determined if a voltage profile evolves during cycling. Therefore,
substantial efforts have been made to stabilize the crystal structure of the Li-excess materials.62
Moreover, the rate capability (i.e. power density) of the materials is moderate, which needs
improvement.65
Figure 1-6 The first-cycle voltage profiles of 0.3Li2MnO3·0.7LiMn0.5Ni0.5O2.61
29
1.4 Ordering paradigm for the oxide cathodes
As we have discussed, nearly all high energy density cathode materials have been sought from
well-ordered close-packed oxides, such as layered Li-TM oxides or spinel-type Li-TM oxides.
In these ordered materials, Li sites and pathways (a 2D slab in the layered oxides and a 3D
network of tetrahedral sites in the spinels) are separated from the TM sublattice, which
provides stability and electron storage capacity.13,30
Having well-ordered structures where there is little or no intermixing between the Li
and the TM sublattice is generally considered critical to obtain high-capacity cathode materials
with good cycle life.7,8 In some cases, improvements in cation ordering have notably improved
the power or energy density.13,66–68 Figure 1-7 shows an example of the effect of cation
(dis)ordering to the performance of the ordered oxides.13 In general, LiNi0.5Mn0.5O2 forms into
a slightly cation-mixed layered structure in which ~ 9 % of Li sites are occupied by TM ions
(accordingly ~9 % TM sites are occupied by Li ions) when standard high-temperature solid
state methods are applied for synthesis. In 2006, Kang et al. demonstrated that removing the
cation mixing by ion-exchange dramatically improves the performance of the layered
LiNi0.5Mn0.5O2.13 The ion-exchanged layered LiNi0.5Mn0.5O2 has almost a perfectly layered
structure with negligible cation mixing, and shows much higher capacity and better rate
capability than a standard LiNi0.5Mn0.5O2 with a partially cation-mixed layered structure.13
Such improvement largely originates from that Li diffusion tends to be faster in a cationordered structure than in a (partially or fully) cation-disordered structure.
In general rocksalt-type oxides, both Li and TM ions occupy a cubic close-packed
lattice of octahedral sites, and Li diffusion proceeds by hopping from one octahedral site to
30
Figure 1-7 The discharge profile at various C rates for the ion-exchanged (IE)-LiNi0.5Mn0.5O2
(above) and sold state (SS)-LiNi0.5Mn0.5O2 (below). In the test, the cell was charged at C/20 to
4.6 V, then held at 4.6 V for 5 hours and discharged at different rates. 1C corresponds to
280 mA/g.13
another octahedral site via an intermediate tetrahedral site.51,69 Li in the intermediate
tetrahedral site is the activated state in Li diffusion. The activated tetrahedral Li ion shares
faces with four octahedral sites: the site previously occupied by the ion itself; the vacancy it
will move into; and two sites that can be occupied by Li, TM, or a vacancy. The energy in this
site, which reflects the Li migration barrier, is largely determined by electrostatic repulsion
between the activated Li+ ion and its face-sharing species, and thus depends on (i) the valence
of the face-sharing species and (ii) the available space for relaxation between the activated Li+
ion and the face-sharing species.13,51,68 This space is measured as the Li slab distance in layered
structures or more generally as the height of the tetrahedron along which the relaxation
occurs.13,51,70
31
Figure 1-8 (a) The octahedral to tetrahedral to octahedral Li diffusion in general rocksalt-type
Li-TM oxides (b) The 1-TM channel in stoichiometric Li-TM oxides.70
As electrostatic repulsion on an activated Li+ ion is too strong when there are two facesharing cations, Li dominantly diffuses with the divacancy mechanism, involving a second
vacancy beside the vacancy the migrating Li will move into.13,51,69,70 In rocksalt-type Li-TM
oxides, two kinds of diffusion channels support this mechanism: 0-TM channels, involving no
face-sharing TM ion, and 1-TM channels, involving one face-sharing TM ion.70–72 Note that in
the traditional stoichiometric layered Li-TM oxides such as LiCoO2, Li diffusion takes place
through the 1-TM channels only.
In general, cation disorder (i.e. cation mixing) limits Li diffusion in the layered
(rocksalt-type) oxides, inhibiting Li cycling from the layered materials.13,66,68,73 The negative
effect of cation mixing on Li diffusion has been understood to originate from the following
reasons. First, cation mixing greatly decreases the relaxation space (i.e. Li slab distance or
tetrahedron height) for an activated Li+ ion against a face-sharing species upon Li migration,
increasing the electrostatic energy in the activated state and thus the Li diffusion barrier (Fig.
1-9).13,70 Moreover, some people have argued that TM ions in Li layers can physically block
the Li diffusion paths as well.74
32
Figure 1-9 (a) The decrease in the Li slab distance (tetrahedron height) by cation mixing: the
layered structure has much larger Li slab spacing than the cation-disordered structure.70 (b) The
calculated activation barrier for Li migration in LiNiO2 with cation mixing (stars) and without
cation mixing (squares).51 The Li slab distances with and without cation mixing are indicated
with thick and thin dashes at different Li contents, respectively. The shift on the left is from
fully lithiated state and the one on the right is from partially delithiated state.51
The negative effects of cation disorder have been experimentally demonstrated as well.
For example in 1998, Obravac et al. demonstrated that mechanochemical synthesis of cationdisordered LiMO2 (M = Ti, Mn, Co, Ni) is possible with ball milling (Fig. 1-10).75 In their
experiments, increasing the ball-milling time resulted in more severe cation mixing in layered
materials such as LiCoO2 and LiNiO2. Then, they demonstrated that the reversible capacity of
layered materials rapidly decreases as their crystal structure changes from a layered rocksalt to
disordered rocksalt structure.75 Furthermore, most disordered rocksalt-type Li-TM oxides such
as α-LiFeO2 are electrochemically inactive, delivering a limited capacity (Fig. 1-11).75–77 Thus,
only a negligible attention has been given to cation-disordered oxides as cathode materials,
based on the experimental observations and the understanding of the negative effects of cation
disorder to Li diffusion.
33
Figure 1-10 (a) The X-ray powder diffraction patterns of LiCoO2 powders prepared at high
temperature after subsequent exposure to different milling times. (b) Voltage profiles of
cathodes prepared from LiCoO2 samples exposed to different milling times.75
Figure 1-11 (a) Charge-discharge curves of high-temperature α-LiFeO2 and low-temperature ɑLiFeO2 (b) Capacity retention of low-temperature α-LiFeO2.76
34
1.5 Overview of the thesis
This thesis explores the potential of cation-disordered Li-TM oxides as promising electrode
materials for rechargeable Li batteries, enlarging the search space of high-energy density
cathode materials to cation-disordered materials.
In the first part, we demonstrate that contrary to the conventional wisdom, cationdisordered materials can be promising electrode materials from the study of novel
Li1.211Mo0.467Cr0.3O2.70 Li1.211Mo0.467Cr0.3O2 forms as a layered rocksalt structure but transforms
almost completely to a disordered rocksalt structure after several cycles. Nevertheless, the
reversible capacity of this material is remarkably high, indicating that Li diffusion can be facile
in the cation-disordered structure.70 Using ab inito computations, we show that this unexpected
behavior is due to percolation of a certain type of active Li diffusion channels (0-TM channels)
that are available in disordered Li-excess materials.70,72
In the second part, we apply 0-TM percolation to design a new class of high capacity
cation-disordered oxides, Li-excess Ni-Ti-Mo oxides, which deliver both high capacities and
high voltage. Combining experiments with ab initio calculations, we investigate their
electrochemical performance, structural changes, and redox mechanism. From the
understandings, strategies to further improve the new materials are presented, providing new
guidelines for the design of high-energy density cation-disordered cathode materials for the
rechargeable Li batteries.
35
-------------Part II-------------
Unlocking the potential of
cation-disordered oxides for
rechargeable lithium batteries
[Reproduced from Ref. 70] Lee, J. et al., Science 343, 519-522 (2014).
36
2.1 Introduction
Most high-energy density cathode materials for Li-ion batteries have been sought from the
well-ordered close-packed oxides, such as layered LiCoO2 and spinel LiMn2O4.8,26,53 In these
ordered materials, Li and TM ions occupy distinct sites such that Li sites and pathways are
separated from the TM sublattice which provides stability and electron storage capacity.
Having well-ordered structures with negligible mixing between the Li and the TM sublattice is
generally considered necessary for obtaining high-capacity cathode materials with good cycle
life.13,66 This is largely because Li diffusion tends to be facile in the well-ordered crystal
structure, which is critical for achieving high energy and high power densities. As discussed in
Part I: Introduction, cation mixing tends to limit Li diffusion because it decreases the Li slab
spacing, increasing the Li diffusion barrier.51,68,70 Moreover, randomly distributed TM ions in
the cation-disordered structure can physically block Li diffusion paths as well. Therefore,
cation-disordered materials are generally disregarded as electrode materials, and have received
only a limited attention (Fig. 2-1).
Figure 2-1 Illustration of the layered rocksalt structure and disordered rocksalt structure: cation
mixing between Li and TM layers lead to a disordered rocksalt structure, which tends to result
in poor cyclability of Li-TM oxides by slowing down Li diffusion.
37
However, in fact there are a few cation-disordered Li-TM oxides with good cycling
performance. For example, cation-disordered Li2VO3 can cycle up to 250 mAh/g (Fig. 2-2),78
and recently it has been demonstrated that cation-disordered Li1.3Mn0.4Nb0.3O2 can cycle up to
300 mAh/g as well.79 These results show that Li diffusion can be facile in the cation-disordered
structure. By understanding the mechanism how these cation-disordered materials allow for
facile Li diffusion, we can design a whole new class of high-capacity cation-disordered
materials for rechargeable lithium battery electrodes.
Figure 2-2 The voltage profile of Li2VO3.78
In this part of the thesis, from the study of Li1.211Mo0.467Cr0.3O2 we demonstrate that
cation-disordered materials can be promising electrode materials if they contain enough Liexcess (x > 1.09 in LixTM2-xO2). Li1.211Mo0.467Cr0.3O2 forms into a layered structure, but
transforms to a cation-disordered structure after several charge-discharge cycles. Nevertheless,
its reversible capacity is remarkably high, delivering as high as ~265 mAh/g. In this work, we
not only demonstrate the high capacity in this material, but also provide a fundamental
understanding of how cation-disordered materials can be promising electrode materials, using
38
ab initio calculations. We believe this work unlocks the potential of cation-disordered oxides
for rechargeable lithium batteries.
2.2. Methodology
2.2.1 Experimental methodology
2.2.1.1 Synthesis
To synthesize Li1.211Mo0.467Cr0.3O2 [original target: Li1.233Mo0.467Cr0.3O2], Li2CO3 (Alfa Aesar,
ACS, 99% min), MoO2 (Alfa Aesar, 99%), and Cr3(OH)2(OOCCH3)7 (Alfa Aesar, Cr 24%)
were used as precursors. 5% excess Li2CO3 from the stoichiometric amount required to
synthesize Li1.233Mo0.467Cr0.3O2 was used to compensate for possible Li loss during high
temperature solid state reaction. The precursors were dispersed into acetone and ball-milled for
24 hours, and then dried overnight in an oven. The precursor mixture was pelletized and then
fired at 1050°C for 15 hours under argon (Ar) gas, followed by furnace cooling to room
temperature. After firing, the pellets were manually ground into fine powder. From inductively
coupled plasma atomic emission spectroscopy (ICP-AES, bomb digestion method), slight Li
deficiency (~1.8%) from the target composition was found in the final product
(Li1.211Mo0.467Cr0.3O2). Carbon coating was applied to the Li1.211Mo0.467Cr0.3O2 powder (i) to
prevent transition metal (TM) dissolution (mainly Mo) at high delithiation states (Fig. 2-8), and
(ii) to improve the electrochemical properties of Li1.211Mo0.467Cr0.3O2 by decreasing its particle
size, through planetary ball-milling in the carbon coating process (Fig. 2-7). Sucrose
39
(C12H22O11, EMD, GR ACS), the carbon precursor, was mixed with the Li1.211Mo0.467Cr0.3O2
powder in the weight ratio of 80 (Li1.211Mo0.467Cr0.3O2): 20 (Sucrose) by planetary ball-milling
(Retsch, PM200) for six hours at 300 rpm in Ar-filled jars. Finally, the mixture was annealed at
700°C for three hours under Ar gas to obtain the carbon-coated Li1.211Mo0.467Cr0.3O2 powder.
The CHN analysis (combustion method) detected 5.17 wt% of carbon from the carbon-coated
Li1.211Mo0.467Cr0.3O2. Figure 2-3 shows the Li1.211Mo0.467Cr0.3O2 particle with and without
carbon coating. As can be seen from the figures, thin layers of amorphous carbon are coated to
the Li1.211Mo0.467Cr0.3O2 particle after carbon coating, which are not seen from the pristine
Li1.211Mo0.467Cr0.3O2 particle.
Figure 2-3 High resolution transmission electrons microscopy images in two directions of the
carbon-coated Li1.211Mo0.467Cr0.3O2 particle (top), and uncoated Li1.211Mo0.467Cr0.3O2 particle
(bottom).70
40
The carbon-coating layers prevent TM dissolution (mainly Mo) from
Li1.211Mo0.467Cr0.3O2 upon cycle. It is known that Mo6+ dissolves into an electrolyte.80 As Mo4+
becomes oxidized to Mo6+ upon charging Li1.211Mo0.467Cr0.3O2, it is necessary to prevent
surface from Mo dissolution for stable cycling performance. Figure 2-4 demonstrates how Mo
dissolution can be significant if without coating. High energy ball-milled Li1.211Mo0.467Cr0.3O2
was cycled for 100 cycles between 1.5−4.3 V at 164 mA/g at room temperature. After 100
cycles, significant coloration is seen from the Li metal anode and the separators, which are
confirmed to be from dissolved Mo compounds based on the energy dispersive X-ray
spectroscopy.
Figure 2-4 The two upper left images show a Li metal foil and a separator from a
Li1.211Mo0.467Cr0.3O2 half-cell that was charged/discharged for 100 cycles between 1.5−4.3 V at
164 mA/g at room temperature. The upper right SEM image shows spherical particles on the
colored area on the separator (dotted circle). The lower image shows an energy dispersive Xray spectroscopy (EDS) spectrum of a particle attached to the separator, indicating that the
particles from the colored area are composed of Mo and Cr (Mo:Cr = 8.875:1).70
41
However, such Mo dissolution can be retarded with carbon coating.
Figure 2-5 shows the Li metal anode of half-cells from uncoated and carboncoated Li1.211Mo0.467Cr0.3O2. Without carbon coating, charging to a higher
capacity beyond 250 mAh/g leads to significant Mo dissolution as can be inferred
from the coloration on the Li metal anode. With carbon coating, the coloration
becomes much less, indicating that Mo dissolution is retarded.
Figure 2-5 Li1.211Mo0.467Cr0.3O2 half-cells were disassembled at points A, B, and C on the first
charge voltage profile of Li1.211Mo0.467Cr0.3O2 at 3.2 mA/g. The three inset images with red
boundaries show Li metal anode foils (attached to a current collector rod) from
Li1.211Mo0.467Cr0.3O2 half cells that were disassembled after four days of resting at the first
charge state of A) 200 mAh/g, B) 250 mAh/g, and C) 300 mAh/g. The inset image with a blue
boundary shows a Li metal anode foil that was disassembled from a carbon-coated
Li1.211Mo0.467Cr0.3O2 half-cell after four days of resting at the same first charge state of
300 mAh/g as point C.70
42
Figure 2-6 The discharge-capacity retention of carbon-coated and uncoated
Li1.211Mo0.467Cr0.3O2 during cycling between 1.5−4.3 V at 16.4 mA/g at room temperature. The
two inset images show Li metal anode foils from uncoated and carbon-coated
Li1.211Mo0.467Cr0.3O2 half-cells after ten cycles. Carbon coating retards transition metal
dissolution, improving the capacity retention of carbon-coated Li1.211Mo0.467Cr0.3O2.70
Preventing Mo dissolution can lead to significant improvement in the
cycle retention of Li1.211Mo0.467Cr0.3O2. Figure 2-6 compares the cycle retention of
Li1.211Mo0.467Cr0.3O2 with and without carbon coating. As can be seen from the
plot, the carbon-coated Li1.211Mo0.467Cr0.3O2 delivers a much higher capacity with
better capacity retention than without carbon coating. The carbon coating process
is further beneficial because during the process the particle size of
Li1.211Mo0.467Cr0.3O2 becomes smaller, improving the reversible capacity by
increasing the surface to volume ratio. Figure 2-7 compares the first cycle voltage
43
profile of the carbon-coated and uncoated Li1.211Mo0.467Cr0.3O2 when cycled
between 1.5−4.3 V at 32.7 mA/g at room temperature. As can be seen from the
profiles, the reversible capacity is higher for the carbon-coated Li1.211Mo0.467Cr0.3O2.
