pubs.acs.org/cmPublished on Web 08/28/2009r 2009 American Chemical Society
Chem. Mater. 2010, 22, 587–603 587
DOI:10.1021/cm901452z
Challenges for Rechargeable Li Batteries†
John B. Goodenough* and Youngsik Kim
Texas Materials Institute, University of Texas at Austin, Austin, Texas 78712
Received May 27, 2009. Revised Manuscript Received July 9, 2009
The challenges for further development of Li rechargeable batteries for electric vehicles are
reviewed. Most important is safety, which requires development of a nonflammable electrolyte with
either a larger window between its lowest unoccupied molecular orbital (LUMO) and highest
occupied molecular orbital (HOMO) or a constituent (or additive) that can develop rapidly a solid/
electrolyte-interface (SEI) layer to prevent plating of Li on a carbon anode during a fast charge of the
battery. A high Liþ-ion conductivity (σLi > 10
-4 S/cm) in the electrolyte and across the electrode/
electrolyte interface is needed for a power battery. Important also is an increase in the density of the
stored energy, which is the product of the voltage and capacity of reversible Li insertion/extraction
into/from the electrodes. It will be difficult to design a better anode than carbon, but carbon requires
formation of an SEI layer, which involves an irreversible capacity loss. The design of a cathode
composed of environmentally benign, low-cost materials that has its electrochemical potential μC
well-matched to the HOMOof the electrolyte and allows access to two Li atoms per transition-metal
cation would increase the energy density, but it is a daunting challenge. Two redox couples can be
accessed where the cation redox couples are “pinned” at the top of the O 2p bands, but to take
advantage of this possibility, it must be realized in a framework structure that can accept more than
one Li atom per transition-metal cation. Moreover, such a situation represents an intrinsic voltage
limit of the cathode, and matching this limit to the HOMO of the electrolyte requires the ability to
tune the intrinsic voltage limit. Finally, the chemical compatibility in the battery must allow a long
service life.
Introduction
It is now almost universally recognized that gaseous
emissions from the burning of fossil fuels and biomass are
not only polluting the air of large, modern cities but are
also creating a global warming with alarming conse-
quences. Moreover, a dependence on foreign oil and/or
gas creates national vulnerabilities that endanger social
stability. These concerns are concentrating attention once
again on national initiatives to reevaluate utilization of
alternative energy sources and replacement of the internal
combustion engine with a wireless electric motor.
Solar radiation, wind, and waves represent energy
sources that are variable in time and diffuse in space.1
These sources require energy storage. Nuclear reactors
provide a constant energy source with associated pro-
blems of radioactive waste disposal.Geothermal energy is
restricted in location. These energy sources also benefit
from electrical energy storage. The energy carriers are the
electricity grid, electromagnetic waves, and chemical en-
ergy. The most convenient form of energy storage is
portable chemical energy, which is the reason for our
addiction to fossil fuels for heat, propulsion, lighting, and
communication. The battery provides the portability of
stored chemical energy with the ability to deliver this
energy as electrical energy with a high conversion effi-
ciency and no gaseous exhaust. Moreover, the alternative
energy sources are preferably converted to d.c. electrical
energy well-matched to storage as chemical energy in a
battery. Whereas alternative energy sources are station-
ary, which allows other means of energy storage to be
competitive with a battery, electric vehicles require the
portable stored energy of a fuel fed to a fuel cell or of a
battery. Therefore, of particular interest is a low-cost,
safe, rechargeable (secondary) battery of high voltage,
capacity, and rate capability.
The higher stored volume and gravimetric energy
density of a Li battery has enabled realization of the
cellular telephone and lap-top computer. However, cost,
safety, stored energy density, charge/discharge rates, and
service life are issues that continue to plague the develop-
ment of the Li battery for the potential mass market of
electric vehicles to alleviate distributed CO2 emissions
and noise pollution.2
A battery consists of a group of interconnected electro-
chemical cells. Here, we focus on batteries for electric
vehicles where cost, gravimetric energy density, and the
performance uniformity of individual cells in a large,
multicell battery are of more concern than the volume
energy density considered critical for hand-held appli-
ances. Moreover, we consider only the choice of active
materials in the individual cells of a secondary battery,
†Accepted as part of the 2010 “Materials Chemistry of Energy Conversion
Special Issue”.
*Author to whom correspondence should be directed. E-mail:
jgoodenough@mail.utexas.edu.

588 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
viz. the anode (negative electrode), the cathode (positive
electrode), and the electrolyte between the electrodes.
Preliminary Considerations
Figure 1 is a schematic of the relative electron energies
in the electrodes and the electrolyte of a thermodynami-
cally stable battery cell having an aqueous electrolyte.
The anode is the reductant, the cathode is the oxidant,
and the energy separation Eg of the lowest unoccupied
molecular orbital (LUMO) and the highest occupied
molecular orbital (HOMO) of the electrolyte is the “win-
dow” of the electrolyte. The two electrodes are electronic
conductors with anode and cathode electrochemical
potentials μA and μC (their Fermi energies εF). An anode
with a μA above the LUMO will reduce the electrolyte
unless a passivation layer creates a barrier to electron
transfer from the anode to the electrolyte LUMO; a
cathode with a μC below the HOMO will oxidize the
electrolyte unless a passivation layer blocks electron
transfer from the electrolyte HOMO to the cathode.
Therefore, thermodynamic stability requires locating
the electrode electrochemical potentials μA and μC within
the window of the electrolyte, which constrains the open-
circuit voltage Voc of a battery cell to
eVoc ¼ μA-μCeEg ð1Þ
where e is the magnitude of the electron charge. A
passivating solid/electrolyte-interface (SEI) layer at the
electrode/electrolyte boundary can give a kinetic stability
to a larger Voc provided that eVoc - Eg is not too large.
On discharge, electrons leave the anode via an external
circuit where they do useful work before entering the
cathode. To retain charge neutrality in the electrodes,
cations are released from the anode to the electrolyte and
the working cation of the electrolyte, the Hþ ion in an
aqueous electrolyte, carries positive charge to the cathode
to provide charge neutrality in the cathode. The process is
reversed on charge in a rechargeable (secondary) battery.
The energy density of a battery cell is ΛVoc; Λ is the
capacity of reversible charge transfer per unit weight (Ah/
g) between the anode and cathode. Λ decreases with the
rate of charge or discharge, i.e. the magnitude of the
electronic current in the external circuit, which must be
matched by the internal ionic current within the battery.
Since the ionic current density of the electrolyte and
electrodes, including the rate of ion transfer across the
electrode/electrolyte interface, is much smaller than
the electronic current density, the electrodes and electro-
lyte have a large surface area and a small thickness.
Nevertheless, at high current densities, the ionic motion
within an electrode and/or across an electrode/electrolyte
interface is too slow for the charge distribution to reach
equilibrium, which is why the reversible capacity de-
creases with increasing current density in the battery
and why this capacity loss is recovered on reducing the
rate of charge and/or discharge.
The high Hþ-ion conductivity required of an aqueous
electrolyte over the practical ambient-temperature range
is only found in liquid or immobilized-liquid water, and
an Eg ≈ 1.3 eV for an aqueous electrolyte limits Voc. In
order to obtain a cell with a higher Voc and therefore a
higher energy density ΛVoc, it is necessary to turn to a
nonaqueous electrolyte with a larger Eg. This observa-
tion, in turn, has led to the Liþ-ion battery since lithium
salts are soluble in some nonaqueous liquids and poly-
mers. However, in this case, the HOMO of the salt as well
as that of the solvent may determine the limiting μC of the
cathode.
Once the window of the Liþ-ion electrolyte has been
determined, it is necessary to design electrodes of high
capacity that have their μA and μCmatched to the LUMO
andHOMOof the electrolyte. Elemental Li0 would be the
ideal anode, but the εF = μA of Li
0 lies above the LUMO
of practical, known nonaqueous electrolytes. Therefore,
use of Li0 as an anode is only possible because a passivat-
ing SEI layer is formed. The SEI layer allows use of Li0 as
an anode in half-cells used to obtain the μA or μC of a
practical electrode relative to the Liþ/Li0 energy level; but
on repeated charge/discharge cycles, breaking of the SEI
layer in selected areas results in the formation of dendrites
that can grow across the electrolyte to short-circuit a cell
of the battery with dangerous consequences. Therefore,
we must design either (1) an anode with a μA matched to
the LUMOof the electrolyte aswell as a cathodewith a μC
matched to the HOMO of the electrolyte or (2) a stable
passivating SEI layer that self-heals rapidly when broken
by the changes in electrode volume that occur in a charge/
discharge cycle; the SEI layer must also permit a fast Liþ-
ion transfer between the electrode and the electrolyte
without blocking electron transfer between the active
particle and the current collector.
In summary, the formidable challenges for the devel-
oper of a rechargeable Li battery for the potential
mass market of electric vehicles are three-fold: to identify
Figure 1. Schematic open-circuit energy diagram of an aqueous electro-
lyte. ΦA and ΦC are the anode and cathode work functions. Eg is the
window of the electrolyte for thermodynamic stability. A μA > LUMO
and/or a μC <HOMO requires a kinetic stability by the formation of an
SEI layer.

Review Chem. Mater., Vol. 22, No. 3, 2010 589
low-cost, environmentally benign materials for the
three active components of a battery cell, viz. (1) a
nonaqueous electrolyte of high Liþ-ion conductivity
(σLi>10
-3 S/cm) over the practical ambient-temperature
range -40 < T < 60
C that has a window allowing
a thermodynamically stable Voc g 4 V and (2) an anode
and (3) a cathode with their μA and μC values well-
matched to the window of the electrolyte as well as each
allowing a fast charge/discharge cycle of large reversible
capacity.
Electrolytes
In addition to a large electrolyte window Eg, the
electrolyte must satisfy several additional requirements
such as:
1) Retention of the electrode/electrolyte interface
during cycling when the electrode particles are
changing their volume.
2) A Liþ-ion conductivity σLi>10
-4 S/cm over the
temperature range of battery operation.
3) An electronic conductivity σe<10
-10 S/cm.
4) A transference number σLi/σtotal ≈ 1, where σtotal
includes conductivities by other ions in the elec-
trolyte as well as σLi þ σe.
5) Chemical stability over ambient temperature ranges
and temperatures in the battery under high power.
6) Chemical stability with respect to the electrodes,
including the ability to form rapidly a passiva-
ting solid/electrolyte-interface (SEI) layer where
kinetic stability is required because the electrode
potential lies outside the electrolyte window.
7) Safe materials, i.e., preferably nonflammable and
nonexplosive if short-circuited.
8) Low toxicity and low cost.
Meeting all these requirements proves to be a
formidable challenge.
Types of Electrolytes. In general, the electrolyte is
specifically designed for a particular battery application.
Table 1 shows several different materials that have been
used as electrolytes for Li batteries.
Organic Liquid Electrolytes. Carbonates are orga-
nic liquids that are reasonably good solvents for Li
salts.3,4 They have an oxidation potential (HOMO) at
ca. 4.7 V3,5,6 and a reduction potential (LUMO) near 1.0
V.7 (All voltages in this paper are referred to the Liþ/Li0
potential.) Moreover, they have a relatively low viscos-
ity, which results in a low activation energy for Liþ-ion
diffusion. Therefore, the most commonly used electro-
lytes are carbonates or carbonate blends consisting of
one or more of the following: propylene carbonate (PC),
ethylene carbonate (EC), diethyl carbonate (DEC),
dimethyl carbonate (DMC), or ethylmethyl carbonate
(EMC). In order to be able to use carbon as the anode,
which has its electrochemical potential above (at a
lower voltage versus Liþ/Li0) that of the LUMO of a
carbonate, the solvents include, in most cases, ethylene
carbonate (EC) because the EC provides a passivating
Table 1. Nonaqueous Electrolytes for Li-Ion Batteries
Electrolytes Example of classical electrolytes
Ionic conductivity
( 10-3 s/cm)
at room temp
Electrochemical
window (V) vs Liþ/Li0
Remark
Reduction Oxidation
Liquid organic 1M LiPF6 in EC:DEC (1:1) 7
3 1.37 4.56 Flammable
1M LiPF6 in EC:DMC (1:1) 10
3 1.37 >5.03
Ionic liquids 1M LiTFSI in EMI-TFSI 2.015 1.015 5.315 Non-flammable
1M LiBF4 in EMI-BF4 8.0
15 0.916 5.316
Polymer LiTFSI-P(EO/MEEGE) 0.124 <0.024 4.724 Flammable
LiClO4-PEO8 þ 10 wt % TiO2 0.02
26 <0.026 5.026
Inorganic solid Li4-xGe1-xPxS4 (x = 0.75) 2.2
28 <0.028 >5.028 Non-flammable
0.05Li4SiO4 þ 0.57Li2S þ 0.38SiS2 1.0
30 <0.030 >8.030
Inorganic liquid LiAlCl4 þ SO2 70
20 - 4.420 Non-flammable
Liquid organic þ
Polymer
0.04LiPF6 þ 0.2EC þ 0.62DMC þ
0.14PAN
4.238 - 4.438 Flammable
LiClO4 þ EC þ PC þ PVdF 3.0
39 - 5.039
Ionic liquid þ
Polymer
1M LiTFSI þ P13TFSI þ
PVdF-HFP
0.1843 <0.043 5.843 Less flammable
Ionic liquid þ Polymer
þ Liquid organic
56 wt % LiTFSI-Py24TFSI þ
30 wt % PVdF-HFP þ
14 wt % EC/PC
0.8144 1.544 4.2 44 Less flammable
Polymer
þ Inorganic solid
2 vol % LiClO4-TEC-19 þ 98 vol%
95 (0.6Li2S þ 0.4Li2S) þ 5Li4SiO4
0.0346 <0.046 >4.546 Non-flammable
Ionic liquid þ Liquid organic19 - - - Non-flammable

590 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
solid/electrolyte-interface (SEI) layer on the surface of a
carbon anode that protects the electrolytes from further
decomposition after SEI formation.8,9 However, carbo-
nate-based solvents are highly flammable with flash
points below 30
C.10 In addition, the preferred salt,
LiPF6, can undergo an autocatalytic decomposition into
LiF and PF5; the PF5 reacts irreversibly with any water
present (PF5 þ H2O=PF3O þ 2HF) and, above 60
C,
with a carbonate electrolyte.11 These reactions degrade
the battery and lead to safety hazards. However, addi-
tives used to lower the operating temperature have been
shown to prevent the autocatalytic decomposition of
LiPF6 salt.