This is due to the smaller particle size of the carbon-coated Li1.211Mo0.467Cr0.3O2
than the uncoated Li1.211Mo0.467Cr0.3O2 (Fig. 2-12)
Figure 2-7 The first-cycle voltage profile of the carbon-coated and uncoated
Li1.211Mo0.467Cr0.3O2 when cycled between 1.5−4.3 V at 32.7 mA/g at room temperature.
44
2.2.1.2 Electrochemistry
The cathode film was composed of the carbon-coated Li1.211Mo0.467Cr0.3O2 powder (70 wt%),
carbon black (20 wt%) as conductive agent, and polytetrafluoroethylene (PTFE) (10 wt%) as
binder. The components were manually mixed for 40 minutes and rolled into a thin film inside
an argon-filled glove box. 1M of LiPF6 in 1:1 volume ratio of ethylene carbonate: dimethyl
carbonate (EC:DMC) solution was used as an electrolyte. Celgard 2500 polypropylene
separator and Li metal foil were used as the separator and the counter electrode, respectively.
The Swagelok cells were assembled inside an argon-filled glove box and tested on a Maccor
2200 at room temperature in galvanostatic mode unless otherwise specified. The loading
density of the cathode film was approximately 3.6 mg/cm2. The current density at 1C
(= 327.486 mAg-1) was based on the theoretical capacity of Li1.233Mo0.467Cr0.3O2
(= 327.486 mAhg-1). The specific capacity was calculated strictly on the amount of
Li1.211Mo0.467Cr0.3O2 (66.318 wt%) in the cathode film.
2.2.1.3 Characterization
The X-ray diffraction (XRD) patterns were collected on a PANalytical multipurpose
diffractometer using Cu Kα radiation in the two-theta range of 5-85°. Kapton film was applied
when running XRD on electrode films. Rietveld refinement was done using PANalytical
X’pert HighScorePlus software. Scanning electron microscopy (SEM) images were collected
on a FEI Phillips XL30 field-emission gun environmental SEM at an accelerating voltage of 15
and 25 kV, respectively. High Resolution Transmission Electron Microscopy (HRTEM)
images were recorded with JEOL 2010F at MIT. High Angle Annular Dark Field (HAADF)
45
Scanning Transmission Electron Microscopy (STEM) images were recorded with Hitachi HD
2700C 200 kV using 1 Å probe of 28mrad semi-convergence angle and ADF detector with
semi-collection angle of 50-28mrad.
2.2.2 Computational methodology
2.2.2.1 Density-functional theory calculations
Structural energies and migration barriers were calculated based on density-functional theory
(DFT)81,82 in the generalized-gradient approximation (GGA) using the PBE exchangecorrelation functional83 and projector-augmented wave (PAW) pseudopotentials,84,85 as
implemented in the Vienna ab-initio Package (VASP).86,87 A plane-wave cutoff of 520 eV was
employed to guarantee numerical convergence in variable-cell calculations. For chromium, a
Hubbard-U correction was employed, using the U-value of reference 89 (U= 3.5 eV).88 Gamma
centered k-point meshes for the integration of the Brillouin zone were selected to converge
energies to 1 meV per formula unit and structural parameters to 10-3 Å . All geometries were
fully optimized with atomic forces below 0.01 eV/Å . The climbing-image nudged elastic band
(CI-NEB) method was used to compute migration barriers based on five images along the
diffusion paths.89–91 All reported migration barriers were converged to 10 meV. The
calculations were conducted for the well-established divacancy diffusion mechanism.51,69 All
NEB calculations resulted in diffusion paths that traverse tetrahedral sites, in agreement with
the mechanism found in layered oxides.
46
2.2.2.2 Structural model
Structural energies and migration barriers were obtained from structural models containing
72 atoms in the fully lithiated state, i.e., 12 Li2MoO3 formula units and 18 LiCrO2 formula
units, respectively (Fig. 2-8). Cation disorder was imposed by randomly distributing Li and
TM ions over all cation sites. For the migration barrier calculations, a divacancy was created in
the fully lithiated structure.
Figure 2-8 Structural model of disordered-Li2MoO3: The simulation cell contains 12 formula
units (72 atoms). In the figure, large red balls represent oxygen atoms, and small green and
violet balls represent the octahedral sites of lithium and molybdenum, respectively.70
47
2.2.2.3 Percolation simulation
Percolation simulations were conducted following the Monte-Carlo algorithm proposed by
Newman and Ziff and employed periodic 8×8×8 supercells of the conventional (4 atom) FCC
cell containing 2048 sites and 1000 MC steps per each combination of cation mixing and Li
content.92 Percolation was approximated by periodic wrapping in one lattice dimension. Since
the wrapping probability is sensitive with respect to the size of the simulation cell, we verified
the convergence of the percolation threshold for the limit of fully disordered structures,
corresponding to 100% cation mixing in Figure 2-9a. As evident from Figure 2-9A, the
inflection point of the wrapping probability rapidly converges with the cell size and therefore is
a good estimate for the percolation probability even in small simulation cells. Finite size effects
are less pronounced for the accessible lithium contents above the percolation threshold, as
shown in Figure 2-9b.
Figure 2-9 0-TM wrapping probability (a) and accessible lithium fraction (b) as a function of
the Li-excess level for different FCC supercell sizes. Results for fully disordered structures are
shown. The legends refer to supercells of the primitive FCC unit cell.70
48
2.3 Experimental results
2.3.1 Characterization of the as-synthesized Li1.211Mo0.467Cr0.3O2
As-synthesized Li1.211Mo0.467Cr0.3O2 has a layered rocksalt structure, which can be inferred
from its XRD pattern in Figure 2-10. In this material, layers of Li and of Li-Mo-Cr are
alternating in the oxygen cubic close-packed framework. The XRD refinement on the assynthesized Li1.211Mo0.467Cr0.3O2 shows no cation mixing between the Li and the Li-Mo-Cr
layers, indicating that this material forms into a well-layered structure.
Figure 2-10 Rietveld refinement on the XRD pattern of pristine Li1.211Mo0.467Cr0.3O2 powder.
The atomic ratio between elements was first set to the ICP result, and then Mo and Cr
occupancy in Li and TM layers were refined. The refinement suggests no transition metals in
Li layers at all. Li occupancy was refined after Mo and Cr occupancy had been determined.70
While Li1.211Mo0.467Cr0.3O2 is targeted to be a solid-solution compound between
Li(Li1/3Mo2/3)O2 (i.e. Li2MoO3) and LiCrO2, one can suspect local phase-separation of
Li1.211Mo0.467Cr0.3O2 into domains of Li(Li1/3Mo2/3)O2 and LiCrO2. However, neither in XRD
or TEM is there any indication of the local domains. The c-lattice parameter of Li2MoO3
49
(~14.91 Å ) and LiCrO2 (~14.43 Å ) are quite different from each other. Thus, if there were such
distinct domains, the XRD signature near 18.4o from LiCrO2 (in blue in Figure 2-11a) would
be visible in the diffraction pattern of Li1.211Mo0.467Cr0.3O2. However, such feature is not
observed in the XRD pattern of Li1.211Mo0.467Cr0.3O2 (black). In addition, the electron
diffraction pattern along ZA [010] of pristine Li1.211Mo0.467Cr0.3O2 in Figure 2-11b shows a
single phase, which further excludes the existence of small LiCrO2 domains (< 5 nm) beyond
the detectability of XRD. Hence, all the evidence points towards a single solid solution.
Figure 2-11 (a) The (003) peak in the XRD pattern on Li1.211Mo0.467Cr0.3O2 [black] and that of
LiCrO2 [blue] (b) The electron diffraction pattern along ZA[010] of Li1.211Mo0.467Cr0.3O2.
The scanning electron microscopy (SEM) images show that the primary particles
around 200 nm in diameter are highly agglomerate in the as-synthesized Li1.211Mo0.467Cr0.3O2,
and the primary particle size decreases to around 100 nm after the carbon-coating process (Fig.
2-12). The decrease in the particle size can also be inferred from the change in the XRD
patterns in which peak broadening is observed after carbon coating.
50
Figure 2-12 The broader XRD peaks of the carbon-coated Li1.211Mo0.467Cr0.3O2 powder (lower
left) compared to the sharper XRD peaks of pristine Li1.211Mo0.467Cr0.3O2 powder (upper left)
indicate that the carbon-coated Li1.211Mo0.467Cr0.3O2 particles are smaller than the uncoated
particles. Furthermore, SEM images (right) confirm that the particles of the carbon-coated
Li1.211Mo0.467Cr0.3O2 are smaller than those of uncoated one. During the carbon coating process,
sucrose (carbon precursor) and the Li1.211Mo0.467Cr0.3O2 powder were mixed in a planetary ballmill for six hours at 300 rpm, resulting in smaller particles after carbon coating.70
2.3.2 The electrochemical properties of Li1.211Mo0.467Cr0.3O2
Figure 2-13 shows the voltage profile of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled
between 1.5−4.3 V at the rate of 16.4 mA/g at room temperature. From the profile, we find a
very high capacity of ~265 mAh/g, corresponding to cycling ~1 Li per formula unit of
Li1.211Mo0.467Cr0.3O2. Considering that conventional well-layered materials show a capacity of
~160 to ~190 mAh/g, we can see that the capacity from Li1.211Mo0.467Cr0.3O2 is remarkably
51
high. The gravimetric energy density of this material is also promising (~660 Wh/kg), which
exceeds those of the conventional cathode materials such as LiCoO2 (~590 Wh/kg), LiMn2O4
(~480 Wh/kg), and LiFePO4 (~595 Wh/kg).14,26,57 From the voltage profile, we find that the
first-charge profile is different from those of further cycles. This indicates that the crystal
structure of Li1.211Mo0.467Cr0.3O2 changes during the first charge.
Figure 2-13 The voltage profile of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled
between 1.5−4.3 V at the rate of 16.4 mA/g at room temperature.
52
Figure 2-14 The 10-cycle voltage profiles of the carbon-coated Li1.211Mo0.467Cr0.3O2 when
cycled between 1.5−4.3 V at the rate of (a) 4.1 mA/g, (b) 8.2 mA/g, and (c) 16.4 mA/g at room
temperature (d) Corresponding specific capacity vs. cycle number plot.70
The rate capability of the carbon-coated Li1.211Mo0.467Cr0.3O2 is also promising. Figure
2-14a, 14b, and 14c shows the 10-cycle voltage profiles of the carbon-coated
Li1.211Mo0.467Cr0.3O2 when cycled between 1.5−4.3 V at the rate of 4.1 mA/g, 8.2 mA/g and
16.2 mA/g, respectively. From the profiles, we find a gentle decrease in the reversible capacity
with increasing rates. The carbon-coated Li1.211Mo0.467Cr0.3O2 still shows a high capacity above
200 mAh/g at 4.1 mA/g, indicating that the rate capability of this material is promising. The
cycle retention of Li1.211Mo0.467Cr0.3O2 is also promising. During 10 cycles, the average loss of
the discharge capacity per cycle is ~2 mAh/g, which is reasonably good for a new material
without any significant engineering (Fig. 2-14d). As the rate further increases, the reversible
53
capacity further decreases (Figs. 2-15 and 2-16). Nevertheless, the carbon-coated
Li1.211Mo0.467Cr0.3O2 can still deliver as much as 135 mAh/g at a very high rate of 1300 mA/g
(4C rate).
Figure 2-15 The voltage profile of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled
between 1.5−4.3 V at a rate of C/10, C/2, 1C, 2C, 4C and 10C (1C = 327 mA/g)
Figure 2-16 Capacity retention of the carbon-coated Li1.211Mo0.467Cr0.3O2 when cycled between
1.5−4.3 V at a rate of C/10, C/2, 1C, 2C, 4C and 10C (1C = 327 mA/g)
54
2.3.3 The structural change in Li1.211Mo0.467Cr0.3O2 during cycling
The electrochemical performance of the Li1.211Mo0.467Cr0.3O2 is promising, but its voltage
profile upon first charge is different from those upon later cycles (Fig. 2-13). Voltage is
determined by the difference in the chemical potential of Li in a cathode and an anode.12 In our
experiment, we used Li metal as the anode. Thus, the chemical potential of Li in the anode
does not change. Thus, the change in the voltage profile must come from the change in the
chemical potential of Li in the cathode during the first cycle. As the chemical potential of Li is
largely determined by the crystal structure of the cathode compounds, the change in the voltage
profile indicates that the crystal structure of the material changes substantially upon first charge.
To analyze the change in the crystal structure during cycling, we performed ex situ
XRD on the carbon-coated Li1.211Mo0.467Cr0.3O2 electrode before cycling, after one cycle, after
two cycles, and after ten cycles between 1.5−4.3V at 32.7 mA/g (Fig. 2-17). From the XRD
patterns, dramatic changes in the pattern are seen after the first cycle. Before cycling, the XRD
pattern shows a pattern of a layered-rocksalt structure (space group: R-3m) in which the Li
layers and the TM layers form alternating layers within an oxygen FCC framework (Fig. 2-1a).
However, after the first cycle, the (003) peak which is an indicative of the layered structure is
almost completely gone,73 and the XRD pattern resembles a pattern of a disordered-rocksalt
structure (space group: Fm-3m). Based on the XRD refinements (Fig. 2-18), we estimate 34 to
52 % of the TM ions to be in Li layers after ten cycles. This indicates substantial cation-mixing
between the Li and TM layers, known to be detrimental to the performance of layered materials.
The structural evolution of Li1.211Mo0.467Cr0.3O2 to a disordered structure can also be confirmed
in real space with scanning transmission electron microscopy (STEM) (Fig. 2-19). The bright
and dark columns in the “before” image correspond to atomic columns of mixed Li-Mo-Cr
55
ions and Li ions, respectively. The Z-contrast decreases after one cycle and is very weak after
10 cycles, indicating increased cation mixing.
As both XRD and STEM show, Li1.211Mo0.467Cr0.3O2 transforms to a disordered
rocksalt-structure after several cycles. Nevertheless, the carbon-coated Li1.211Mo0.467Cr0.3O2
still delivers a high capacity up to 265 mAh/g after substantial cation mixing. It has been
understood that cation-disordered structures do not allow for facile Li diffusion, largely
because the Li slab spacing (i.e. tetrahedron height) becomes very small in the structures.13,66,68
Indeed, the Li slab spacing of Li1.211Mo0.467Cr0.3O2 decreases considerably from ~2.63 Å to
~2.39 Å after cation mixing. Thus, the high capacity after cation mixing is very
counterintuitive and exciting, because the mechanism by which this material delivers a very
high capacity will allow us to design promising cation-disordered materials as high-capacity
lithium battery electrodes.
Figure 2-17 The ex situ XRD pattern of the carbon-coated Li1.211Mo0.467Cr0.3O2 electrode
before cycle, after one cycle, after two cycles, and after 10 cycles between 1.5−4.3V at
32.7 mA/g.
56
Figure 2-18 Rietveld refinement on the XRD pattern of a carbon-coated Li1.211Mo0.467Cr0.3O2
electrode after ten cycles between 1.5−4.3 V at 32.7 mA/g. Kapton film was applied to prevent
oxidation, resulting in humps below 30°. From XRD refinement, it is difficult to precisely
quantify and distinguish the level of Mo migration and Cr migration to Li layers. Therefore, we
refined the XRD pattern with two different methods: (a) Refine Mo occupancy first then Cr
occupancy. Here, refinement suggests that only Mo migrates (0.2634 Mo in Li layers). (b)
Assume all Cr are in Li layers then refine Mo occupancy. Refinement is much improved by
assuming 0.097988 Mo in addition to 0.3 Cr in Li layers. In both methods, Li occupancy was
later refined by distributing 0.92 Li (based on the cumulative charge-discharge capacity after
ten cycles) to Li and TM layers. However, the values are more subject to error. The two ways
of refining give two extremes of the transition metal (TM) migration level. Refinement based
on the method A suggests a TM migration level of ~34.3% (i.e. ~34.3% of overall TM ions are
in Li layers), and refinement based on the method B suggests a TM migration level of ~51.88%.