12
Ionic Liquids. Room-temperature ionic liquids
(RTILs)13-17 have been recently considered as alternative
electrolytes for Li-ion batteries because they offer several
advantages over carbonate-based electrolytes: a high
oxidation potential (∼5.3 V vs Liþ/Li0), nonflammabil-
ity, a low vapor pressure, better thermal stability, low
toxicity, high boiling points, and a high Li-salt solubility.
Unfortunately, they have a higher viscosity, which re-
duces their Liþ-ion conductivity. Ionic liquids based on
imidazolium-based cations would appear to be the most
appropriate candidates for Li batteries due to their lower
viscosity and a high Li-salt solubility at room tempera-
ture. However, these ionic liquids have poor stability at
voltages below 1.1 V,18 so that additives such as ECorVC
must be added to introduce a stable SEI layer on a carbon
anode. An alternative approach is to increase σLi by
adding a liquid carbonate to an ionic liquid, but at a
concentration that retains the nonflammability of the
ionic liquid.19 With this strategy, it is also possible to
increase the oxidation voltage (lower HOMO energy) of
the hybrid electrolyte from that of the carbonate. In spite
of extensive research, no RTILs have yet been introduced
into large power batteries.
Inorganic Liquid Electrolytes. The inorganic liquid
electrolyte based on LiAlCl4 and SO2 proposed by Stas-
sen and Hambitzer20,21 has a good room-temperature
σLi=7  10
-2 S/cm and is nonflammable, but its electro-
lyte window appears to be too small to be competitive.
Solid Polymer Electrolytes. A solid electrolyte can act
as the separator of the electrodes, and a solid polymer
electrolyte can also retain contact over an electrode/
electrolyte interface during modest changes of the elec-
trode volume with the state of charge of the battery.
Polyethylene oxides (PEOs) containing a lithium salt
(LiPF6 or LiAsF6)
22-24 are low-cost, nontoxic, Liþ-ion
polymer electrolytes with good chemical stability, but the
Liþ-ion conductivity, σLi<10
-5 S/cm at room tempera-
ture, is too low for a power-battery system. The introduc-
tion of oxide particles (e.g., Al2O3, TiO2, SiO2, or
ZrO2)
25-27 creates a more amorphous polymer matrix
by inhibiting chain crystallization and attracting Liþ
from its salt. The result is an enhanced σLi and Li-ion
transference number, but σLi is still not comparable to
that of the carbonate electrolytes.
Inorganic Solid Electrolytes. Inorganic solid Liþ-ion
conducting materials having a σLi>10
-4 S/cm28-31 have
been considered for Li-based electrolytes, as has been
extensively reviewed,32 because they have a wide electro-
chemical window and additionally meet the electrolyte
requirements from 2 to 7, not 1. For these reasons,
laboratory-size all-solid-state Li-ion batteries have been
investigated.33-35 However, the first of the additional
electrolyte requirements has excluded inorganic solid
Liþ-ion electrolytes from consideration for large-scale
batteries having solid electrodes. They have only been
used in thin-film battery applications.36
Hybrid Electrolyte System. Hybrid electrolytes are
blends of organic liquid electrolytes, ionic liquids, poly-
mer electrolytes, and/or inorganic solid electrolytes:
Polymer þ organic liquid (polymer gel)37-40
Ionic liquid þ polymer electrolyte (ionic liquid poly-
mer gel)41-45
Ionic liquid þ polymer electrolyte þ liquid organic
electrolyte43,44
Ionic liquid þ liquid organic electrolyte19
Polymer electrolyte þ inorganic solid electrolyte46-48
The mixtures of two or more electrolytes are investi-
gated in attempts to exploit the advantages of each
constituent, but the disadvantages of each also appear.
For example, the ionic conductivity is increased in the
polymer gel electrolytes, but they are still flammable and
have the irreversible capacity loss below 1 V associated
with formation of a passivation layer.40
Electrode-Electrolyte Compatibility. Although ther-
modynamic stability of the electrolyte vis
a vis the elec-
trodes is possible where the μA and μC of the electrodes lie
within the window of the electrolyte, nevertheless chemi-
cal reactions between the electrode and the electrolyte
may occur. For example, the reversible electrochemical
intercalation of Li into LixVS2 was originally frustrated
by use of the electrolyte LiClO4 in PC,
49 but the electro-
lyte LiPF6 in EC/DEC allows full electrochemical cycling
between LiVS2 and VS2.
50
The cathode spinel Li1-x[Mn2]O4 provides another
type of electrode-electrolyte reaction; an electrode sur-
face disproportionation reaction 2Mn3þ=Mn2þþMn4þ
results in dissolution of the Mn2þ from the electrode into
the electrolyte.51 This reaction, unless suppressed, gives
an irreversible capacity loss of the cathode and migration
of the Mn2þ across the electrolyte to the anode during
charge to block Liþ-ion insertion into the anode. The
result is an intolerable limitation of the service life of
the cell.
In addition to chemical stability vis
a vis the electrodes
and higher temperatures, the electrolyte should not be
decomposed by an anode μA at a higher energy than the
electrolyte LUMOor a cathode μC at a lower energy than
the HOMO. However, if μA or μC lie outside the window
of the electrolyte, kinetic stability may be achieved by
formation of a passivating SEI layer on the surface of the
electrode, but at the expense of the loss in capacity to form
the layer. Moreover, during a fast charge, the concentra-
tion of Liþ ions may build up on the surface of the SEI
layer, and where a change in volume of the electrode

Review Chem. Mater., Vol. 22, No. 3, 2010 591
breaks the SEI layer, Li0 may be plated out before the
break is healed. Li plating can result in dendrites that
grow across the electrolyte. This problem creates a safety
issue that has haunted the use of a carbon anode in large-
scale power batteries. These problems need to be mana-
ged if safety standards are to be met with any anode,
including carbon, that has its μA above the LUMO of the
electrolyte.
A cathode μC at a lower energy than the electrolyte
HOMO must be distinguished from an intrinsic voltage
limit of the cathode, as is discussed below. As with
anodes, passivating layers on cathodes are best formed
in situ so that electronic contact with the cathode current
collector is not broken. Preliminary work52-54 on pas-
sivating SEI layers on oxide cathodes has found them
to be unstable. This field has yet to be adequately
researched.
Electrodes
The design of an electrode involves tailoring of the μA
of an anode or μC of a cathode to the LUMO or HOMO
of the Liþ-ion electrolyte to be used; the electrode must
also be chemically stable in the electrolyte. To date,
practical electrodes have all had host structures into/from
which guest Li atoms can be inserted/extracted reversibly.
Factors Determining μA and μC. The energy of a given
μA or μC may correspond to the Fermi energy in an
itinerant-electron band, as is the case for carbon, or the
energy of a redox couple of a transition-metal cation.
Tailoring of the energy of a redox couple depends not
only on the formal valence state of the cation, but also on
the covalent component of its nearest-neighbor bonding,
which is influenced by the placement and character of any
counter cations and by the Madelung energy of the ionic
component of the bonding, which is influenced by the
structure. In addition, the position of a redox couple
relative to the bottom of a broad conduction band or to
the top of an anion p band may determine the intrinsic
voltage limit versus Liþ/Li0 of a given electrode. This
problem arises for a μC where the active redox couple is
“pinned” at the top of the anion p bands. Pinning of redox
couples and the intrinsic voltage limit are concepts de-
scribed below. However, we first demonstrate in Figure 2
the range of voltages that are exhibited by host structures
Figure 2. (a) Voltage profiles versus Liþ/Li0 of the discharge curves of LixC6, LixTiS2 and Lix[Ti2]S4, LixCoO2, and LixCoPO4. (b) Schematic of their
corresponding energy vs density of states showing the relative positions of the Fermi energy in an itinerant electron band for LixC6, the Ti
4þ/Ti3þ redox
couple for LixTiS2 and Lix[Ti2]S4, the Co
4þ/Co3þ redox couple for LixCoO2, and the Co
3þ/Co2þ redox couple for LixCoPO4.

592 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
into/from which Liþ ions have been inserted reversibly.
Carbon, LixTiS2, and LixCoO2 are all layered com-
pounds, the [Ti2]S4 spinel host of Lix[Ti2]S4 is strongly
bonded in 3D, and LixCoPO4 shows the influence of the
countercation of the (PO4)
3- polyanion on the Co3þ/
Co2þ couple relative to the Co4þ/Co3þ couple in the
layered LixCoO2 (charges on ions represent formal
valence states, not actual charges). The upper voltage
limits in the sulfides are much lower than those in the
oxides. In Figure 3, we also show how the Mn4þ/Mn3þ
couple of Lix[Mn2]O2 in a spinel framework is shifted by
more than 1 eV where the Liþ ions change their position
from octahedral to tetrahedral sites as 2 > x > 1
decreases to 1 > x > 0. The influence of structure is
exemplified by the comparison in Figure 4 of the vol-
tages from the Fe3þ/Fe2þ couple in the olivine LixFe-
PO4, the NASICON structure of Li3þxFe2(PO4)3, and
some diphosphates.55
Host Structures. The sulfur atoms of TiS2 form a close-
packed hexagonal arraywith Ti occupying alternate (001)
planes of octahedral sites. The TiS6/3 sheets of edge-
shared octahedra are held together by van der Waals
forces. Steele56 originally suggested that intercalation of
Li into the empty octahedral sites between the TiS6/3
sheets would be reversible, which made TiS2 a potential
cathode for a rechargeable Li battery.Whittingham57was
the first to demonstrate fast, reversible Li insertion into
TiS2 over the solid-solution range 0 e x e 1 of LixTiS2.
However, attempts to make a TiS2/Li
0 battery failed
because dendrite formation on the Li0 anode caused
explosive failure. Nevertheless, these early experiments
demonstrated that compounds into which Li can be
inserted/extracted reversibly are candidate electrodes
for the rechargeable Li battery.
The horizontal, dashed lines of Figure 5 are the energies
relative to the Liþ/Li0 potential of the LUMO and
HOMO of the EC/DEC solvent containing the more
benign LiPF6 as the Li
þ-ion salt. This figure shows that
the energy μC of the Ti
4þ/Ti3þ redox couple of LixTiS2 is
not well-matched to the HOMO of this electrolyte. Rea-
lization that TiS2 approaches the voltage limit versus Li
þ/
Li0 of a layered sulfide suggested exploration of Li
insertion into layered oxides.58 However, layered oxides
are only found where a transition-metal cation forms an
MdO bond as with the vanadyl VdO and molybdyl
ModO cations of V2O5 and MoO3.
59,60 On the other
hand, LiMO2 oxides forming an ordered rock-salt struc-
ture with Li and transition-metal M atoms on alternate
(111) octahedral-site planes invited investigation of
reversible Li extraction.58,61 Removal of Li from ordered
LiMO2 allows operating on anM
4þ/M3þ couple of lower
energy than the Ti4þ/Ti3þ couple of TiS2. However,
removal of Li leaves a metastable compound, and M
cations stable in tetrahedral sites either move into the
partially occupied Li layer or transform the structure to
spinel on removal of half of the Li. Moreover, good order
of the Li and M atoms in the initial LiMO2 is required.
Nevertheless, removal of half of the Li from well-ordered
LiCoO2 at a μC≈ 4.0 V versus Li
þ/Li0 for the Co4þ/Co3þ
couple proved stable. This couple has a goodmatch to the
HOMO of the LiPF6 in EC/DEC electrolyte, but only
one Li for two cobalt represents a reduced capacity; and
Co is too expensive and toxic for a large-battery mass
market.
Graphite has a layered structure that seemed to offer an
intercalation anode to replace Li0, but early attempts to
use graphite were frustrated by reduction of the electro-
lyte on Li insertion.62 As is shown in Figure 5, the μA of
carbon lies well above the LUMO of a carbonate electro-
lyte, which is why identification of a Li intercalation
compound is not a sufficient condition for a viable
electrode. On the other hand, incorporation of ethylene
carbonate (EC) into the carbonate electrolyte promotes
formation of an SEI layer on the carbon that provides a
kinetic stability.8 An irreversible capacity loss on the
Figure 3. Voltage profile versus Liþ/Li0 of the spinel LixMn2O4
(2g xg 0). The Liþ ions are shifted cooperatively from the tetrahedral
to the octahedral sites as x increases though x = 1. Adapted from
ref 67.
Figure 4. Positions of the Fe3þ/Fe2þ redox couples relative to the Fermi
energyof lithium in different phosphates.Reprintedwith permission from
ref 55. Copyright 1997 The University of Texas at Austin.

Review Chem. Mater., Vol. 22, No. 3, 2010 593
initial charge of the carbon anode is associated with the
formation of a thin, amorphous SEI layer on the
carbon that stabilizes reversible Li insertion/extraction
on subsequent charge/discharge cycles, see Figure 6, with
a reversible capacity of 370 mA h/g. Disordered carbon
rather than graphitic carbon provides a better capacity.63
With a passivated carbon anode and LiCoO2 as the
cathode, members of the Sony corporation launched the
hand-held wireless revolution with their introduction of
the wireless telephone.64
The next step was to recognize that framework struc-
tures offer strong 3D bonding as well as interstitial space
for the insertion of Liþ ions. For example, the A[B2]X4
spinels contain B cations in octahedral sites andA cations
in tetrahedral sites of a close-packed-cubic X-atom array.