From the two values, we can expect that at least ~34% to ~52% of TM ions have migrated to
Li layers after ten cycles, indicating substantial cation mixing in Li1.211Mo0.467Cr0.3O2. Note that
the TM migration level can slightly vary with the rate of cycling which changes the delithiation
level at the end of charge. To determine the average tetrahedron height of the disordered57
Li1.211Mo0.467Cr0.3O2, eight tetrahedron heights (four from the tetrahedron in Li layers, and four
from that in TM layers) were averaged out, as Li1.211Mo0.467Cr0.3O2 is not yet completely
disordered. Method A yields 2.39701 Å (× 3), 2.38362 Å (× 3), 2.41831 Å (× 1), and
2.37791 Å (× 1), resulting in the average tetrahedron height of 2.39226 Å . Method B yields
2.39816 Å (× 3), 2.38775 Å (× 3), 2.40875 Å (× 1), and 2.37756 Å (× 1), resulting in the
average tetrahedron height of 2.393 Å . Therefore, both methods suggest the average
tetrahedron height of disordered- Li1.211Mo0.467Cr0.3O2 to be ~2.39 Å .70
Figure 2-19 Left: the scanning transmission electron microscopy images along the [010] zone
axis in a carbon-coated Li1.211Mo0.467Cr0.3O2 particle before cycling and after 1 and 10 cycles
between 1.5−4.3 V at 16.4 mA/g. Right: Corresponding line profiles of the Z-contrast
information with the measured spacing of Li-Mo-Cr layers.70
58
2.3.4 Redox mechanism of Li1.211Mo0.467Cr0.3O2
(a)
(b)
(c)
Figure 2-20 (a) The cyclic voltammetry profile of the carbon-coated Li1.211Mo0.467Cr0.3O2, (b)
Cr EELS L3/L2 ratio in Li1.211Mo0.467Cr0.3O2 before cycling, after the first charge, and after the
first cycle. (c) Integrated spin density of Mo and Cr in Li1.211Mo0.467Cr0.3O2 as a function of the
delithiation level, obtained by DFT calculations.
To understand the redox mechanism of Li1.211Mo0.467Cr0.3O2, we performed the cyclic
voltammetry (CV) test. From the CV test (Fig. 2-20a), we see a redox peak at ~2.4 V, and a
small peak at ~ 4 V after Li1.211Mo0.467Cr0.3O2 becomes cation-disordered. The L3/L2 intensity
ratio in electron energy loss spectroscopy (EELS) on the carbon-coated Li1.211Mo0.467Cr0.3O2
shows that Cr oxidation state increases mainly after Li1.211Mo0.467Cr0.3O2 is being charged to
above 4 V (~240 mAh/g, C/20), as shown in Figure 2-20b, which suggests that Mo 4+/Mo6+
redox (~248 mAh/g) is largely being used before Cr redox is being used (~80 mAh/g): the error
59
bar on each data point in Figure 2-20b gives the distribution of the results as obtained from ten
different particles on the sample region with less than 0.3 inelastic mean free path. The blue
and yellow bands in Figure 2-20b are the distributions of Cr EELS L3/L2 ratios for Cr3+ and
Cr4+, respectively, from different Cr oxides standards in the literature.93 These results indicate
that Cr is only being oxidized above 4.0 V and to some oxidation state between +3 and 4+.
Therefore, we determine the redox peak at ~2.4 V to be of Mo4+/Mo6+ and the peak at ~4 V to
be of Cr3+/Cr4+. Note that the ~2.4 V redox reaction is consistent with the Mo4+/Mo6+ redox
voltage in Li2MoO3.80
Additionally, we computationally assessed the redox mechanism of the cationdisordered Li1.211Mo0.467Cr0.3O2. Our calculations show that Mo is mostly being oxidized
before Cr oxidation, which is in agreement with our experimental observations. Only after
more than ~75 % of Mo4+ is oxidized to Mo6+, does the oxidation of Cr3+ start together with the
remainder of Mo oxidation. This can be seen from the change in the average integrated spin
densities of Mo and Cr in Figure 2-20c. While the spin on Mo decreases continuously upon Li
removal, the spin on Cr3+ only decreases slightly at the top of charge. Therefore, combining CV,
EELS and DFT calculations, we propose that Mo4+/Mo6+ redox is mainly being utilized, and
Cr3+/Cr4+ redox can be used as well after Mo4+ ions are fully oxidized to Mo6+.
60
2.4 Computational results
Li1.211Mo0.467Cr0.3O2 delivers a high capacity and energy density even after substantial cation
mixing, indicating that Li diffusion can be facile in the cation-disordered structure. This is in
stark contrast to the conventional wisdom that an ordered crystal structure is necessary for
facile Li diffusion. Then, how can we understand this counterintuitive behavior? To explain
this, we need to understand the Li diffusion mechanism in general rocksalt-type oxides.
2.4.1 Li diffusion channels in cation-disordered oxides
In general rocksalt-type oxides, both Li and TM ions occupy the octahedral sites, and Li
diffusion takes place by traversing an intermediate tetrahedral site (Fig. 2-21a).13,51,69–71 Li in
this tetrahedral site is the activated state in Li diffusion, whose electrostatic energy largely
determines the Li diffusion barrier. Therefore, the species in the face-sharing octahedral sites
and the tetrahedron height that serves as the relaxation space for the activated Li+ ion govern
the Li mobility through each diffusion channel.13,51,66,68
As electrostatic repulsion on an activated Li ion is too strong when there are two facesharing cations, Li dominantly diffuses with the divacancy mechanism, involving a second
vacancy beside the vacancy the migrating Li will migrate into.51,69,71 As discussed in the Part I:
Introduction, two types of diffusion channels support this mechanism in rocksalt-type oxides
thus in cation-disordered oxides: 0-TM channels, involving no face-sharing TM ion (Fig 221b), and 1-TM channels, involving one face-sharing TM ion (Fig. 2-21c).70–72 Note that the 1TM channels are responsible for Li diffusion in typical layered Li-TM oxides.
61
Figure 2-21 (a) o-t-o diffusion: Two tetrahedral paths connect each pair of neighboring
octahedral sites (b-d) The activated Li+ ion can face-share with no octahedral transition metals
(0-TM channel) (b), one transition metal (1-TM channel) (c), or two transition metals (2-TM
channel) (d): note that the 2-TM channels do not support divacancy diffusion mechanism, and
thus diffusion barrier through the 2-TM channels are extremely high.70
2.4.2 Li diffusion barriers in cation-disordered oxides: active 0-TM
channels
To investigate which channels in disordered-Li1.211Mo0.467Cr0.3O2 allow for reasonable hopping
rates, Li migration barriers for 1- and 0-TM channels were calculated using density-functional
theory, according to the divacancy mechanism.
62
Figure 2-22 Calculated Li migration barriers along 1-TM (Mo4+) channels (red squares), 1-TM
(Cr3+) channels (blue triangles), and 0-TM (Li+) channels (black circles) as a function of the
average tetrahedron height of model disordered structure (disordered Li2MoO3, disordered
LiCrO2). Error bars denotes standard deviation. The shaded area highlights typical tetrahedron
heights of disordered materials.70
The red and blue dashed lines in Figure 2-22 show the mean migration barriers along a
1-TM channel as a function of the average tetrahedron height of model disordered structures
(disordered-Li2MoO3, disordered-LiCrO2) when the face-sharing octahedral species is Mo4+
and Cr3+, respectively. Note that migration barriers in disordered structures vary with the local
atomic environment, which accounts for a distribution of migration barriers. The mean barrier
increases as the tetrahedron height (h) decreases, and reaches ~510 meV along a 1-Mo4+
channel and ~490 meV along a 1-Cr3+ channel at h ~ 2.39 Å , the average tetrahedron height in
disordered-Li1.211Mo0.467Cr0.3O2. Note that these barriers increases as the TM ion becomes
oxidized in charge.13,51 Considering that typical 1-TM barriers in layered oxides are
63
~300 meV,51 such high barriers in disordered-Li1.211Mo0.467Cr0.3O2 indicate limited Li diffusion
along 1-TM channels. This is because the small tetrahedron height in disorderedLi1.211Mo0.467Cr0.3O2 confines the activated Li+ ion close to a face-sharing high-valent
octahedral TM ion in 1-TM channels, resulting in strong electrostatic repulsion on the Li+ ion.
The black line in Figure 2-22 shows the mean migration barriers along 0-TM channels.
In contrast to the high barriers in 1-TM channels, the low ~290 meV barrier at h ~ 2.39 Å
indicates that Li migration along 0-TM channels will still be facile in disorderedLi1.211Mo0.467Cr0.3O2, with a ~4400 times [exp(500 meV/kBT)/exp(290 meV/kBT)] higher
hopping rate than along 1-TM channels at room temperature. The low valence of a face-sharing
octahedral Li+ ion (vs. Mo4+ or Cr3+) results in much weaker electrostatic repulsion on the
activated Li+ ion in 0-TM channels. At highly charged states, tetrahedral Li may form in some
0-TM channels since high delithiation should leave no face-sharing octahedral Li at all.
However, the mean migration barrier between two 0-TM tetrahedral sites was calculated to be
~415 meV, indicating that Li can easily escape from these sites.
.
Herein lies the real issue of cation-disordered structures: 1-TM channels, which
account for the great Li mobility in the layered Li-TM oxides which currently dominate the
battery industry, become nearly inactive in the cation-disordered materials due to their small
tetrahedron heights. In contrast, 0-TM channels are active in disordered rocksalts, but are much
less frequent than 1-TM channels. Nevertheless, as we will show below, 0-TM channels start to
enable facile macroscopic Li diffusion in disordered structures, once enough Li excess is
introduced.
64
2.4.3 Percolation of 0-TM diffusion channels
For 0-TM channels to allow for macroscopic Li diffusion, they must be continuously
connected through the entire material, forming a percolating network uninterrupted by 1- and
2-TM channels. To establish a general understanding of 0-TM percolation, we investigated 1)
when 0-TM channels percolate in a rocksalt-type Li-TM-oxide and 2) which fraction of the Li
ions become part of a percolating network of 0-TM channels.
Figure 2-23a shows the probability to find a percolating network of 0-TM channels (0TM network) in a rocksalt-type Li-TM-oxide as a function of Li content (x in LixTM2-xO2) and
cation mixing (TMLi layers/TMTM layers × 100 %), as obtained by Monte-Carlo simulations. The
probability (color-coded) steeply increases from zero (red) to one (blue) across the black line in
Fig. 2-23a (percolation threshold), varying from x ~ 1.13 for layered oxides to x ~ 1.09 for
fully disordered oxides. As 0-TM channels require a locally Li-rich environment, excess Li
(x ≥ ~1.09) is crucial to open the percolating 0-TM network.
To estimate the contribution of a percolating 0-TM network to macroscopic Li diffusion,
we investigated how Li excess and cation mixing affect the Li content in the network (Fig. 223b), which we refer to as accessible Li.70,72 This Li can diffuse through the network without
traversing 1- or 2-TM channels, while inaccessible Li must traverse 1- or 2-TM channels to
reach the percolating 0-TM network. The three black lines in Fig. 2-12b are the contour lines
where the accessible Li content is 0.8 Li, 1 Li, and 1.2 Li per LixTM2-xO2. For x ≤ 1, no
percolating 0-TM network exists (Fig. 2-23a), thus there is no accessible Li content, explaining
why stoichiometric Li-TM oxides have low capacity when cation disordered.73,75–77 However,
the accessible Li content gradually increases as x exceeds ~1.09 (percolation threshold), and
becomes as high as 1 Li as x exceeds ~1.22 regardless of cation mixing. Increasing Li excess
65
adds more 0-TM channels to a percolating 0-TM network, improving the network’s
connectivity.70,72
Figure 2-23 (a) Computed probability to find a percolating network of 0-TM channels (color)
vs. Li content (x in LixTM2-xO2) and cation mixing (TMLi layers/TMTM layers ×100%) (b) The
accessible Li contents by a percolating 0-TM network (color) vs. Li content and cation mixing.
In the simulation, cations were randomly distributed at each cation mixing level.70
66
The above results explain how Li diffusion can be facile in the cation-disordered
Li1.211Mo0.467Cr0.3O2. Li1.211Mo0.467Cr0.3O2 is a Li-excess material with x = 1.233 in LixTM2-xO2.
With this Li content, 0-TM channels will be percolating (Fig. 2-23a), accessing as high as
~1 Li per formula unit (Fig. 2-23b). Therefore, even as 1-TM channels become nearly inactive
after cation mixing, a large fraction of Li in the material can still be cycled through the
percolating active 0-TM network.
2.5 Discussions
The principle of creating a percolating 0-TM network can be applied to design other high
performing disordered Li-TM-oxides for two reasons. First, the 0-TM activated state is only
surrounded by Li sites, making the effect of the TM species on the activation energy less
pronounced. Secondly, as shown in Table 2-1, the tetrahedron height of most disordered
rocksalts is such that 0-TM channels are calculated to be active (Fig. 2-22). Therefore, a
percolating 0-TM network will likely enable facile Li diffusion in other disordered materials,
assuming no other kinetic barriers become limiting. Note that the few cation-disordered
materials in the literature that were electrochemically active are indeed Li excess materials,
whereas stoichiometric disordered materials are usually not electrochemically active, which is
consistent with our understanding.73,75–78,94
Disordered Li-excess rocksalts have considerable advantages over layered materials.
First, we find that the changes in lattice parameters and volume, as function of Li concentration,
are very small in disordered materials (< 1 % in Li1.211Mo0.467Cr0.3O2), which will lead to less
mechanical stresses and capacity loss in an electrode (Fig. 2-24). Furthermore, as their cation
distribution is more homogenous, they tend to have less significant changes of local
67
environment of the Li ions as a function of state of charge. This change in environment is
particularly problematic in layered structures where the Li slab spacing decreases significantly
when large amounts of Li are removed, leading to a substantial reduction of Li mobility.69,95,96
However, in the cation-disordered structure, homogenously distributed cations should lead to a
Li diffusivity that is more independent of the Li concentration, as is the case for electrode
materials with the spinel and olivine type structures. One issue that requires more investigation
is whether cation disordering will lead to a more sloped voltage profile than for well-ordered
materials, as one would expect from the wider distribution of Li-site energies in a cationdisordered material. However, this variance in the Li-site energy may be counteracted by a less
effective Li-Li interaction, which is responsible for the slope of the voltage curve in layered
materials.56 Therefore, careful tailoring of the TM-Li to Li-Li ion interaction may mitigate this
effect. Given the insights presented in this work, it may not be surprising that the highest
capacity layered materials are highly Li-excess materials45,61,97 which become more disordered
in the first few cycles due to a particular overcharge mechanism.60,98,99
Materials
75
LiTiO2
LiMnO275
α-LiFeO2100
Li0.542Co1.458O275
Li0.8388Ni1.1612O275
LiNi0.5Ti0.5O2101
LiCo0.5Ti0.5O2102
a (Å )
Tetrahedron height (Å )
Ref.
4.149(1)
4.179(4)
4.157(1)
4.143(2)
4.074(1)
4.1453(6)
4.14
2.395(4)
2.412(9)
2.400(1)
2.392(0)
2.352(1)
2.3933(2)
2.39
75
75
100
75
75
101
102
Table 2-1 The average tetrahedron height of disordered rocksalt-type Li-TM-oxides, deduced
from the literatures. The tetrahedron height of a disordered Li-TM-oxide is equivalent to its
(111) plane distance, which can be derived by dividing the a-lattice parameter of each material
by √3.70
68
Figure 2-24 (a) The XRD patterns of the carbon-coated Li1.211Mo0.467Cr0.3O2 after ten cycles,
after ten cycles then charged to Li0.6165Mo0.467Cr0.3O2, and after ten cycles then charged to
Li0.3082Mo0.467Cr0.3O2 when cycled between 1.5−4.3 V at 32.7 mA/g. (b) The c- and a-lattice
parameter in disordered Li1.211Mo0.467Cr0.3O2 (based on the space group of R-3m) at different
delithiation states (x=0.291, 0.5945, 0.90275 in Li1.211-xMo0.467Cr0.3O2). The change in the
lattice parameters is very small, leading to negligible volume change (< ~0.12 %) upon
delithiation.70
2.6 Conclusion
In conclusion, through the high capacity of carbon-coated Li1.211Mo0.467Cr0.3O2, which
transforms from a layered to a cation-disordered oxide during electrochemical cycling, we have
shown that Li diffusion can be facile in cation-disordered materials. Using DFT calculations
and percolation theory, we have explained this by a percolating network of transition states
with no transition metals around them (0-TM channels), which provides macroscopic diffusion.
The 0-TM percolation threshold is found to be x ~ 1.09 in LixTM1-xO2, but accessing 1 Li per
formula unit requires x ≥ ~1.22, in great agreement with our experimental results. Our results
may explain why disorder has not been pursued as a strategy before: Most materials
69
synthesized are near stoichiometry (LiTMO2), which is well below the percolation threshold
for 0-TM diffusion. Thus, these materials quickly lose their capacity upon disorder as it renders
typical 1-TM channels inactive, while 0-TM channels are not percolating.73,75–77 As a result,
cation disorder may have appeared as a counterintuitive strategy. In contrast, our analysis
points to cation-disordered materials as an exciting new class of materials that can deliver high
capacity and high energy density, and therefore offers one of the best hopes to substantially
improve the performance of rechargeable Li batteries.