The B cations are ordered to give a 3D-bonded [B2]X4
framework in which the interstitial space is intercon-
nected by edge-sharing octahedral sites that share faces
with the tetrahedral A sites; see Figure 7. Murphy and
colleagues 65 removed Cu from Cu[Ti2]S4 and then in-
serted Li into the [Ti2]S4 spinel framework. In this sulfide,
Liþ ions initially enter the octahedral sites of the inter-
stitial space rather than the tetrahedral sites, so the
voltage versus x profile of Lix[Ti2]S4, 0 e x e 1, was
essentially identical to that of the layered LixTiS2.
57
Independent work at Oxford66 showed that Li can be
inserted into the oxospinels; but in oxides the Li species
occupy the tetrahedral sites in Li1-x[B2]O4. On insertion
of Li into Li[B2]O4, Coulomb interactions between the
Liþ ions displace all the Liþ ions to the octahedral sites. In
the spinel Li[Mn2]O4, the high-spinMn
3þ ions are Jahn-
Teller ions, and cooperative orbital ordering for a ratio
Mn3þ/Mn4þ > 0.5 distorts the cubic structure to tetra-
gonal to give a coexistence of two phases rather than a
solid solution and therefore a flat Voc ≈ 3.0 V for Li1þx-
[Mn2]O4. Subsequently, Thackeray et al.
67 showed that on
removal of the Li from the tetrahedral sites of Li1-x-
[Mn
2
]O4, the Voc versus x profile was at 4.0 V versus Li
þ/
Li0 for the sameMn4þ/Mn3þ redox couple. These observa-
tions, summarized in Figure 3, showed that shifting the Liþ
ions from octahedral to tetrahedral sites produces a 1 eV
step in the Mn4þ/Mn3þ redox couple as a result of the
inductive effect of the Liþ ions. However, it also shows
that use of the oxospinels limits the operative capacity
to one Li per two B-site cations as in layered LiCoO2.
Although Mn is cheaper and environmentally more
Figure 5. Voltageversus capacityof several electrodematerials relative to
the window of the electrolyte 1 M LiPF6 in EC/DEC (1:1).
Figure 6. (a) Voltage curves of graphite tested in 1MLiClO4 in PC and 1
MLiAsF6 inPC:EC (1:1) electrolytes. The electrolyte is reducedatV≈ 0.7
V in 1MLiClO4 inPC.AnSEI layer is formed in theEC-based electrolyte
between 0.8 and0.4VversusLiþ/Li0, whichallows further intercalationof
Liþ ions after an initial capacity loss. Adapted from refs 62 and 8. (b)
Schematic presentation of the formation of the SEI layer by decomposi-
tion of EC-based electrolytes. Adapted from ref 9.
Figure 7. Two quadrants of the cubic spinel A[B2]X4 showing
the occupied tetrahedral sites (8a), occupied octahedral sites (16d),
and unoccupied octahedral sites (16c). The Li species of Li1-x[B2]O4
occupy 8a tetrahedral sites, and those of Li1þx[B2]O4 occupy only
unoccupied octahedral sites (16c). The Li species of Lix[Ti2]S4 occupy
only unoccupied octahedral sites (16c) for all x of 0 e x e 2. Adapted
from ref 77.

594 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
benign than Co, it has proven necessary to substitute some
Li and Ni forMn to suppress Li order at Li0.5[Mn2]O4 and
dissolution of Mn to the electrolyte on repeated charge/
discharge cycling; the resulting further loss of capacity is
only partially regained by the ability to replace some
oxygen with fluorine.68,69
Another 3D framework is the M2(XO4)3 structure of
hexagonal Fe2(SO4)3, which consists of clusters of two
MO6 octahedra bridged by three corner-sharing (XO4)
tetrahedra; the octahedra of these clusters share corners
with the tetrahedra of neighboring clusters to create an
open 3D host structure capable of accepting up to 5 Li
atoms per formula unit into its interstitial space; see
Figure 8. This framework is referred to as the NASICON
structure since Na1þ3xZr2(P1-xSixO4)3 was shown
70 to
support fast Naþ-ion conduction, i.e. to be a NA-ion
Superior Ionic CONductor. Substitution of an (XO4)
polyanion for oxygen opens the host framework. More-
over, the observation71 of a 0.6 V increase in theVoc from
Li3þxFe2(MoO4)3 to Li3þxFe2(SO4)3, each operating on
the Fe3þ/Fe2þ couple, demonstrated that significant tun-
ing of the energy of a redox couple can also be achieved
through the inductive effect by changing the counter-
cation in the polyanion. With this framework, it was
possible to determine the relative energies of several redox
couples in an oxide, Figure 9, and how these shift together
on replacing (PO4) with (SO4).
72,73
This exploration led to identification74 of the olivine
framework, Figure 10, of FePO4 in which insertion of Li
into its 1D channels gives a flat Voc = 3.45 V versus Li
þ/
Li0 for LixFePO4, 0 e x e 1, as a result of a small
Figure 9. (a) Positions of the Fe3þ/Fe2þ redox couples relative to the Fermi energy of lithium in the NASICON structure with different polyanion
countercations. Adapted from ref 55. (b) Positions of some Mn/Mnþ1 redox couples in LixM2(PO4)3. Adapted from ref 73.
Figure 8. NASICON framework of LixM2(XO4)3 that is built withMO6
octahedra linked by corners to XO4 tetrahedra and vice versa. Adapted
from ref 70.
Figure 10. Olivine structure of LiFePO4 showing Li in 1D channels.
Adapted from ref 74.

Review Chem. Mater., Vol. 22, No. 3, 2010 595
displacive structural change of the framework between
LiFePO4 and FePO4. Despite the two-phase character for
0< x<1,which results in a poor electronic conductivity,
and the 1D channels for the Li motion, small particles of
carbon-coated C-LixFePO4 result in safe cathodes with
reversible capacities that do not fade significantly on
cycling thousands of times. However, the Voc is not
optimal for the LiPF6 in the EC/DEC electrolyte. Never-
theless, a cell voltage Voc ≈ 3 V can be obtained with
LiFePO4 as a cathode and a carbon anode, which is
excellent for many applications. However, to ensure
safety for large-scale power applications, it would be
preferable to have an anode with a μA ≈ 1.3 V versus
Liþ/Li0 if LiPF6 in EC/DEC is used as the electrolyte.
Such an anode with a LiFePO4 cathode would, however,
only give a cellVoc≈ 2 V, whichmight not be competitive
with a nickel/metal-hydride battery having an aqueous
electrolyte.
Johnston75 had shown that the spinel Li[Ti2]O4 is a
superconductor and the system Li[LixTi2-x]O4, 0 e x e
1/3, had been well-characterized,76,77 but the identifica-
tion of Li1þx[Li0.33Ti1.67]O4 as a potential anode with a
flat Voc = 1.5 V versus Li
þ/Li0 was first made by Ferg et
al.78 Although Li4Ti5O12 represents a thermodynamically
stable anode having no passivation layer, it has only a
modest capacity and its use as an anode requires identi-
fication of a cathode with a better match to the HOMO
of the electrolyte than LiFePO4.
Intrinsic Voltage Limits. Pinning of a redox couple at
the top of an anion p band provides an intrinsic voltage
limit for a cathode.79 Pinning occurswhere, as is illustrated
inFigure 11, the energyof a redox couple crosses the topof
the anion p bands. At this crossover, the electronic states
of a dn redox couple change from primarily cation d, i.e.