70
-------------Part III-------------
A new class of high capacity
cation-disordered oxides for
rechargeable lithium batteries:
Li-Ni-Ti-Mo oxides
71
3.1 Introduction
From the study of Li1.211Mo0.467Cr0.3O2, important progress has been made in the oxide space,
which enlarges the search space of high energy density cathode materials to cation-disordered
lithium transition metal oxides (Li-TM oxides). While cation-disordered materials were widely
are disregarded as electrode materials,73,75–77 they can be promising cathode materials if
containing enough Li excess (x > 1.09 in LixTM2-xO2).70,72,79,103–105
To briefly summarize our understanding, the Li-excess content in a close-packed oxide
is important as it can strongly affect Li diffusion. In the close-packed oxides, Li diffusion
occurs between two connected octahedral sites through an intermediate tetrahedral site.13,51,70–72
A Li ion in this tetrahedral site is the activated state in Li diffusion, whose electrostatic energy
largely determines the Li diffusion barrier. 13,51,70–72 As the activated Li ion feels weaker
electrostatic repulsion when it avoids face-sharing high valent TM ions, the diffusion barrier
through a channel with no face-sharing TM ions around the activated state (0-TM channels) is
lower than through other types of channels.70–72 In the cation-disordered oxides, Li diffusion
can be facile only through these 0-TM channels. Nevertheless, for the channels to support
macroscopic Li diffusion, they must be percolating in the disordered structure, which requires a
Li-excess composition. As the Li-excess level increases, the percolating network of 0-TM
channels becomes more extensive, enabling a higher fraction of Li ions in a disordered
structure to cycle through the network.70,72
While this understanding is very exciting, the energy density of Li1.211Mo0.467Cr0.3O2
(from which the understanding is developed) is in fact not particularly high even with a very
high capacity of ~ 265 mAh/g.70 This originates from a low redox potential (~ 2.6 V) of
Mo4+/Mo6+ couple in the material. Therefore, in the second part of this thesis, we try to search
72
for high voltage cation-disordered materials to achieve higher energy densities. For the design
of such materials, we directly apply the percolation theory, demonstrating the validity of
percolation theory. In this work, we present a new class of high capacity and high voltage
cation-disordered oxides: lithium nickel titanium molybdenum oxides (Li-Ni-Ti-Mo oxides).
Combining electrochemistry with in situ X-ray diffraction, electron energy loss spectroscopy,
and X-ray absorption near edge spectroscopy, we investigate their electrochemical properties,
redox mechanism and structural changes upon cycling. From this understanding, strategies to
improve the new materials are presented, setting new guidelines for the design of high energy
density cation-disordered cathode materials for rechargeable lithium batteries.
73
3.2 Methodology
3.2.1 Experimental methodology
3.2.1.1 Synthesis
To synthesize Li1+x/100Ni1/2-x/120Ti1/2-x/120Mox/150O2 (x = 0, 5, 10, 15, 20), Li2CO3 (Alfa Aesar,
ACS, 99% min), NiCO3 (Alfa Aesar, 99 %), TiO2 (Alfa Aesar, 99.9 %), and MoO2 (Alfa Aesar,
99%) were used as precursors. Other than for LiNi0.5Ti0.5O2, a stoichiometric amount of
precursors were used. For LiNi0.5Ti0.5O2, 5 % excess Li precursor and 4 % excess Ni precursor
were used, because it resulted in the purest disordered rocksalt phase with a composition close
to the desired composition (Fig. 2, Table 1). The precursors were dispersed into acetone and
ball-milled for 15 hours, and then dried overnight in an oven. The mixture of the precursors
was pelletized and then sintered at 750°C for two hours in air, followed by furnace cooling to
room temperature. After the sintering, the pellets were manually ground into fine powder.
3.2.1.2 Electrochemistry
To prepare a cathode film, the powder of the Li-Ni-Ti-Mo oxides and carbon black (Timcal,
Super P) were first mixed by a planetary ball mill (Retsch PM200) in the weight ratio of 70:20
for two hours at 300 rpm. Then, polytetrafluoroethylene (PTFE, DuPont, Teflon 8C) was added
to the mixture as a binder, such that the cathode film consists of the Li-Ni-Ti-Mo oxide powder,
carbon black, and PTFE in the weight ratio of 70:20:10. The components were manually mixed
74
for 30 minutes and rolled into a thin film inside an argon-filled glove box. To assemble a cell
for all cycling tests, except for in situ X-ray diffraction, 1 M of LiPF6 in ethylene carbonate
(EC) and dimethyl carbonate (DMC) solution (1:1, Techno Semichem), Celgard 2500
polypropylene separator, and Li metal foil (FMC) were used as the electrolyte, the separator,
and the counter electrode, respectively. Swagelok-type cells were assembled inside an argonfilled glove box and tested on a Maccor 2200 at room temperature in the galvanostatic mode
otherwise specified. Cyclic voltammetry tests were performed on a Solartron electrochemical
potentiostat (1470E) between 1.5−4.1 V (or 1.5−4.5 V) at 0.1 mV/s. The loading density of the
cathode film was ~5 mg/cm2. The specific capacity was calculated on the amount of the Li-NiTi-Mo oxides (70 wt %) in the cathode film.
3.2.1.3 X-ray diffraction (XRD), scanning electron microscopy (SEM), and
electron energy loss spectroscopy (EELS)
The X-ray diffraction (XRD) patterns for the as-prepared compounds were collected on a
PANalytical multipurpose diffractometer (Cu source) in the 2ϴ range of 5−85°. Rietveld
refinement was completed using PANalytical X’pert HighScore Plus software. Scanning
electron microscopy (SEM) images were collected on a Zeiss Merlin High-resolution SEM.
Elemental analysis on the compounds was performed with direct current plasma emission
spectroscopy (ASTM E 1097-12). Electron energy loss spectroscopy (EELS) spectra were
obtained from thin specimens on a JEOL 2010F equipped with a Gatan spectrometer, using
parallel incident electron beam and semi-collection angle of 8 mrad in TEM diffraction mode.
EELS quantification was performed by using a signal integration window of 50 eV, HartreeSlater model of partial ionization cross section, and power law background subtraction.
75
3.2.1.4 In situ X-ray diffraction
For in situ XRD, an in situ cell was designed with a Be window for X-ray penetration. The cell
was configured with a Li1.2Ni1/3Ti1/3Mo2/15O2 electrode film as the working electrode, Li metal
foil as the counter electrode, 1M of LiPF6 in EC:DMC (1:1) solution as the electrolyte, and
glass fiber as the separator. Galvanostatic charge-discharge of the in situ cell was performed on
a Solartron electrochemical potentiostat (SI12287) between 1.5−4.8 V at 10 mA/g. The in situ
XRD patterns were obtained in one hour intervals from a Bruker D8 Advanced Da Vinci Mosource diffractometer (Mo source) in the 2Ө range of 7−36°. Rietveld refinement on the in situ
XRD patterns was performed using PANalytical X’pert HighScore Plus software for every
other scan.
3.2.1.5 Ex situ X-ray absorption near edge spectroscopy (XANES)
Ni, Ti and Mo K-edge XANES measurements were performed in transmission made using
beamline 20BM at the Advanced Photon Source. The incident energy was selected using a Si
(111) monochromator. The energy calibration was performed by simultaneously measuring the
spectra of the appropriate metal foil. Harmonic rejection was accomplished using a Rh-coated
mirror. The samples for the measurements were prepared with the Li1.2Ni1/3Ti1/3Mo2/15O2
electrode films (a) before cycling, (b) after the first charge to 4.8 V at 20 mA/g, and (c) after
the first charge to 4.8 V then discharge to 1.5 V at 20 mA/g. The loading density of the films
was ~5 mg/cm2. Additionally, spectra of some reference standards were measured in
transmission mode, to facilitate interpretation of the XANES data. Data reduction was carried
out using the Athena software.106
76
3.2.2 Computational methodology
First principles calculations were carried out with density functional theory (DFT) using the
spin polarized generalized gradient approximation (GGA).83 Hubbard U parameters (GGA + U)
were used to correct the self-interaction of GGA,107 using U values of 6.0 eV for Ni and U
values of 4.4 eV for Mo.108 The projector-augmented wave pseudopotentials were used for all
energy calculations as implemented to the Vienna Ab initio Simulation Package (VASP).86
To determine the cation-disordered Li1.2Ni0.33Ti0.33Mo0.13O2 structure, a large number of
Li/Ni/Ti/Mo cation orderings were generated by using the genetic algorithm method109,110
within a 5 × 3 × 2 supercell containing thirty formula units of monoclinic LiMO2 primitive cell
(space group: C2/m). To model the cation-disordered structure, the large supercell was selected
and the compositions of cation were fixed to similar composition of Li1.2Ni0.33Ti0.33Mo0.13O2
for every layer along the c direction. One hundreds of Li/Ni/Ti/Mo orderings with the lowest
electrostatic energy were calculated with GGA+U. Among them, the most stable configuration
was selected as the cation-disordered Li1.2Ni0.33Ti0.33Mo0.13O2 structure. The cation-disordered
Li0.467Ni0.37Ti0.37Mo0.15O2-α (α = 0.2) structure was also determined by same technique within a
3 × 3 × 3 supercell of the monoclinic LiMO2 primitive cell. The various oxygen/oxygenvacancy orderings were considered within Li0.467Ni0.33Ti0.33Mo0.13O2-α (α = 0.2) to determine
the structure of Li0.467Ni0.33Ti0.33Mo0.13O1.8 with same technique. The Li/Li-vacancy orderings
were also generated for Li1.2-xNi0.33Ti0.33Mo0.13O2 and Li1.11-xNi0.37Ti0.37Mo0.15O2 with same
methods. Thirty Li/Li-vacancy orderings with the lowest electrostatic energy were calculated
within GGA+U at each composition. The DFT energies of the most stable configurations at
each composition were used to calculate the voltage profile with following equation;
⟨𝑉⟩ = −
E[Lix1 MO2 ]−E[Lix2 MO2 ]−(x1 −x2 )E[Li]
(x1 −x2 )F
77
,
where E is the DFT energy of the structure and F is the Faraday constant.
The oxidation states were determined by comparing calculated magnetizations (average
net moments) of Ni, Ti, Mo and oxygen ions and the number of unpaired spins of Ni2+ (2), Ni3+
(1), Ni4+ (0), Ti3+ (1), Ti4+ (0), Mo5+ (1), Mo6+ (0), O2- (0), and O- (1).111,112 The contributions of
Ni, Ti, Mo and oxygen ions to redox reaction were determined by the change of their
magnetization as follows;
contribution of A= ΔM
ΔMA
,
Ni+ΔMTi+ΔMMo+ΔMO
where ΔM is the change of the magnetization between x1 and x2.
The oxygen loss potential of Li1.2-xNi0.33Ti0.33Mo0.13O2-y was also calculated with
following equation;
⟨𝑉⟩ = −
y
2
𝑒𝑥𝑝
2
E[Lix1 MO2 ]−E[Lix2 MO2−𝑦 ]− (E𝑐 [O2 ]−TSO (300K))−(x1 −x2 )E[Li]
(x1 −x2 )F
,
where E is the DFT energy, Ec[O2] is the corrected DFT energy of O2,113 T is temperature
(300 K) and SO𝑒𝑥𝑝
(300 K) is the entropy of O2 gas at 300 K and 1 atm as obtained from
2
experiments.114,115 We considered two possible transformed structures for the oxygen loss
potential as follows;116
(a) Li0.867Ni0.33Ti0.33Mo0.13O2  0.4 Li + 0.1 O2 + Li0.519Ni0.37Ti0.37Mo0.15O2
: Oxygen loss with lattice densification
(b) Li0.867Ni0.33Ti0.33Mo0.13O2  0.4 Li + 0.1 O2 + Li0.467Ni0.33Ti0.33Mo0.13O1.8
: Oxygen loss with oxygen vacancy formation
78
3.3 Experimental results
3.3.1 Design strategy from percolation theory
Percolation theory predicts that sufficient Li excess is necessary to obtain high capacity from
cation-disordered oxides.70,72 Thus, to design high capacity cation-disordered materials, we
need to find a disordered host structure that can accommodate Li excess. In this work,
LiNi0.5Ti0.5O2 was chosen as the host material for Li excess, containing Ni2+ and Ti4+. This is
because LiNi0.5Ti0.5O2 is a disordered rocksalt phase with an active Ni2+/Ni4+ redox
couple,117,118 which tends to deliver high voltage of ~3.8 V in the oxide cathodes.13,67,119,120
However, LiNi0.5Ti0.5O2 shows limited reversible capacity,117,118 which is due to poor Li
diffusion in its structure by the absence of 0-TM percolation.70,72 Thus, introducing Li excess
for 0-TM percolation can transform the material into promising cation-disordered cathode
materials that not only deliver a high capacity but also high voltage (Fig. 3-1a).
Figures 3-1b to 1d plot the theoretical specific capacities of the hypothetical solid
solution compounds between LiNi0.5Ti0.5O2 and Li1+xM1-xO2 (M = Ti4+, Nb5+, Mo6+), all of
which lead to LiNi0.5Ti0.5O2-based Li-excess cation-disordered materials. Each figure plots
three different capacities: Li capacity that assumes full extraction of available Li ions,
Ni2+/Ni4+ redox capacity, and 0-TM capacity, which is defined as the Li capacity accessible by
the percolating 0-TM network. In all cases, there is a crossover between the Ni2+/Ni4+ redox
capacity and the 0-TM capacity. By a high-valence ‘M’ cation as a charge compensator
(Li1+xM1-xO2), additional Li-excess can be accommodated, making the percolating 0-TM
network more extensive. Hence, the 0-TM capacity increases. However, as excess Li and ‘M’
cation replace some of the Ni sites, the Ni2+/Ni4+ capacity decreases. From Figures 3-1B to 1D,
79
it is seen that charge-compensating excess Li+ with Mo6+ sacrifices the least Ni2+/Ni4+ capacity.
This is because Mo6+ has the highest valency among the charge compensators, thus can
accommodate excess Li with the least amount of Mo6+, preserving the most Ni sites for the
Ni2+/Ni4+ capacity. Therefore, to maintain both the Ni2+/Ni4+ and 0-TM capacity high, we
selected Mo6+ as the charge compensator and investigated the performance of the cationdisordered Li-Ni-Ti-Mo oxides.
Figure 3-3 (a) The crystal structure of disordered rocksalt-type solid solution compounds
between LiNi0.5Ti0.5O2 and Li1+xM1-xO2 (M = Ti4+, Nb5+, Mo6+) (b-d) Theoretical capacities are
given when M is (b) Ti4+ (c) Nb5+ and (d) Mo6+. Each figure plots three different capacities: Li
capacity that assumes full extraction of available Li-ions, Ni2+/Ni4+ redox capacity, and 0-TM
capacity that is the Li capacity accessible by the percolating 0-TM network.
80
3.3.2 Characterization of Li-Ni-Ti-Mo oxides
Figure 3-4 The X-ray diffraction patterns of LiNi0.5Ti0.5O2 (LNTO), Li1.05Ni11/24Ti11/24Mo1/30O2
(LNTMO5), Li1.1Ni5/12Ti5/12Mo1/15O2 (LNTMO10), Li1.15Ni3/8Ti3/8Mo1/10O2 (LNTMO15), and
Li1.2Ni1/3Ti1/3Mo2/15O2 (LNTMO20). Insets are the a-lattice parameters of the compounds.
81
To synthesize the Li-Ni-Ti-Mo oxides (Li1+x/100Ni1/2-x/120Ti1/2-x/120Mox/150O2), we applied
standard solid-state methods as described in the experimental methodology section. Figure 3-2
shows the X-ray diffraction (XRD) patterns of Li1+x/100Ni1/2-x/120Ti1/2-x/120Mox/150O2 (x = 0, 5, 10,
15, 20). Hereafter, LiNi0.5Ti0.5O2 (x = 0) will be referred to as LNTO, and Li1+x/100Ni1/2-x/120Ti1/2x/120Mox/150O2
with x = 5, 10, 15, and 20 will be referred to as LNTMO5, LNTMO10,
LNTMO15, and LNTMO20, respectively. The XRD patterns of a disordered rocksalt in Figure
3-2 and the elemental analysis on the compounds in Table 3-1 show that the target phases are
successfully synthesized (Fig. 3-3, Table 3-2). Insets in Figure 3-2a are the a-lattice parameters
of each compound. The lattice parameter increases slightly with Li excess. This trend is
consistent with the hypothetical Li1.6Mo0.4O2 having bigger average cationic radius (0.726 Å )
than LiNi0.5Ti0.5O2 (0.704 Å ). Thus, introducing excess Li to LiNi0.5Ti0.5O2 by incorporating
Li1.6Mo0.4O2 should increase the a-lattice parameter.
Table 3-1 Target vs. actual Li:Ni:Ti:Mo atomic ratio as determined by direct current plasma
emission spectroscopy.
Li excess
%
0
Target ratio
Li:Ni:Ti:Mo
1:0.5:0.5:0
Actual ratio
Li:Ni:Ti:Mo
0.99:0.51:0.5:0
5
1.05:0.458:0.458:0.033 1.04:0.45:0.47:0.035
10
1.1:0.417:0.417:0.067
1.08:0.42:0.43:0.069
15
1.15:0.375:0.375:0.1
1.15:0.365:0.385:0.1
20
1.2:0.333:0.333:0.133
1.2:0.32:0.35:0.135
82
Figure 3-3 Rietveld refinements on the XRD patterns of (a) LiNi0.5Ti0.5O2 [LNTO], (b)
Li1.05Ni11/24Ti11/24Mo1/30O2 [LNTMO5] (c) Li1.1Ni5/12Ti5/12Mo1/15O2 [LNTMO10], (d)
Li1.15Ni3/8Ti3/8Mo1/10O2 [LNTMO15], and (d) Li1.2Ni1/3Ti1/3Mo2/15O2 [LNTMO20]: structural
parameters from the refinements are listed in Table 3-2.
83
Li excess (%)
Space group
Rwp
Site 4a
(x, y, z) = (0, 0, 0)
Li occupancy
Ni occupancy
Ti occupancy
Mo occupancy
Site 4b
O occupancy
(x, y, z) = (0.5, 0.5, 0.5)
a (Å )
Volume (Å 3)
Derived density (kg/l)
0
5
10
15
20
3.38
0.497
0.262
0.254
0
6.40
0.5160
0.2269
0.2345
0.0163
Fm-3m
3.79
0.5488
0.2109
0.2153
0.0342
3.73
0.5750
0.1899
0.1925
0.0507
2.13
0.5985
0.1582
0.1708
0.0671
1
1
1
1
1
4.1426
71.094
4.39
4.1444
71.186
4.43
4.145
71.216
4.27
4.1451
71.22
4.22
4.1452
71.226
4.11
Table 3-2 Structural parameters from the Rietveld refinements in Figure 3-3: crystallographic
information file of Fm-3m LiFeO2 (ICSD collection code 51208) was first used as an input file.
The atomic occupancies were initially set to the atomic ratio obtained from the elemental
analysis by direct current plasma emission spectroscopy, based on which the lattice parameters
were initially refined. Then, we further refined the lattice parameters and the atomic
occupancies simultaneously: TM occupancies were first refined, and then Li occupancy was
refined. O occupancy did not change after the refinement for all the compounds.
Scanning electron microscopy (SEM) shows that small primary particles, less than 200
nm in diameter (d), are highly agglomerated in secondary particles in all the compounds (Fig.
3-4). The average primary particle size is the smallest for LNTO (d~80 nm) and the largest for
LNTMO20 (d~150 nm). After high-energy ball milling the compounds with carbon black for
the electrode fabrication, the primary particle size becomes slightly less than d~100 nm on
average and the size distribution becomes wider, as can be seen from the image of high-energy
ball milled LNTMO20.
84
Figure 3-4 Scanning electron microscope (SEM) images of LNTO, LNTMO5, LNTMO10,
LNTMO15, LNTMO20, and high-energy ball-milled LNTMO20 (HB-LNTMO20).
85
3.3.3 The electrochemical properties of Li-Ni-Ti-Mo oxides
Figure 3-5 (a) The first-cycle voltage profile of LNTO, LNTMO5, LNTMO10, LNTMO15 and
LNTMO20 [1.5−4.5 V, 20 mA/g], and (b) 20-cycle capacity retention of the compounds.
86
To test the cycling performance of the materials, we performed the galvanostatic chargedischarge tests. Figure 3-5a shows the first-cycle voltage profiles of LNTO, LNTMO5,
LNTMO10, LNTMO15, and LNTMO20 when they are cycled between 1.5-4.5 V at 20 mA/g.
The charge-discharge capacity increases with Li excess from ~110 mAh/g to ~225 mAh/g. The
shape of the voltage curves also evolves with Li excess, with the beginning of the first charge
coming at lower voltage and the 4.3 V plateau becoming longer with higher Li excess, all of
which lead to higher charge capacity. Then, substantial increase in the discharge capacity is
achieved with higher Li excess. The first discharge capacity of LNTO is only 109 mAh/g, but
that of LNTMO20 is as high as 223 mAh/g. Such increase in the reversible capacity with Li
excess is consistent with the increase of 0-TM capacity from percolation theory (Fig. 3-1d).70,72
It is notable that the capacity of LNTMO20 exceeds its theoretical Ni2+/Ni4+ capacity
(= 201.6 mAh/g). This indicates that not only Ni2+/Ni4+ but also other redox couples are active
in LNTMO20. The trend of higher capacity with Li excess continues upon further cycles as
shown in Figure 3-5b.
As LNTMO20 delivers the best performance among the Li-Ni-Ti-Mo oxides, we chose
LNTMO20 as a representative and compared it with LNTO. Figures 3-6a and 6b show the 10cycle voltage profiles of LNTO and LNTMO20 when cycled between 1.5-4.5 V at 20 mA/g.
LNTMO20 delivers much higher capacity (~230 mAh/g) and energy density (~680 Wh/kg,
~2800 Wh/l) than LNTO (~110 mAh/g, ~350 Wh/kg, ~1540 Wh/l). While the capacity above
3 V is higher for LNTMO20, most gains in the discharge capacity come at voltages lower than
3 V, particularly from the ~2.2 V plateau that becomes more obvious with cycles. This results
in the average discharge voltage of ~3 V for LNTMO20. It is notable that the charge-discharge
profile of LNTMO20 is asymmetric with polarization especially near the end of discharge. This
indicates some degree of kinetic limitation in LNTMO20,121 although its performance is still
much better than that of LNTO.
87
Figure 3-6 The voltage profiles for 10 cycles of (a) LNTO, and (b) LNTMO20 [1.5−4.5 V,
20 mA/g, room temperature].
88
Figure 3-7 The voltage profile of (a) LNTO and (b) LNTMO20 when charged and discharged
once at 10 mA/g, and then at 20, 40, 100, 200, and 400 mA/g for the subsequent cycles.
89
Figures 3-7 and 7b show the rate capability of LNTO and LNTMO20, respectively.
Cells made of each compound were charged and discharged once at 10 mA/g, and then at 20,
40, 100, 200, and 400 mA/g for the subsequent cycles. From the profiles, we find that
LNTMO20 delivers higher capacity than LNTO at all rates. As the rate increases from 10 to
400 mA/g, the discharge capacity decreases from 250 mAh/g (750 Wh/kg) to 120 mAh/g
(365 Wh/kg) for LNTMO20 and from 120 mAh/g (366 Wh/kg) to 50 mAh/g (145 Wh/kg) for
LNTO. Note that the capacity of LNTMO20 at 400 mA/g is comparable to that of LNTO at
10 mA/g. This shows that LNTMO20 has improved rate capability than that of LNTO, which
is consistent with the percolation theory.70,72 However, a notable decrease in the capacity of
LNTMO20 with higher rates implies that its rate capability is somewhat limited as well.
To analyze the kinetics in LNTMO20, we performed the galvanostatic intermittent
titration test (GITT). Figure 3-8a shows the first discharge voltage profile of LNTMO20 from
the GITT. Upon first charge to 270 mAh/g and discharge to 270 mAh/g, every step of 9 mAh/g
was galvanostatically charged or discharged at 20 mA/g, and then the test cell was relaxed for
five hours between each step. Time-dependent polarization, inversely correlated with Li
diffusion, is most significant at the end of discharge.122,123 The polarization appears to depend
on the charge cutoff voltage (Fig. 3-8b). When the cutoff voltage is 4.1 V (black solid), the
galvanostatic charge-discharge profiles are symmetric with only minor polarization. When the
material is charged to 4.5 V (blue dash), discharge comes with substantial polarization as in
Figures 3-6b and 8a. This indicates that Li diffusion in LNTMO20 depends on the structural
changes that occur at high voltage.
90
Figure 3-8 (a) The first discharge voltage profile of LNTMO20 from a galvanostatic
intermittent titration test after charging to 270 mAh/g: the inset zooms in the time range
between 240 h and 270 h. (b) The voltage profile of LNTMO20 when charged-discharged three
times between 2.0−4.1 V (black) then between 1.5−4.5 V (blue) at 20 mA/g.
91
3.3.4 The structural evolution of Li1.2Ni1/3Ti1/3Mo2/15O2 during cycling
Figure 3-9 (a) The in situ XRD patterns of LNTMO20 upon two galvanostatic chargedischarge cycles between 1.5−4.8 V at 10 mA/g (b) The corresponding voltage profile (c) and
the a-lattice parameter from single phase XRD refinements are shown. (d) The (002) peak is
zoomed in for the intensity comparison.
We performed in situ X-ray diffraction (XRD) to investigate the structural evolution of
LNTMO20 upon charge and discharge. Figure 3-9a shows the in situ XRD patterns of
LNTMO20 upon two galvanostatic charge-discharge cycles between 1.5−4.8 V at 10 mA/g.
The corresponding voltage profile and the a-lattice parameters from single-phase XRD
refinements are shown in Figures 3-9b and 9c, respectively. Upon first charge, the (002) peak
shifts to a higher angle, indicating a decrease in the a-lattice parameter, but roughly in three
steps. Upon first charge to ~110 mAh/g, which accompanies a sloped voltage profile, the peak
92
continuously shifts to a higher angle. However, further peak shift is negligible in charging to
~215 mAh/g, which occurs at the 4.3 V plateau. After this region, the peak further shifts to a
higher angle with charging. This indicates that the disordered lattice shrinks at the beginning
and end of the first charge, but there is an interval in the middle where it barely shrinks. Upon
first discharge, the (002) peak shifts to a lower angle and gains intensity continuously as in
Figures 3-9a and 9d, indicating Li insertion to LNTMO20 upon discharge. Two features are
seen in the XRD patterns upon first discharge. First, the (002) peak quickly shifts to a lower
angle by discharging to ~100 mAh/g, but further shift is small. Second, after the first discharge,
the peak is at a lower angle (~19.6°) than where it was before cycle (~19.8°), showing the
lattice expansion of LNTMO20 after the first cycle. During the second cycle, the a-lattice
parameter decreases upon charge and increases upon discharge until the 2.2 V plateau is
reached, after which the lattice expansion is small.
3.3.5 Investigation on the redox mechanism
To study the redox mechanism of LNTMO20, we performed X-ray absorption near edge
spectroscopy (XANES) measurements. Figures 3-10a, 10b, and 10c show the Ni K-edge, Ti Kedge, and Mo K-edge XANES spectra of LNTMO20, respectively. Each figure shows spectra
before cycling (black), after the first charge to 4.8 V (blue: ~300 mAh/g charged), and after the
first discharge to 1.5 V (red: ~250 mAh/g discharged). From Figure 3-10a, it is seen that the Ni
edge shifts from an energy close to in LiNi2/3Sb1/3O2 used as a standard for Ni2+ to a higher
energy similar to Ni3+ in NaNiO2 upon first charge to 4.8 V. After the first discharge to 1.5 V,
the Ni edge returns to its starting position.119 This indicates that Ni2+ is oxidized up to Ni~3+
upon first charge to 4.8 V, then reduces back to Ni2+ after the first discharge. As the Ni2+/Ni3+
93
capacity corresponds to ~100 mAh/g, our finding suggests that the remaining charge capacity
comes from either oxygen loss and/or oxygen oxidation, both of which are known to occur in
Li-excess materials.45,97,124,125
Figure 3-10 The X-ray absorption near-edge structures of (a) Ni, (b) Ti, and (c) Mo in
LNTMO20 before cycle [black], after the first charge to 4.8 V [blue, ~300 mAh/g charged],
and after the first discharge to 1.5 V [red, ~250 mAh/g discharged] at 20 mA/g.
94
From the absorption spectra in Figures 3-10b and 10c, it is seen that the Ti and Mo
edges barely shift during charging and discharging, indicating that changes in the Mo and Ti
oxidation states during cycles, if any, are small. However, the pre-edge peak of Mo XANES
increases after the first charge, and remains significantly increased after the first discharge.
This shows that Mo environment has deviated from the regular octahedral coordination during
the cycle, which might originate from a distortion of the Mo-O octahedral or from some degree
of Mo6+ migration from octahedral to tetrahedral sites.126. Comparison with the spectra of
MoO2 and MoO3 shows that the Mo edge position does not shift down in energy after the first
discharge. This observation strongly suggests that the majority of the Mo ions remain 6+.
Likewise, any reduction in the Ti oxidation state on discharge is small on average.
To investigate if oxygen loss occurs from LNTMO20, we performed electron energy
loss spectroscopy (EELS) on the surface of the LNTMO20 particles before and after cycling.
Figure 3-11a shows the Ti L-edge and O K-edge from the EELS spectrums of LNTMO20
before cycling (black) and after 20 cycles (red) between 1.5-4.5 V at 20 mAh/g. Comparing the
EELS quantifications of the atomic ratio of O to Ti, we find a considerable decrease in the ratio
by ~39 % after cycling. This indicates that oxygen loss has occurred from the surface of
LNTMO20 upon cycling, which can contribute to additional charge capacity beyond the
Ni2+/Ni~3+ capacity. In addition, we observe that the Ti L-edge is chemically shifted to low
energy direction by ~1.5 eV relative to O K-edge after cycling (the inset in Fig. 3-11a). This
indicates Ti reduction below 4+ at the surface region,127 which is not captured by the Ti
XANES.
95
Figure 3-11 Electron energy loss spectra of Ti L-edge and O K-edge in LNTMO20 before
cycling [black] and after 20 cycles between 1.5−4.5 V [red] at 20 mA/g: the inset focuses on
the Ti L-edge.
Figure 3-12 The first-cycle cyclic voltammetry profiles of LNTMO20 when voltage-swept
between 1.5−4.5 V [black solid] and 1.5−4.1 V [red dash] at 0.1 mV/s.
96
Oxygen loss from LNTMO20 can also be inferred from the cyclic voltammetry (CV)
tests. Figure 3-12b shows the first cycle CV profile of LNTMO20. When the oxidation cutoff
voltage is 4.1 V (red), we observe a main reduction peak at ~3.7 V and a minor reduction peak
at ~2.7 V. However, when the cutoff is increased to 4.5 V (black), an additional reduction peak
at ~2.2 V is observed in the CV profile, which is likely associated with reduction of a second
TM species. This shows that charging above 4.1 V triggers a reaction which allows reduction
of species that were previously not reducible. In the case of LNTMO20, reduction of Mo6+ or
Ti4+ upon discharge is likely triggered by oxygen loss, similar to reduction of Mn4+ that
becomes possible in Li-excess Ni-Mn-Co oxides after oxygen loss.40,97,128 Although the Mo or
Ti XANES do not show a clear evidence of the decrease in the average Mo or Ti oxidation
states after the first discharge (Figs. 3-10b, 10c), the apparent discrepancy between CV (or
EELS) and XANES implies that oxygen loss may be significant at the surface but not in the
bulk. Otherwise, XANES should detect the overall decrease in the average Mo or Ti oxidation
state after the first discharge.
Based on the information from the XANES spectra, we can approximate the limit for
the oxygen loss capacity of LNTMO20 during the first cycle. The Ni XANES shows that Ni2+
is oxidized to Ni~3+ upon first charge to 4.8 V, which gives ~100 mAh/g in capacity (Fig. 310a). The remaining first charge capacity (~200 mAh/g) can originate from both oxygen loss
and oxygen oxidation. Depending on the oxygen loss capacity, the average Mo or Ti oxidation
state should change after the first charge to 300 mAh/g then discharge to 250 mAh/g (Fig. 313). For example, if we assume uniform oxygen loss in a LNTMO20 particle, no loss in the
TM content, and that Ni, Ti, and O remain as Ni2+, Ti4+, and O2- after the first discharge, the
average Mo oxidation state after the first discharge should be 5.92+, 5.53+, 5.13+, and 4.35+
with increasing oxygen loss capacity of 70, 90, 110, and 150 mAh/g, respectively (Fig. 3-13).
According to the Mo and Ti XANES, neither Mo nor Ti is reduced on average after the first
97
discharge (Figs. 3-10b, 10c). Thus, it is likely that oxygen loss does not account for all the
extra capacity beyond the Ni2+/Ni3+ capacity, possibly less than 90 mAh/g, and the remaining
first charge capacity originates from oxygen oxidation. However, note that the local oxygenloss level can be different between the surface and the bulk of a LNTMO20 particle: our
estimation is on the overall oxygen loss capacity.
Figure 3-13 Back-of-the-envelope calculation: Average Mo oxidation state expected after the
first charge to 300 mAh/g then discharge to 250 mAh/g as a function of the oxygen-loss
capacity during the first charge. It was assumed that Ni, Ti, and O stay as Ni2+, Ti4+, and O2after the first discharge, and that there is no loss in the TM content during the first cycle.