(d þ p)n, to primarily anion p, i.e. (p þ d)n character, and
where the couple has a primarily anion p character, the
cation will appear to be in a lower valence state, dnþ1 after
oxidation of the couple. Nevertheless, the antibonding
(pþ d)n orbitals retain a d orbital symmetry and behave as
a redox couple. However, as the percentage of anion
p character increases as the redox couple falls further
below the top of the anion p bands or with increasing
oxidation beyond a critical value, the antibonding char-
acter of the states at the top of the anion p bands fades and
holes occupy bonding anion p states; for larger concentra-
tions of purely anion p holes, the holes become trapped in
dianion antibonding states, e.g. (S2)
2- or (O2)
2-.
The Co4þ/Co3þ and Ni4þ/Ni3þ redox couples in the
layered oxides Li1-xCoO2 and Li1-xNiO2 are pinned at
the top of the O 2p bands. The systemLi1-xCoO2 shows a
flat Voc ≈ 4.0 V versus Li
þ/Li0 for 0 < x e 0.5 because
there is a coexistence of a polaronic, high-spin Co4þ in a
low-x phase80 and an itinerant-electron, low-spin Co4þ/
Co3þ phase near x = 0.5. For x > 0.5, peroxide forma-
tion at the surface leads to a loss of O2 by the surface
reaction
2ðO2Þ
2-
¼ 2O2-þO2v ð2Þ
A Voc ≈ 4.0 V is the intrinsic upper voltage limit for Li1-x-
CoO2. Decomposition of the compound occurs at higher
voltages.
The low-spin Cox
4þCo1-x
3þ:π*6-xσ*0 redox couple of
Li1-xCoO2 contains holes in the antibonding, itinerant π*
orbitals of t orbital parentage; the low-spin Nix
4þNi1-x
3þ:
t6σ*1-x couple of Li1-xNiO2 contains electrons in the
antibonding σ* orbitals of e orbital parentage, and for
x > 0.6 in Li1-xNiO2, holes become trapped in peroxide
ions. The initial voltage with the σ* orbitals is a little lower,
at V ≈ 3.8 V in layered Li1-xNiO2, which is why a greater
concentration of Ni4þ valence is found before O2 evolution
than with Co4þ in Li1-xCoO2. It is the cubic ligand-field
splitting of the octahedral-site 3d orbitals that raises the
Ni4þ/Ni3þ couple above that of the low-spin Co4þ/Co3þ
couple. Substitution of half of the Ni by Mn in Li-
(Ni0.5Mn0.5)O2 gives the formal valence states Ni
2þ and
Mn4þ. The Mn5þ/Mn4þ couple lies well-below the top of
Figure 11. Schematic representation of a slightly oxidized redox couple for different positions relative to the top of the anion p bands. (a) Itinerant versus
polaronic character of hole states of a couple on the approach to the topof the anionpband, (b) pinned couplewith predominantly antibonding (a.b.) anion
phole states andpredominantly cationdbonding (b.) states, and (c) couple too far below topof anionpband for significant cationd character inhole states.
Reprinted with permission from ref 79. Copyright 2009 Elsevier.

596 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
the O 2p bands, but the Ni3þ/Ni2þ couple is pinned at the
top of the O 2p bands. In Li1-x(Ni0.5Mn0.5)O2, the holes
occupy a σ* band of e2-2x parentage at the Ni:t6σ*2-2x, so
there is no step in the Fermi energy EF on passing from the
Ni3þ/Ni2þ to theNi4þ/Ni3þ couple.Moreover, theMn-Ni
interaction raises theNi4þ/Ni3þ redox couple relative to the
top of the O 2p bands to give access to the entire Ni4þ/Ni3þ
couple. Nevertheless, σ* orbitals of (e þ p)2-2x parentage
change to (p þ e)2-2x parentage as x increases.
The limiting lower voltage of an anode occurs where a
redox couple crosses the bottom of the broad cation
conduction band; e.g. the 4s band for the first-row
transition-metal atoms. This situation is illustrated by
Li insertion into the layered LiMS2 sheets to create a
coexistence of LiMS2 and Li2MS2 phases and therefore a
flat V versus x profile. The bottom of the M 4s band of
the monosulfides MS lowers progressively as theM atom
nuclear charge increases from M=Ti to M=Ni.81 In
LiTiS2 and LiVS2, the bottom of the 4s band is at ca. 0.15
eV below the Liþ/Li0 Fermi energy; it is only a little lower
at 0.85 eV in LiCrS2.
82 As is evident in Figure 12, an SEI
layer was formed rapidly on Li0.8þxMS2 (M = Ti, V) in
the voltage range 0.5 < V < 0.9 V versus Liþ/Li0 on the
first discharge; the Ti3þ/Ti2þ and V3þ/V2þ couples gave,
respectively, a V ≈ 0.5 and 1.0 V versus Liþ/Li0.
However, insertion of Li into LiCrS2 yielded, in addition
to the SEI layer, Cr0 þ Li2S at 0.85 V. The crystal-field
splitting of the 3d orbitals lifts the σ-bonding e orbitals
of high-spin Cr2þ:t3e1 above the bottom of the 4s band
in LiCrS2 even though CrS has been obtained chemically
by Jellinek.83 Since overlap of the S 3p orbitals with
the Cr 4s is larger than with the Cr 3d orbitals, the
covalent component of the Cr-S bond lifts the bottom
of the 4s band of CrS above the energy of the 3d4
configuration to allow access to the Cr2þ valence state.
In Li1þxCrS2, the sulfur atoms bond to Cr on one side
and to Li on the other. Moreover, the Li species are
forced into tetrahedral sites in Li2CrS2, and failure to
access Cr2þ implies that a strong covalent component
of the tetrahedral-site Li-S bond is reducing the cova-
lent component in the Cr-S bond to leave the bottom
of the Cr 4s band below the 3d4 energy. This observa-
tion illustrates how the bonding of a countercation can,
through the inductive effect, change the intrinsic limit-
ing lower voltage associated with a transition-metal
cation.
Effect of Cation Substitutions. A countercation can,
through the inductive effect, not only change an intrin-
sic limiting voltage associated with a transition-metal
cation but also be used to tune the energy of an operative
redox couple. This tuning phenomenon is illustrated in
Figure 5 by comparison of the voltages associated with
Figure 12. (a) Voltage profile for the discharge and charge curves on Li intercalation into Li0.8TiS2, Li0.8VS2, and LiCrS2 tested in the 1 M LiPF6 in EC:
DEC (1:1) electrolyte. (b) Corresponding positions of the bottom of the 4s band, the top of the S 3p band, theM3þ/M2þ, and theM4þ/M3þ redox couples
relative to the Fermi energy of lithium. Adapted from ref 82.