98
3.4 Computational results
3.4.1 Oxygen loss mechanisms
As we have discussed, oxygen loss occurs from LNTMO20, which likely is most substantial
near the surface of the LNTMO20 particles (Fig. 3-11). Two oxygen loss mechanisms have
been proposed in the literature. When oxygen is released from the particle surface, either (i)
oxygen vacancies or (ii) under-coordinated TM ions at the surface may diffuse into the bulk of
the crystal structure.97,124,129,130 The former mechanism introduces oxygen vacancies in the bulk
lattice after oxygen loss.97 The latter mechanism results in an increased TM content in the bulk,
and is therefore commonly referred to as lattice densification.124,129,130
To study which mechanism accompanies oxygen loss from LNTMO20, we performed
DFT calculations on the densified Li0.52Ni0.37Mo0.15O2 structures and Li0.47Ni0.33Ti0.33Mo0.13O1.8
structures with oxygen vacancies, both of which represent possible structures after partially
delithiated LNTMO20 loses oxygen by following equations.
(a) Li0.867Ni0.33Ti0.33Mo0.13O2  0.4 Li + 0.1 O2 + Li0.519Ni0.37Ti0.37Mo0.15O2
: Oxygen loss with lattice densification
(b) Li0.867Ni0.33Ti0.33Mo0.13O2  0.4 Li + 0.1 O2 + Li0.467Ni0.33Ti0.33Mo0.13O1.8
: Oxygen loss with oxygen vacancy formation
Based on our calculations, we find that the densified Li0.519Ni0.37Ti0.37Mo0.15O2
structures are energetically more stable than the Li0.467Ni0.33Ti0.33Mo0.13O1.8 structures with
oxygen vacancies even though they have same composition of supercell as Li 14Ti10Ni10Mo4O54:
The energy difference is 105 meV per LixMO2 formula unit (Fig. 3-14). Thus, the oxygen loss
99
potential in Figure 3-15a is derived from the reaction (a): oxygen loss with lattice densification.
Figure 3-14 The calculated energies of densified Li0.52Ni0.37Ti0.37Mo0.15O2 structures and
Li0.47Ni0.33Ti0.33Mo0.13O1.8 structures with oxygen vacancies: From the plot, it is seen that the
energies of the densified Li0.52Ni0.37Ti0.37Mo0.15O2 structures are lower than those of
Li0.47Ni0.33Ti0.33Mo0.13O1.8 structures with oxygen vacancies. This indicates that oxygen loss
with densification is thermodynamically more favorable than that with oxygen vacancies in the
lattice.
100
3.4.2 Calculated voltage profiles from possible redox mechanisms
Figure 3-15 (a) The voltage profiles of Li1.2-xNi0.33Ti0.33Mo0.13O2: The black curve is an
experimental profile from the galvanostatic intermittent titration test (GITT) during the first
charge, and the red line is calculated voltage profile of Li1.2-xNi0.33Ti0.33Mo0.13O2 that assumes
no oxygen loss during the first charge: the dotted red arrows specify the region of Ni oxidation
and O oxidation. Finally, the dashed blue line indicates the oxygen-loss potential of Li1.2xNi0.33Ti0.33Mo0.13O2-y. (b) The calculated contribution of Ni [black] and oxygen [red] oxidation
upon delithiation in Li1.2-xNi0.33Ti0.33Mo0.13O2 when no oxygen loss is assumed (the red curve in
Figure 3-15a)
101
To further understand the behavior of Li1.2-xNi0.33Ti0.33Mo0.13O2 (LNTMO20) during cycling,
we investigated the redox mechanism in LNTMO20 with density functional theory (DFT)
calculations. Figure 3-15a compares the computed first charge voltage profiles of LNTMO20
against the experimental profile. The red curve is the computed profile assuming no oxygen
loss in the crystal structure of LNTMO20 upon first charge. The blue line is the computed
profile assuming that partially-delithiated LNTMO20 (Li0.867Ni0.33Ti0.33Mo0.13O2) loses oxygen
(−0.1 O2) upon further charge (−0.4 Li), leading to a densified disordered phase of
Li0.52Ni0.37Mo0.15O2 that is thermodynamically more stable than a disordered phase with
oxygen vacancies (Li0.47Ni0.33Ti0.33Mo0.13O1.9) (Fig. 3-14). Finally, the black curve is the
experimental profile from GITT. Figure 3-15b shows the contribution of Ni and oxygen
oxidation upon delithiation in LNTMO20 when no oxygen loss is assumed, which results in the
red profile in the Figure 3-15a.
In general, the experimental voltage (black) is lower than the calculated voltage without
oxygen loss (red) and higher than the calculated voltage with oxygen loss (blue). The much
lower (calculated) voltage from oxygen loss than without oxygen loss shows that there is a
clear driving force for oxygen loss from LNTMO20 after a certain level of delithiation. The
calculated voltage with oxygen loss is quite lower that the experimental profile, implying that
the actual first charge mechanism is likely to include some degree of oxygen loss based on the
thermodynamic driving force. Note that the oxygen loss mechanism in experiments can be
more complicated than what is assumed in our calculations. Therefore, there can be a
discrepancy in the oxygen loss voltage between the calculations and experiments.
Although there exists a clear driving force for oxygen loss after a certain level of
delithiation, the computed voltage without oxygen loss (thus with Ni and O oxidation only) is
within the experimental voltage window of 1.5−4.5 V (or 1.5−4.8 V). Thus, not to mention the
Ni oxidation which is experimentally confirmed by XANES (Fig. 3-10a), oxygen oxidation can
102
also take place at high voltage along with oxygen loss. It is notable that the oxygen oxidation is
predicted to take place before Ni2+ becomes completely oxidized to Ni4+ (Fig. 3-15b). In theory,
full Ni2+/Ni4+ capacity allows for 0.66 Li delithiation. However, oxygen oxidation already takes
place after ~0.35 Li delithiation, which demonstrates an overlap between Ni 3d and O 2p bands
in LNTMO20.
3.4.3 Reduction of Mo and Ti after oxygen loss
After oxygen loss, not only Ni3+, Ni4+, and O-, but also Mo6+ and Ti4+ can be reduced upon
discharge. To study which species (Mo6+ vs. Ti4+) is first reduced after oxygen loss, we
computationally studied the redox mechanism of Li1.11-xNi0.37Ti0.37Mo0.15O2 which represents a
densified disordered phase of LNTMO20 after some degree of oxygen loss. Figure 3-16a
shows the voltage profile of Li1.11-xNi0.37Ti0.37Mo0.15O2, and Figure 3-16b shows the average net
moments of Ni, Ti, Mo and oxygen ions in Li1.11-xNi0.37Ti0.37Mo0.15O2 at each composition. The
average net moment of Ni ions continuously increases from 1.06 to 1.77 as x decreases from
1.04 to 0.22, indicating reduction of Ni~3+ to Ni2+ in average. Simultaneously, the average net
moment of O ions decreases from 0.17 to 0.03 as x decreases from 1.04 to 0.44, indicating
reduction of O~1.8- to O2- in average. Although the net moments of Ti ions remain unchanged
for whole compositions in Li1.11-xNi0.37Ti0.37Mo0.15O2, the average net moment of Mo ions
rapidly increases as x decreases from 0.22 to 0. This indicates that Mo reduction occurs before
Ti reduction but after O and Ni reduction. Thus, for the Ti reduction to occur after oxygen loss,
it is likely that Mo reduction has already taken place to some degree.
103
Figure 3-16 (a) The voltage profiles of Li1.11-xNi0.37Ti0.37Mo0.15O2. (b) The average net moments
of Ni, Ti, Mo and oxygen ions in Li1.11-xNi0.37Ti0.37Mo0.15O2 (x = 0, 0.074, 0.148, 0.222, 0.296,
0.370, 0.444, 0.519, 0.593, 0.667, 0.741, 0.815, 0.889, 0.963 and 1.037).
104
3.5 Discussions
3.5.1 Redox mechanism
When we designed the Li-Ni-Ti-Mo oxides, full Ni2+/Ni4+ oxidation was assumed to be
possible. In experiments, LNTMO20 can be charged beyond the full Ni2+/Ni4+ capacity.
However, the Ni XANES suggests that Ni2+ can be oxidized to only up to Ni~3+ (Fig. 3-10a),
and the rest of the first charge capacity originates from either oxygen loss or/and oxygen
oxidation. Because the Mo and Ti XANES do not show significant reduction of the average
Mo and Ti oxidation state after the first discharge, we suspect that oxygen loss does not
account for all the extra capacity beyond the Ni2+/Ni~3+ capacity, and oxygen oxidation may be
responsible for the remaining charge capacity. Based on this understanding, we propose the
following first charge mechanism for LNTMO20 as a representative of the Li-Ni-Ti-Mo oxides:
after the Ni2+/Ni~3+ oxidation, oxygen loss mainly occurs until the surface becomes passivated
against the oxygen loss, and then oxygen oxidation dominantly takes place at higher voltage.
Note that clear distinction between the oxygen loss region and the oxygen oxidation region
may not exist as both can happen simultaneously, such as oxygen loss at the surface and
oxygen oxidation in the bulk.129,130
The proposed mechanism is consistent with the change in the lattice parameter of
LNTMO20 during the first charge (Fig. 3-17). Upon first charge to ~110 mAh/g, the lattice
parameter decreases continuously. This can be explained with the Ni2+/ Ni~3+ oxidation
(~100 mAh/g) because Ni3+ (r = 0.56 Å ) and Ni4+ (r = 0.48 Å ) are smaller than Ni2+ (r = 0.69
Å ), thus the lattice parameter decreases with the smaller ions. Upon further charge to
~215 mAh/g, the lattice parameter barely decreases. This can be related to oxygen loss because
105
charging with oxygen loss slows down the increase in the oxidation state of the remaining ions
in the crystal structure. As the oxidation state thus the size of the remaining ions stays similar,
the lattice shrink can be retarded.30,39 Note that the capacity from this region is ~105 mAh/g,
which roughly agrees with our maximum estimated oxygen loss capacity (~90 mAh/g) from
the XANES results. Finally, charging beyond ~215 mAh/g decreases the lattice parameter. This
can be explained by oxygen oxidation that can shrink the oxygen framework either by making
the oxygen ions smaller in size or by introducing peroxo-like species whose oxygen-to-oxygen
bond distance is shorter.45,125
Although the Mo and Ti XANES do not show clear evidence of Mo or Ti reduction in
LNTMO20 after the first discharge (Figs. 3-10b, 10c), we believe that after oxygen loss, Mo6+
and Ti4+ can be reduced upon discharge in addition to Ni3+, Ni4+ and O−, especially near the
surface. First, EELS Ti Li-edge collected from the surface region shows a chemical shift to low
energy direction by ~1.5 eV relative to O K-edge after cycling, indicating Ti reduction below
4+ (Fig. 3-11).127 Furthermore, the activation of the Mo or Ti redox couples can be inferred
from the CV test, which shows reduction of an additional species when the oxidation cutoff
voltage is increased to 4.5 V and LNTMO20 shows oxygen loss (Fig. 3-12). Finally,
LNTMO20 delivers ~50 % of its discharge capacity below 3 V during the 1.5−4.5 V cycling
test (Fig. 3-6), and its discharge plateau at ~2.2 V matches to Ti4+ reduction in the literature.117
As our calculations predict Mo6+ reduction to occur before Ti4+ reduction (Fig. 3-16), it is
likely that oxygen loss allows for both Mo6+ and Ti4+ reduction upon discharge, particularly
near the surface. The apparent discrepancy between EELS/CV (showing reduction of Mo6+ or
Ti4+) and XANES (showing no change in Mo6+ or Ti4+) may be explained by significant
oxygen loss near the surface but not in the bulk. More careful studies to characterize the redox
activity are underway.
106
Figure 3-17 Proposed first charge mechanism of LNTMO20: Ni2+/Ni~3+ oxidation, oxygen loss,
and oxygen oxidation largely in sequence. The voltage profile of LNTMO20 from the in situ
XRD test and the corresponding a-lattice parameters from single-phase XRD refinements are
overlayed.
3.5.2 Electrochemical performance
As the percolation theory predicts (Fig. 3-1d), the reversible capacity and the rate capability
improves with Li excess. In particular, LNTMO20 delivers high capacity and energy density
(250 mAh/g, 750 Wh/kg, 3080 Wh/l) at 10 mA/g, which is double the capacity and energy
density of LNTO (120 mAh/g, 366 Wh/kg, 1610 Wh/l) at the same rate. Nevertheless, Li
diffusion is still somewhat limited in LNTMO20, resulting in large polarization and limited
rate capability (Fig. 3-8a). This can also be explained with the structural changes from oxygen
loss.
107
As we have discussed, oxygen loss occurs from LNTMO20, which likely is most
substantial near the surface of the LNTMO20 particles (Fig. 3-11). Two oxygen loss
mechanisms have been proposed in the literature. When oxygen is released from the particle
surface, either (i) oxygen vacancies or (ii) under-coordinated TM ions at the surface may
diffuse into the bulk of the crystal structure.97,124,129,130 The former mechanism introduces
oxygen vacancies in the bulk lattice after oxygen loss.97 The latter mechanism results in an
increased TM content in the bulk, and is therefore commonly referred to as lattice
densification.124,129,130
Based on our calculations and the literatures, oxygen loss with lattice densification is
thermodynamically more favorable than oxygen loss with oxygen vacancies in the lattice (Fig.
3-14).124,129,130 Therefore, it is likely that the surface of LNTMO20 becomes densified after
oxygen loss (Fig. 3-18). This can impede Li diffusion in LNTMO20 because densification
lowers the Li-excess level by increasing the TM content, resulting in poorer 0-TM percolation
in the disordered structure (Fig. 3-1d).70,72 For example, the EELS measurement on the surface
of LNTMO20 particles shows a considerable decrease (~39 %) in the O/Ti intensity ratio after
the cycling between 1.5−4.5 V (Fig. 3-11). If we assume no loss in the TM content upon
oxygen loss with densification, the decrease in the ratio by ~39 % can be interpreted as the
change in the composition at the surface from Li1.2-xTM0.8O2 (20 % Li excess) to Li0.7-xTM1.3O2
(−30 % Li excess), which is well below the threshold for 0-TM percolation (~9 % Li excess).
This demonstrates how greatly the Li-excess level can decrease at the surface by oxygen loss,
thus 0-TM percolation becomes poorer after oxygen loss. While further work is necessary to
clearly confirm this hypothesis, it is consistent with Li transport becoming limited when
LNTMO20 is charged above ~110 mAh/g, the threshold capacity after which oxygen loss
occurs (Fig. 3-8b).
108
Figure 3-18 Illustrations of a LNTMO20 particle before and after oxygen loss with
densification near the surface: oxygen loss with densification reduces the Li-excess level, and
thus decreases the 0-TM capacity.
3.5.3 Strategies for improvements
The observation that Ni2+ can be oxidized up to only Ni~3+ indicates a great overlap between
the Ni 3d and O 2p bands in LNTMO20. Variations in local environments for Ni and O in the
disordered structure may result in varying orbital overlap, reversing some Ni and O states in
energy (Fig. 3-15). In such a scenario, high capacity could not be achieved without oxygen loss
or oxidation. While oxygen oxidation can be beneficial because it delivers capacity at high
voltage,45,125,130 oxygen loss with densification can be detrimental because it degrades Li
diffusion by lowering the Li-excess level in the disordered structure. Therefore, avoiding the
oxygen loss seems necessary for any disordered Li-excess materials. Here, we propose two
different approaches to avoid oxygen loss from the Li-Ni-Ti-Mo oxides and thus improve the
materials.
109
First, decreasing the lattice parameter by cation substitution can be beneficial. This is
because a smaller lattice parameter results in a greater orbital overlap between the Ni 3d and O
2p orbitals, which increases the covalency of the Ni-O bonding.7,12 As the covalency increases,
the Ni 3d band (eg* band) which has anti-bonding characteristic becomes higher in energy,
resulting in less band overlap between the Ni 3d and O 2p bands. This can maximize the
Ni2+/Ni4+ capacity so that oxygen loss or oxidation is not required to achieve high capacity. For
example, if it were possible to completely use the Ni2+/Ni4+ capacity, LNTMO20 could have
delivered ~200 mAh/g without oxygen loss or oxidation at all.
Second, surface coating can be beneficial.131,132 If surface coating can retard oxygen
loss, oxygen oxidation will immediately take over instead. In this way, high capacity and
energy density may be achieved only by the Ni redox and O redox, bypassing the oxygen loss
with densification. Thus, the 0-TM percolation stays intact for facile Li diffusion in the Li-NiTi-Mo oxides.
110
3.6 Conclusion
In conclusion, we have designed a new class of high-capacity cation-disordered oxides: Li-NiTi-Mo oxides. As 0-TM percolation theory predicts, the reversible capacity and the rate
capability improve with Li excess. In particular, Li1.2Ni1/3Ti1/3Mo2/15O2 (20 % Li excess)
delivers up to 250 mAh/g and 750 Wh/kg (~3080 Wh/l) at 10 mAh/g, which is double the
capacity and energy density of LiNi1/2Ti1/2O2 (0 % Li excess) at the same rate. Through a
combination of the in situ XRD, XAS, EELS, and electrochemistry, we propose that first
charging Li1.2Ni1/3Ti1/3Mo2/15O2 to 4.8 V is accompanied by Ni2+/Ni~3+ oxidation, oxygen loss,
and oxygen oxidation largely in this sequence, after which Mo6+ and/or Ti4+ can be reduced
upon discharge. Furthermore, we argue that oxygen loss with densification can impede Li
diffusion in Li-Ni-Ti-Mo oxides especially near the surface, because densification lowers the
Li-excess level, resulting in the poorer 0-TM percolation in the disordered materials. Finally,
we proposed that preventing oxygen loss will improve the new disordered materials by
preserving the 0-TM percolation for facile Li diffusion. We believe that the line of thought in
this work provides important guidelines for the design of high energy density cation-disordered
cathode materials for rechargeable lithium batteries.
111
-------------Part IV-------------
Conclusions
112
To develop high-performance Li-ion batteries, the hunt for cathode materials with high-energy
densities has been intense. This is because such cathode materials will enable smaller and
lighter Li-ion batteries for various complex applications, such as for electric vehicles and grid
energy storage.
Most high-energy density cathodes have been sought from well-ordered close-packed
oxides such as layered (rocksalt-type) Li-TM oxides, and there has been a common wisdom
that a well-ordered crystal structure is necessary for facile Li diffusion in the cathode materials,
leading to good cyclability of the materials. Accordingly, cation-disordered (i.e. disordered
rocksalt) Li-TM oxides have received only a limited attention as cathode materials.
In this thesis, it was demonstrated that cation-disordered Li-TM oxides can be
promising cathode (or electrode) materials, which enlarges the search space for high-energy
density cathode materials to cation-disordered oxides.
In the first part of this thesis, the potential of the cation-disordered oxides was
demonstrated by the counterintuitive performance of Li1.211Mo0.467Cr0.3O2. This material forms
into a well-layered structure but transforms to a cation-disordered structure after several
charge-discharge cycles. While a common wisdom expects poor cycling performance of this
material, this material delivers a very high capacity (~265 mAh/g) and energy density
(~660 Wh/kg), which is very rarely achieved even by the well-layered materials. Using ab
initio computations, we explained that the counterintuitive behavior originates from percolation
of active 0-TM channels in disordered Li-excess materials. When diffusing through a 0-TM
channel, an activated Li+ ion does not face-share high-valent transition metal ions. Therefore,
the diffusion barrier through the channel can stay low (thus allows for facile Li migration) even
in a cation-disordered structure in which the relaxation space for the activated Li+ ion against
face-sharing species is much smaller than in a layered structure: other types of channels are
inactive in the cation-disordered structure due to the small relaxation space. Nevertheless, for
113
the active 0-TM channels to support macroscopic lithium diffusion in the cation-disordered
materials, the number of 0-TM channels must be great enough for the channels to percolate in
the cation-disordered structure, which requires a Li-excess composition (x > 1.09 in LixTM2xO2). As
the Li-excess level increases, the percolating network of 0-TM channels become more
extensive, enabling a higher fraction of Li ions in a disordered structure to cycle through the
network. The Li-excess level of Li1.211Mo0.467Cr0.3O2 is x ~ 1.233 (~23 % Li excess) which is
high enough for the percolating 0-TM network to allow for as much as ~1 Li per f.u.
(~265 mAh/g) to cycle through the network. Therefore, even as the material disorders during
cycling, the materials can still cycle as much as ~1 Li per f.u. through the percolating 0-TM
network.
Based on the above understanding, a new class of high capacity cation-disordered
oxides (Li-Ni-Ti-Mo oxides) was designed in the second part of this thesis. The Li-Ni-Ti-Mo
oxides are Li-excess cation-disordered oxides with a general composition of Li1+x/100Ni1/2x/120Ti1/2-x/120Mox/150O2
(0 < x < 30). As the 0-TM percolation theory predicts, the reversible
capacity and rate capability improve with Li excess. In particular, Li1.2Ni1/3Ti1/3Mo2/15O2 (20 %
Li excess) delivers a high capacity and energy density up to 250 mAh/g and 750 Wh/kg (~3080
Wh/l) at 10 mA/g, which is double the capacity and energy density of LiNi0.5Ti0.5O2 (0 % Li
excess) at the same rate. However, large polarization in the voltage profile and limited rate
capability from Li1.2Ni1/3Ti1/3Mo2/15O2 indicate that lithium transport needs to be further
improved. Combining various characterization techniques, it was proposed that first charging
Li1.2Ni1/3Ti1/3Mo2/15O2 to 4.8 V is accompanied by Ni2+/Ni~3+ oxidation, oxygen loss, and
oxygen oxidation largely in this sequence, after which both Mo6+ and Ti4+ can be reduced upon
discharge. Furthermore, we argued that oxygen loss with densification can impede lithium
diffusion in the Li-Ni-Ti-Mo oxides especially near the surface, because densification lowers
the Li-excess level, leading to poorer 0-TM percolation in the disordered materials. Finally, we
114
proposed that preventing oxygen loss will improve the new disordered materials by preserving
the 0-TM percolation.
The search of high-capacity cation-disordered materials has just started. While much
effort will be necessary for these materials to become ultimate cathode materials that replace
existing materials, I believe the work in this thesis has unlocked the potential of cationdisordered oxides for rechargeable lithium battery cathodes, thus offers hope for substantial
improvement in the performance of rechargeable lithium batteries.
115
116
References
1.
Tarascon, J. M. & Armand, M. Issues and challenges facing rechargeable lithium
batteries. Nature 414, 359–67 (2001).
2.
Dunn, B., Kamath, H. & Tarascon, J.-M. Electrical Energy Storage for the Grid: A
Battery of Choices. Science 334, 928–935 (2011).
3.
Denholm, P., Ela, E., Kirby, B. & Milligan, M. The Role of Energy Storage with
Renewable Electricity Generation The Role of Energy Storage with Renewable
Electricity Generation. 1–61 (2010).
4.
Etacheri, V., Marom, R., Elazari, R., Salitra, G. & Aurbach, D. Challenges in the
development of advanced Li-ion batteries: a review. Energy Environ. Sci. 3243–3262
(2011).
5.
Ceder, G. Opportunities and challenges for first-principles materials design and
applications to Li battery materials. MRS Bull. 35, 693–702 (2010).
6.
Armand, M. & Tarascon, J.-M. Building better batteries. Nature 451, 652–657 (2008).
7.
Goodenough, J. B. & Kim, Y. Challenges for Rechargeable Li Batteries. Chem. Mater.
22, 587–603 (2010).
8.
Whittingham, M. S. Lithium batteries and cathode materials. Chem. Rev. 104, 4271–301
(2004).
9.
Goodenough, J. B. & Park, K. S. The Li-ion rechargeable battery: A perspective. J. Am.
Chem. Soc. 135, 1167–1176 (2013).
10.
Aravindan, V., Gnanaraj, J., Madhavi, S. & Liu, H.-K. Lithium-Ion Conducting
Electrolyte Salts for Lithium Batteries. Chemistry. Eur. J. 17, 14326−14346 (2011).
11.
Aurbach, D. et al. Design of electrolyte solutions for Li and Li-ion batteries: A review.
Electrochim. Acta 50, 247–254 (2004).
12.
Aydinol, M. K., Kohan, A. F. & Ceder, G. Ab initio study of lithium intercalation in
metal oxides and metal dichalcogenides. Phys. Rev. B 56, 1354–1365 (1997).
13.
Kang, K., Meng, Y. S., Bréger, J., Grey, C. P. & Ceder, G. Electrodes with high power
and high capacity for rechargeable lithium batteries. Science 311, 977–980 (2006).
14.
Kang, B. & Ceder, G. Battery materials for ultrafast charging and discharging. Nature
458, 190–193 (2009).
15.
Nykvist, B. & Nilsson, M. Rapidly falling costs of battery packs for electric vehicles.
Nat. Clim. Chang. 5, 100–103 (2015).
117
16.
Hautier, G. et al. Phosphates as Lithium-Ion Battery Cathodes : An Evaluation Based on
High-Throughput ab Initio Calculations. Chem. Mater. 23, 3495−3508 (2011).
17.
Jain, A., Hautier, G., Ong, S. P., Dacek, S. & Ceder, G. Relating Q1 Q2 voltage and
thermal safety in Li-ion battery cathodes: a high-throughput computational study. Phys.
Chem. Chem. Phys. 17, 5942–5953 (2015).
18.
Yoon, W.-S., Grey, C. P., Balasubramanian, M., Yang, X.-Q. & McBreen, J. In Situ Xray Absorption Spectroscopic Study on LiNi0.5Mn0.5O2 Cathode Material during
Electrochemical Cycling. Chem. Mater. 15, 3161–3169 (2003).
19.
Ohzuku, T. Electrochemistry and Structural Chemistry of LiNiO2 (R-3m) for 4 Volt
Secondary Lithium Cells. J. Electrochem. Soc. 140, 1862 (1993).
20.
Dahn, J., Fuller, E., Obrovac, M. & Vonsacken, U. Thermal stability of LixCoO2,
LixNiO2 and λ-MnO2 and consequences for the safety of Li-ion cells. Solid State Ionics
69, 265–270 (1994).
21.
Dahn, J. R., Vonsacken, U., Juzkow, M. W. & Aljanaby, H. Rechargeable LiNiO2
Carbon Cells. J. Electrochem. Soc. 138, 2207–2211 (1991).
22.
Kobayashi, H., Arachi, Y., Emura, S. & Tatsumi, K. Investigation on lithium deintercalation mechanism for Li1-yNi1/3Mn1/3Co1/3O2. J. Power Sources 146, 640–644
(2005).
23.
Choi, J. & Manthiram, a. Role of Chemical and Structural Stabilities on the
Electrochemical Properties of Layered LiNi1/3Mn1/3Co1/3O2 Cathodes. J. Electrochem.
Soc. 152, A1714 (2005).
24.
Bréger, J. et al. Effect of High Voltage on the Structure and Electrochemistry of
LiNi0.5Mn0.5O2: A Joint Experimental and Theoretical Study. Chem. Mater. 18, 4768–
4781 (2006).
25.
Kim, D. K. et al. Spinel LiMn2O4 nanorods as lithium ion battery cathodes. Nano Lett. 8,
3948–3952 (2008).
26.
Shaju, K. M. & Bruce, P. G. A stoichiometric nano-LiMn2O4 spinel electrode exhibiting
high power and stable cycling. Chem. Mater. 20, 5557–5562 (2008).
27.
Cho, J., Kim, T.-J., Kim, Y. J. & Park, B. Complete blocking of Mn3+ ion dissolution
from a LiMn2O4 spinel intercalation compound by Co3O4 coating. Chem. Commun.
1074–1075 (2001).
28.
Van Der Ven, A., Marianetti, C., Morgan, D. & Ceder, G. Phase transformations and
volume changes in spinel LixMn2O4. Solid State Ionics 135, 21–32 (2000).
29.
Yang, M.-C. et al. Electronic, Structural, and Electrochemical Properties of LiNixCuyMn
2–x–yO4 (0 < x < 0.5, 0 < y < 0.5) High-Voltage Spinel Materials. Chem. Mater. 23,
2832–2841 (2011).
118
30.
Ma, X., Kang, B. & Ceder, G. High Rate Micron-Sized Ordered LiNi0.5Mn1.5O4. J.
Electrochem. Soc. 157, A925 (2010).
31.
Kang, B. & Ceder, G. Electrochemical Performance of LiMnPO4 Synthesized with OffStoichiometry. J. Electrochem. Soc. 157, A808 (2010).
32.
Chung, S.-Y., Bloking, J. T. & Chiang, Y.-M. Electronically conductive phosphoolivines as lithium storage electrodes. Nat. Mater. 1, 123–8 (2002).
33.
Chung, S.-Y. & Chiang, Y.-M. Microscale Measurements of the Electrical Conductivity
of Doped LiFePO4. Electrochem. Solid-State Lett. 6, A278 (2003).
34.
Delmas, C., Maccario, M., Croguennec, L., Le Cras, F. & Weill, F. Lithium
deintercalation in LiFePO4 nanoparticles via a domino-cascade model. Nat. Mater. 7,
665–71 (2008).
35.
Wang, F. et al. Conversion Reaction Mechanisms in Lithium Ion Batteries: Study of the
Binary Metal Fluoride Electrodes. J. Am. Chem. Soc. 133, 18828−18836 (2011).
36.
Ong, S. P., Wang, L., Kang, B. & Ceder, G. Li-Fe-P-O2 Phase Diagram from First
Principles Calculations. Society 77, 1798–1807 (2008).
37.
Ong, S., Chevrier, V. & Ceder, G. Comparison of small polaron migration and phase
separation in olivine LiMnPO4 and LiFePO4 using hybrid density functional theory.
Phys. Rev. B 83, 1–7 (2011).
38.
Wang, L., Maxisch, T. & Ceder, G. A first-principles approach to studying the thermal
stability of oxide cathode materials. Chem. Mater. 19, 543–552 (2007).
39.
Bervas, M., Yakshinskiy, B., Klein, L. C. & Amatucci, G. G. Soft-Chemistry Synthesis
and Characterization of Bismuth Oxyfluorides and Ammonium Bismuth Fluorides. J.
Am. Ceram. Soc. 89, 645–651 (2006).
40.
Thackeray, M. M. et al. Li2MnO3-stabilized LiMO2 (M = Mn, Ni, Co) electrodes for
lithium-ion batteries. J. Mater. Chem. 17, 3112 (2007).
41.
Johnson, C. S., Li, N., Lefief, C., Vaughey, J. T. & Thackeray, M. M. Synthesis,
Characterization and Electrochemistry of Lithium Battery Electrodes: xLi2MnO3·(1−x)
LiMn0.333Ni0.333Co0.333O2 (0 ≤ x ≤ 0.7). Chem. Mater. 20, 6095–6106 (2008).
42.
Li, J. et al. Synthesis and Characterization of the Lithium-Rich Core-Shell Cathodes
with Low Irreversible Capacity and Mitigated Voltage Fade. Chem. Mater.27,
3366−3377 (2015).
43.
Cho, J., Kim, Y.-W., Kim, B., Lee, J.-G. & Park, B. A breakthrough in the safety of
lithium secondary batteries by coating the cathode material with AlPO4 nanoparticles.
Angew. Chem. Int. Ed. Engl. 42, 1618–1621 (2003).
119
44.
Sun, Y.-K. et al. A novel concentration-gradient Li[Ni0.83Co0.07Mn0.10]O2 cathode
material for high-energy lithium-ion batteries. J. Mater. Chem. 21, 10108 (2011).
45.
Sathiya, M. et al. Reversible anionic redox chemistry in high-capacity layered-oxide
electrodes. Nat. Mater. 12, 827–35 (2013).
46.
Lee, H. W. et al. Ultrathin spinel LiMn2O4 nanowires as high power cathode materials
for Li-ion batteries. Nano Lett. 10, 3852–3856 (2010).
47.
Li, L. & Cells, M. Dissolution of Spinel Oxides and Capacily Losses in 4 V. Society 143,
2204–2211 (1996).
48.
Wang, L.-F., Ou, C.-C., Striebel, K. a. & Chen, J.-S. Study of Mn Dissolution from
LiMn2O4 Spinel Electrodes Using Rotating Ring-Disk Collection Experiments. J.
Electrochem. Soc. 150, A905 (2003).
49.
Sun, Y. K., Hong, K. J., Prakash, J. & Amine, K. Electrochemical performance of nanosized ZnO-coated LiNi0.5Mn1.5O4 spinel as 5 V materials at elevated temperatures.
Electrochem. commun. 4, 344–348 (2002).
50.
Oh, P. et al. Superior Long-Term Energy Retention and Volumetric Energy Density for
Li-Rich Cathode Materials. Nano Lett. 14, 5965−5972 (2014).
51.
Kang, K. & Ceder, G. Factors that affect Li mobility in layered lithium transition metal
oxides. Phys. Rev. B 74, 1–7 (2006).
52.
Cho, J., Kim, Y. J. & Park, B. Novel LiCoO2 Cathode Material with Al2O3 Coating for a
Li Ion Cell. Chem. Mater. 12, 3788–3791 (2000).
53.