Review Chem. Mater., Vol. 22, No. 3, 2010 597
the Ti4þ/Ti3þ couples in the NASICON framework of
Li1þxTi2(PO4)3 and the spinel Li4Ti5O12. A change from
2.5 to 1.5 V shows that the countercation can have a
profound effect on a redox energy through the inductive
effect. Moreover, comparison of the Ni4þ/Ni3þ redox
energy in layered LiNiO2 where incomplete access is
found at 3.8 V with that in the spinel Li1-x[Ni0.5Mn1.5]O4
where complete access is found at 4.75 V, see Figure 13a,
shows that covalent bonding with the countercation can
also increase the intrinsic limiting voltage of a cathode by
lowering the top of theO 2p bands.We now inquire about
the effect of cation substitution for the active cation on
the intrinsic limiting voltage where the parent-cation
redox energy is pinned at the top of an anion p band.
For this purpose, we compare the influence of Mn4þ
versus Ti4þ substitutions on the energies of the Ni3þ/
Ni2þ and Ni4þ/Ni3þ couples in layered and spinel oxides.
Comparison with the influence of Cr3þ substitutions for
vanadium in the layered LiV1-yCryS2 sulfides is also in-
structive.
Removal of Li from the tetrahedral sites of the spinel
Li1-x[Ni0.5Mn1.5]O4 initially probes the Ni
3þ/Ni2þ and
then the Ni4þ/Ni3þ redox couples pinned at the top of the
O 2p bands. Figure 13a shows that the voltage increases
gradually with x from about 4.74 to 4.77 V in the first
charge curve. The irreversible capacity loss in the first
cycle corresponds to the formation of an SEI layer by the
oxidation of the electrolyte. The irreversible capacity loss
on each cycle is not characteristic of formation of a stable
SEI layer; it represents oxidation of the electrolyte and an
unstable SEI layer. Since the Ni4þ/Ni3þ couple of LiNiO2
gives an initial voltage of 3.8 V, we may expect a voltage
of 4.8 V versus Liþ/Li0 for extraction of Li from the
tetrahedral sites of the spinel if the top of the O 2p band
and the pinned Ni4þ/Ni3þ couple are both stabilized by
the shift of the Liþ ions from octahedral to tetrahedral
sites. Although the presence of the Mn4þ ions raises the
Ni4þ/Ni3þ redox energy relative to the top of the O 2p
bands, the voltage exceeds the HOMO of the electrolyte
at about 4.75 V before the full Ni4þ/Ni3þ couple is
accessed. In the spinel, the limitation is the HOMO of
the electrolyte; it is not the intrinsic voltage limit of the
cathode.
Figures 13b-d show what happens to the voltage
profile asMn is replaced by Ti in Li[Ni0.5Mn1.5-yTiy]O4.
84
The octahedral-site Mn5þ/Mn4þ couple appears to be
close enough to the top of the O 2p bands to be nearly
pinned; the Mn5þ ion is stable in tetrahedral sites with
a basic countercation. Interaction between the Mn5þ/
Mn4þ and Ni4þ/Ni3þ pinned couples allows access to
antibonding states at the top of the O 2p bands corre-
sponding to a formal Ni4þ/Ni3þ couple. Missing this
interaction, the Ni4þ/Ni3þ and Ni3þ/Ni2þ redox couples
fall below the top of theO 2pband.As a result, the voltage
increases with increasing Ti content in Li[Ni0.5Mn1.5-y-
Tiy]O4. However, the irreversible capacity loss is bigger
with increasing Ti content and at higher Ti content (y g
1.0) insertion of Li does not permit access to even the
Figure 13. Voltage profiles of Li1-x[Ni0.5Mn1.5-yTiy]O4: (a) y=0, (b) y=0.3, (c) y=0.5, and (d) y=1.The capacity decreases and the voltage increases
with higher Ti content. Adapted from ref 84.

598 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
Ni3þ/Ni2þ redox couple; there appears a large irrever-
sible curve at 4.9 V at y = 1.0. This is due to the fact that
the two nickel redox couples fall further below the top
of the O 2p bands with increasing Ti content in Li[Ni0.5-
Mn1.5-yTiy]O4. Hence, the irreversible flat curve at 4.9 V
corresponds to the irreversible access to the top of the O 2p
bands, which indicates the intrinsic voltage limit of the spinel
oxides. However, this can be confused with the oxidation
potential of the electrolytes; it is estimated around 4.75 V.
Figure 14 compares the voltage profiles of the layered
oxidesLi1-x(Ni0.5Mn0.5)O2
85 andLi0.9-x(Ni0.45Ti0.55)O2.
86
In these examples, the voltages are well below the 4.75 V
of the HOMO of the electrolyte. The Mn4þ raises the
energies of the two nickel couples relative to the top of
the O 2p bands, so the Ni4þ valence state is accessed
reversibly. On the other hand, the Ti4þ apparently lowers
the Ni3þ/Ni2þ couple relative to the top of the O 2p bands
sufficiently to limit the intrinsic voltage of the layered
oxide to under 4.0 V. Even the Ni3þ/Ni2þ couple is not
completely accessed in the presence of Ti4þ.
Figure 15 compares the voltage profiles of layered
oxides initially containing Ni2þ in the presence of Mn4þ
with layered sulfides initially containing V3þ in the pre-
sence ofCr3þ. In the oxides, theNi3þ/Ni2þ andNi4þ/Ni3þ
couples are pinned at the top of the O 2p bands; in the
sulfides, the V4þ/V3þ and V5þ/V4þ couples are pinned at
the top of the S 3p bands. Pinning of the redox couples of
nickel gives an initial V ≈ 3.7 V for Li1-x(Ni0.5Mn0.5)O2,
which is similar to the 3.8 V found61 for the Ni4þ/Ni3þ
couple of Li1-xNiO2. There is no step in the voltage
profile at x = 0.5 where the Fermi energy falls from the
Ni3þ/Ni2þ to theNi4þ/Ni3þ couple. This lack of a step is a
result of the pinning of the couples and the itinerant
character of the holes. It is to be contrasted with the steps
found, for example, in the NASICON structure,87 as is
illustrated in Figure 16. Finally, as already noted, the
Ni4þ valence state is accessed without the evolution of O2
because of the presence of Mn4þ. Similarly, Li1-x-
(V
0.5
Cr0.5)S2 shows a reversible charge/discharge profile
that varies smoothly through x = 0.5 because the V4þ/
V3þ andV5þ/V4þ couples are both pinned at the top of the
S 3p bands.79
Since only one Li can be removed from a layered oxide
and the nickel redox couples are both accessible in the
presence of Mn4þ, it was logical to investigate the oxides
Li(Ni0.25-yMn0.75-zLiyþz)O2
88,89 that are more easily
prepared with the Li well-ordered into the Li layers. Lu
and Dahn88 extracted Li from nominal Li1-x(Ni0.25-
Li0.167Mn0.583)O2 to obtain the voltage profile of the
second panel of Figure 15; we compare it with that for
Li1-x(V0.25Cr0.75)S2. At x = 0.5, the Ni
4þ valence is
reached in the oxide, the V5þ valence is reached in the
sulfide. In each, the voltage profile is flat for 0.5 < x<1
where the Fermi energy falls below the pinned redox
couple of antibonding states at the top of the anion p
bands; see Figure 17. A flat V = 4.5 V places the Fermi
energy above the HOMO of the electrolyte, which we
estimated to be at V ≈ 4.75 V. This observation means
that the flat voltage profile signals the voltage limit has
been reached, i.e. an EF in the O 2p band, rather than an
oxidation of the electrolyte. This situation must surely be
the case in the sulfide. At the intrinsic voltage limit, a
second phase appears in the electrode. Once the second
phase has been segregated in the oxide on the initial
charge, the electrode cycles with a reduced capacity in
the majority phase. In the sulfide, formation of the
second phase appears to be more reversible as if initially
disulfide ions are created on the surface before segrega-
tion of Li2S þ Cr2S3. Similarly, some peroxide ions may
form reversibly on the oxide before segregation of Li2O
and MnO2.