Cho, J. Correlation between AlPO4 nanoparticle coating thickness on LiCoO2 cathode
and thermal stability. Electrochim. Acta 48, 2807–2811 (2003).
54.
Xie, J. et al. Orientation dependence of Li–ion diffusion kinetics in LiCoO2 thin films
prepared by RF magnetron sputtering. Solid State Ionics 179, 362–370 (2008).
55.
Cho, J., Kim, Y. J. & Park, B. LiCoO2 Cathode Material That Does Not Show a Phase
Transition from Hexagonal to Monoclinic Phase. J. Electrochem. Soc. 148, A1110
(2001).
56.
Van der Ven, A. First-Principles Evidence for Stage Ordering in LixCoO2. J.
Electrochem. Soc. 145, 2149 (1998).
57.
Ueda, A. Solid-State Redox Reactions of LiCoO2 (R-3m) for 4 Volt Secondary Lithium
Cells. J. Electrochem. Soc. 141, 2010 (1994).
58.
Cho, Y. & Cho, J. Significant Improvement of LiNi0.8Co0.15Al0.05O2 Cathodes at 60°C
by SiO2 Dry Coating for Li-Ion Batteries. J. Electrochem. Soc. 157, A625 (2010).
120
59.
McCalla, E., Lowartz, C. M., Brown, C. R. & Dahn, J. R. Formation of layered-layered
composites in the Li-Co-Mn oxide pseudoternary system during slow cooling. Chem.
Mater. 25, 912–918 (2013).
60.
Yabuuchi, N., Yoshii, K., Myung, S.-T., Nakai, I. & Komaba, S. Detailed studies of a
high-capacity electrode material for rechargeable batteries, Li2MnO3LiCo1/3Ni1/3Mn1/3O2. J. Am. Chem. Soc. 133, 4404–4419 (2011).
61.
Johnson, C. S. et al. The significance of the Li2MnO3 component in ‘composite’
xLi2MnO3·(1−x)LiMn0.5Ni0.5O2 electrodes. Electrochem. commun. 6, 1085–1091 (2004).
62.
Zheng, J. et al. Mitigating voltage fade in cathode materials by improving the atomic
level uniformity of elemental distribution. Nano Lett. 14, 2628–2635 (2014).
63.
Croy, J. R., Gallagher, K. G., Balasubramanian, M., Long, B. R. & Thackeray, M. M.
Quantifying Hysteresis and Voltage Fade in xLi2MnO3·(1-x)LiMn0.5Ni0.5O2 Electrodes
as a Function of Li2MnO3 Content. J. Electrochem. Soc. 161, A318–A325 (2013).
64.
Mohanty, D. et al. Unraveling the Voltage Fade Mechanism in High-Energy-Density
Lithium-ion Batteries: Origin of the Tetrahedral Cations for Spinel Conversion. Chem.
Mater. 26, 6272−6280 (2014).
65.
Kang, S. H. & Thackeray, M. M. Enhancing the rate capability of high capacity
xLi2MnO3·(1 - x)LiMO2 (M = Mn, Ni, Co) electrodes by Li-Ni-PO4 treatment.
Electrochem. commun. 11, 748–751 (2009).
66.
Rougier, A. Effect of cobalt substitution on cationic distribution in LiNi1 − yCoyO2
electrode materials. Solid State Ionics 90, 83–90 (1996).
67.
Lu, Z., MacNeil, D. D. & Dahn, J. R. Layered Li[NixCo1−2xMnx]O2 Cathode Materials
for Lithium-Ion Batteries. Electrochem. Solid-State Lett. 4, A200 (2001).
68.
Rougier, A., Gravereau, P. & Delmas, C. Optimization of the Composition of the Li1zNi1+zO2 Electrode Materials: Structural, Magnetic, and Electrochemical Studies. J.
Electrochem. Soc. 143, 1168–1175 (1996).
69.
Van der Ven, A. Lithium Diffusion in Layered LixCoO2. Electrochem. Solid-State Lett.
3, 301 (1999).
70.
Lee, J. et al. Unlocking the potential of cation-disordered oxides for rechargeable
lithium batteries. Science 343, 519–522 (2014).
71.
Van Der Ven, A., Bhattacharya, J. & Belak, A. a. Understanding Li diffusion in Liintercalation compounds. Acc. Chem. Res. 46, 1216–1225 (2013).
72.
Urban, A., Lee, J. & Ceder, G. The Configurational Space of Rocksalt-Type Oxides for
High-Capacity Lithium Battery Electrodes. Adv. Energy Mater. 4, 1400478 (2014).
121
73.
Kang, K. et al. Synthesis and Electrochemical Properties of Layered Li0.9Ni0.45Ti0.55O2.
Chem. Mater. 15, 4503−4507 (2003).
74.
Li, J., Li, J., Luo, J., Wang, L. & He, X. Recent Advances in the LiFeO2-based Materials
for Li-ion Batteries. Int. J. Electrochem. Sci. 6, 1550–1561 (2011).
75.
Obrovac, M. N., Mao, O. & Dahn, J. R. Structure and electrochemistry of LiMO2 (M =
Ti , Mn , Fe , Co , Ni) prepared by mechanochemical synthesis. Solid State Ionics 112,
9–19 (1998).
76.
Sakurai, Y., Arai, H. & Yamaki, J. Preparation of electrochemically active α-LiFeO2 at
low temperature. Solid State Ionics 113-115, 29–34 (1998).
77.
Ozawa, K. et al. Structural modifications caused by electrochemical lithium extraction
for two types of layered LiVO2 (R-3m). J. Power Sources 174, 469–472 (2007).
78.
Pralong, V., Gopal, V., Caignaert, V., Duffort, V. & Raveau, B. Lithium-Rich RockSalt-Type Vanadate as Energy Storage Cathode: Li2-xVO3. Chem. Mater. 24, 12−14
(2012).
79.
Yabuuchi, N. et al. High-capacity electrode materials for rechargeable lithium batteries:
Li3NbO4-based system with cation-disordered rocksalt structure. Proc. Natl. Acad. Sci.
112, 7650−7655 (2015).
80.
Park, K. S. et al. LiFeO2-Incorporated Li2MoO3 as a Cathode Additive for Lithium-Ion
Battery Safety. Chem. Mater. 24, 2673–2683 (2012).
81.
Hohenberg, P. & Kohn, W. Inhomogeneous electron gas. Phys. Rev. B 7, 1912–1919
(1973).
82.
Kohn, W. & Sham, L. J. Self-consistent equations including exchange and correlation
effects. Phys. Rev. 140, A1133–A1138 (1965).
83.
Perdew, J. P. J., Burke, K. & Ernzerhof, M. Generalized Gradient Approximation Made
Simple. Phys. Rev. Lett. 77, 3865–3868 (1996).
84.
Blochl, P. E. Projecter augmented-wave method. Phys. Rev. B 50, 17953–17979 (1994).
85.
Kresse, G. From ultrasoft pseudopotentials to the projector augmented-wave method.
Phys. Rev. B 59, 1758–1775 (1999).
86.
Kresse, G. & Furthmüller, J. Efficiency of ab-initio total energy calculations for metals
and semiconductors using a plane-wave basis set. Comput. Mater. Sci. 6, 15–50 (1996).
87.
Kresse, G. Efficient iterative schemes for ab initio total-energy calculations using a
plane-wave basis set. Phys. Rev. B 54, 11169–11186 (1996).
88.
Jain, A. et al. A high-throughput infrastructure for density functional theory calculations.
Comput. Mater. Sci. 50, 2295–2310 (2011).
122
89.
Jonsson, H., Mills, G. & Jacobsen, W. Classical and Quantum Dynamics in Condensed
Phase Simulations. (1998).
90.
Henkelman, G. & Jónsson, H. Improved tangent estimate in the nudged elastic band
method for finding minimum energy paths and saddle points. J. Chem. Phys. 113, 9978–
9985 (2000).
91.
Henkelman, G., Uberuaga, B. P. & Jónsson, H. A climbing image nudged elastic band
method for finding saddle points and minimum energy paths. J. Chem. Phys. 113, 9901–
9904 (2000).
92.
Newman, M. E. J. & Ziff, R. M. Efficient Monte Carlo algorithm and high-precision
results for percolation. Phys. Rev. Lett. 85, 4104–4107 (2000).
93.
Daulton, T. L. & Little, B. J. Determination of chromium valence over the range Cr(0)Cr(VI) by electron energy loss spectroscopy. Ultramicroscopy 106, 561–573 (2006).
94.
Zhang, L. & Noguchi, H. Novel Layered Li-Cr-Ti-O Cathode Materials Related to the
LiCrO2-Li2TiO3 Solid Solution. J. Electrochem. Soc. 150, A601 (2003).
95.
Van der Ven, A., Aydinol, M., Ceder, G., Kresse, G. & Hafner, J. First-principles
investigation of phase stability in LixCoO2. Phys. Rev. B 58, 2975–2987 (1998).
96.
Shirafuji, J. & Field, J. E. CoO2, The End Member of the LiCoO2 Solid Solution.
Electrochem. Soc. 143, 1114–1123 (1996).
97.
Lu, Z. & Dahn, J. R. Understanding the Anomalous Capacity of
Li/Li[NixLi(1/3−2x/3)Mn(2/3−x/3)]O2 Cells Using In Situ X-Ray Diffraction and
Electrochemical Studies. J. Electrochem. Soc. 149, A815 (2002).
98.
Zheng, J. et al. Structural and Chemical Evolution of Li- and Mn-Rich Layered Cathode
Material. Chem. Mater. 27, 1381−1390 (2015).
99.
Yan, P. et al. Evolution of Lattice Structure and Chemical Composition of the Surface
Reconstruction Layer in Li1.2Ni0.2Mn0.6O2 Cathode Material for Lithium Ion Batteries.
Nano Lett. 15, 514–522 (2015).
100. Hewston, T. a. & Chamberland, B. L. A Survey of first-row ternary oxides LiMO2 (M =
Sc-Cu). J. Phys. Chem. Solids 48, 97–108 (1987).
101. Prabaharan, S. R. S., Michael, M. S., Ikuta, H., Uchimoto, Y. & Wakihara, M.
Li2NiTiO4 - A new positive electrode for lithium batteries: Soft-chemistry synthesis and
electrochemical characterization. Solid State Ionics 172, 39–45 (2004).
102. Yang, M. et al. Cation disordered rock salt phase Li2CoTiO4 as a potential cathode
material for Li-ion batteries. J. Mater. Chem. 22, 6200 (2012).
103. Sakuda, A. et al. Rock-salt-type lithium metal sulphides as novel positive-electrode
materials. Sci. Rep. 4, 4883 (2014).
123
104. Chen, R. et al. Disordered Lithium-Rich Oxyfluoride as a Stable Host for Enhanced Li+
Intercalation Storage. Adv. Energy Mater. 5, 1401814 (2015).
105. Ren, S. et al. Improved Voltage and Cycling for Li+ Intercalation in High-Capacity
Disordered Oxyfluoride Cathodes. Adv. Sci. 1500128 (2015).
106. Ravel, B. & Newville, M. ATHENA, ARTEMIS, HEPHAESTUS: Data analysis for Xray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 12, 537–541 (2005).
107. Dudarev, S. L., Savrasov, S. Y., Humphreys, C. J. & Sutton, a. P. Electron-energy-loss
spectra and the structural stability of nickel oxide: An LSDA+U study. Phys. Rev. B 57,
1505–1509 (1998).
108. Jain, A. et al. Formation enthalpies by mixing GGA and GGA + U calculations. Phys.
Rev. B 84, 1–10 (2011).
109. Woodley, S., Battle, P., Gale, J. & Catlow, C. The prediction of inorganic crystal
structures using a genetic algorithm and energy minimisation. Phys. Chem. Chem. Phys.
1, 2535–2542 (1999).
110. Woodley, S. M. & Catlow, R. Crystal structure prediction from first principles. Nat.
Mater. 7, 937–946 (2008).
111. Reed, J., Ceder, G. & Van Der Ven, a. Layered-to-Spinel Phase Transition in LixMnO2.
Electrochem. Solid-State Lett. 4, A78 (2001).
112. Hwang, B. J., Tsai, Y. W., Carlier, D. & Ceder, G. A combined
computational/experimental study on LiNi1/3Co1/3Mn1/3O2. Chem. Mater. 15, 3676–3682
(2003).
113. Wang, L., Maxisch, T. & Ceder, G. Oxidation energies of transition metal oxides within
the GGA+U framework. Phys. Rev. B - Condens. Matter Mater. Phys. 73, 1–6 (2006).
114. Chase, M. NIST-JANAF Thermodynamical Tables. American Chemical Society: New
York 12, (1998).
115. Kang, S., Mo, Y., Ong, S. P. & Ceder, G. Nanoscale stabilization of sodium oxides:
Implications for Na-O2 batteries. Nano Lett. 14, 1016–1020 (2014).
116. Tran, N. et al. Mechanisms Associated with the ‘Plateau’ Observed at High Voltage for
the Overlithiated Li1.12(Ni0.425Mn0.425Co0.15)0.88O2 System. Chem. Mater. 20, 4815–4825
(2008).
117. Zhang, L. et al. Synthesis and electrochemistry of cubic rocksalt Li–Ni–Ti–O
compounds in the phase diagram of LiNiO2–LiTiO2–Li[Li1/3Ti2/3]O2. J. Power Sources
185, 534–541 (2008).
118. Trócoli, R., Cruz-Yusta, M., Morales, J. & Santos-Peña, J. On the limited electroactivity
of Li2NiTiO4 nanoparticles in lithium batteries. Electrochim. Acta 100, 93–100 (2013).
124
119. Twu, N. et al. Designing New Lithium-Excess Cathode Materials from Percolation
Theory: Nanohighways in LixNi2-4x/3Sbx/3O2. Nano Lett. 15, 596–602 (2015).
120. Ong, S. P. et al. Voltage, stability and diffusion barrier differences between sodium-ion
and lithium-ion intercalation materials. Energy Environ. Sci. 4, 3680−3688 (2011).
121. Yu, H.-C. et al. Designing the next generation high capacity battery electrodes. Energy
Environ. Sci. 7, 1760−1768 (2014).
122. Weppner, W. & Huggins, R. A. Determination of the Kinetic Parameters of MixedConducting Electrodes and Application to the System Li3Sb. J. Electrochem. Soc. 124,
1569−1578 (1978).
123. Kim, J. C., Seo, D., Chen, H. & Ceder, G. The effect of antisite disorder and particle
size on Li intercalation kinetics in monoclinic LiMnBO3. Adv. Energy. Mater. 5, 1–17
(2015).
124. Armstrong, A. R. et al. Demonstrating oxygen loss and associated structural
reorganization in the lithium battery cathode Li[Ni0.2Li0.2Mn0.6]O2. J. Am. Chem. Soc.
128, 8694–8 (2006).
125. Sathiya, M. et al. High Performance Li2Ru1-yMnyO3 (0.2 <y<0.8) Cathode Materials for
Rechargeable Lithium-Ion Batteries: Their Understanding. Chem. Mater. 25, 1121–1131
(2013).
126. Hara, D. et al. Charge–discharge reaction mechanism of manganese molybdenum
vanadium oxide as a high capacity anode material for Li secondary battery. J. Mater.
Chem. 13, 897–903 (2003).
127. Stemmer, S. et al. Oxidation states of titanium in bulk barium titanates and in (100)
fiber-textured (BaxSr1-x)Ti1+yO3+z thin films. Appl. Phys. Lett. 79, 3149–3151 (2001).
128. Ates, M. N., Mukerjee, S. & Abraham, K. M. A high rate Li-rich layered MNC cathode
material for lithium-ion batteries. RSC Adv. 5, 27375–27386 (2015).
129. Koga, H. et al. Different oxygen redox participation for bulk and surface: A possible
global explanation for the cycling mechanism of Li1.20Mn0.54Co0.13Ni0.13O2. J. Power
Sources 236, 250–258 (2013).
130. Koga, H. et al. Operando X-ray Absorption Study of the Redox Processes Involved
upon Cycling of the Li-Rich Layered Oxide Li1.20Mn0.54Co0.13Ni0.13O2 in Li Ion Batteries.
J. Phys. Chem. C 118, 5700–5709 (2014).
131. Oh, P., Ko, M., Myeong, S., Kim, Y. & Cho, J. A Novel Surface Treatment Method and
New Insight into Discharge Voltage Deterioration for High-Performance 0.4Li2MnO30.6LiNi1/3Co1/3Mn1/3O2 Cathode Materials. Adv. Energy Mater. 4, 1400631 (2014).
125
132. Kim, H., Kim, M. G., Jeong, H. Y., Nam, H. & Cho, J. A New Coating Method for
Alleviating Surface Degradation of LiNi0.6Co0.2Mn0.2O2 Cathode Material: Nanoscale
Surface Treatment of Primary Particles. Nano Lett. 15, 2111−2119 (2015).
126
127