The third panel of Figure 15 shows that on further
decrease of the Ni concentration and increase of the Mn
concentration in nominal Li1-x(Ni0.17Li0.22Mn0.61)O2,
the onset of the flat V = 4.5 is introduced at a smaller x
with a subsequent reversible capacity similar to that of
Li1-x(Ni0.25Li0.167Mn0.583)O2 whereas Li1-x(V0.1Cr0.9)S2
exhibits an initial capacity fade that becomes reversible
with a reduced capacity after several cycles. These ob-
servations are consistent with the coexistence of two
phases in the electrode where the voltage becomes flat
with a reversible cycling once the second phase is segre-
gated out.
Figure 14. Voltage profiles of Li1-x[Ni0.5Mn0.5]O2 and Li0.9-x[Ni0.45Ti0.55]O2. Adapted from refs 85 and 86.

Review Chem. Mater., Vol. 22, No. 3, 2010 599
Indeed, Thackeray et al. 90 have argued that the attempt
to introduce excess Li homogeneously into the transition-
metal layers does not occur; but a coexistence of Li2-
MnO3 = Li(Li0.33Mn0.67)O2 layers is interleaved with
Li(Ni0.5-yMn0.5þy)O2 layers with the transformation
Li2MnO3 ¼ Li2OþMnO2 ð3Þ
occurring at V = 4.5 V.
Cathode SEI Layers. Extensive research has been de-
voted to characterization of the SEI layer formed on
lithium and on carbon anodes by reduction of the elec-
trolyte LiPF6 in EC/DEC;
8,63 this amorphous Li-electro-
lyte layer is complex, and the rate at which it is healed
after it is broken by changes in the electrode volume on
charge and discharge is difficult to measure accurately.
Preliminary work52,53 on the SEI layers formed on oxide
cathodes by an oxidative reaction of the electrodewith the
carbonate electrolyte indicates that these SEI layers are
generally unstable; the electrolyte is not protected from
further oxidation on subsequent cycling. A continued
electrode-electrolyte reaction on cycling thickens the
SEI layer, and progressive fading of the reversible capa-
city of the cathode is related to the thickening of the SEI
layer. Moreover, ambiguity in the measurement of the
onset of the oxidation reaction has given a reported
HOMO of the electrolyte LiPF6 in EC/DEC located at
4.5 ( 0.2 eV versus Liþ/Li0. This ambiguity may be
enhanced by a dependence of the oxidation voltage on
the SEI product, which can vary from one electrode
material to another. However, confusion between the
intrinsic voltage limit of an electrode and the HOMO
voltage may also contribute to this ambiguity.
Attempts to create a stable SEI passivation layer on an
oxide cathode have used two approaches: one seeks to
Figure 15. Voltageprofiles for the charge anddischarge curvesoncyclingofLi intercalation intoLi1-x[NiyLi(1/3-2y/3)Mn(2/3-y/3)]O2 andLi1-x[VyCr1-y]S2,
respectively. The flat voltage curves at 4.5 and 2.8 V indicate intrinsic voltage limits for the layered oxide and layered sulfide. Adapted from refs 89 and 79.

600 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
identify an additive91 to the electrolyte such as the EC
component for the carbon anode; the other attempts to
coat the cathode particles with a main-group oxide that is
permeable to Liþ ions.92 The former approach forms the
SEI layer in situ after the electrode particles have made
contact with the carbon of the particle/carbon composite
electrode, so the SEI layer formed does not interfere with
electronic contact between particles and the current col-
lector. The second approach has obtained some improve-
ment in cyclability, but complete coverage of the active
particles with a passivation layer before fabricating the
particle/carbon composite electrode would seem to in-
hibit electronic contact with the current collector. How to
coat the cathode/electrolyte interface with a stable SEI
layer while retaining electronic contact with the current
collector is a continuing challenge for cathodes that
would provide a V > 4.8 V versus Liþ/Li0 in the electro-
lyte LiPF6 in EC/DEC.
Capacity. The energy density of the cell of a recharge-
able battery is the product of the voltage V and the
capacity Λ of reversible charge transfer per unit weight
in amp hours per gram between the anode and the
cathode. The capacity of each electrode may be measured
separately versus Liþ/Li0 in a half-call with Li0 as the
anode. Three types of reversible electrode reactions have
been considered: (1) Li insertion into a transition-metal
oxide or sulfide host, (2) Li insertion into elements, and
(3) Li displacement reactions.
Transition-Metal Oxide or Sulfide Hosts. A transition-
metal oxide or sulfide host may be a layered compound or
a framework structure with 1D, 2D, or 3D interconnected
interstitial space for the guest Li atoms. Several examples
have been discussed above. The voltage given by the host
electrode was seen to be the energy of the operative
transition-metal redox couple. A flat voltage profile
versus Li concentration is preferred; it is found where
two phases coexist rather than where there is a solid
solution between the charged and discharged host. With
this strategy, the capacity is generally restricted to no
more than one Li per transition-metal atom; but where
the redox couple of a transition-metal atom is pinned at
the top of an anion p band, two redox couples per that
transition-metal ion may be accessed. This situation was
illustrated by the Ni in Li1-x(Ni0.5Mn0.5)O2 and by V in
Li1-x(V0.5Cr0.5)S2. However, as these examples illustrate,
it is not possible to take advantage of this accessibility
unless the host can accommodate more than one Li per
transition-metal atomwithout a voltage step. Framework
structures with a large interstitial space are needed to
obtain a capacity ofmore than oneLi per transition-metal
Figure 16. Voltage steps in the NASICON structure of Li3FeV(PO4)3.
Plateau A of the first discharge corresponds to the Fe3þ/Fe2þ redox
couple at 2.8 V; plateauB to theV3þ/V2þ redox couple at 1.7V; plateauC
in the second discharge to the V4þ/V3þ redox couple at 3.7 V. Adapted
from ref 87.
Figure 17. (a) Positions of theMnþ1/Mn redox couples relative to the Fermi energy of lithium in (a) Li0.5(V0.25Cr0.75)S2 and (b) Li0.5(Ni0.25Li0.17Mn0.68)O2;
2.8 eV in the layered sulfide and 4.5 eV in the layered oxide correspond to the flat curves in their voltage profiles in Figure 15.

Review Chem. Mater., Vol. 22, No. 3, 2010 601
atom, but such frameworks tend to be unstable if the large
cations that they form around are replaced by smaller
Liþ ions. On the other hand, the NASICON M2(XO4)3
framework is capable of receiving reversibly up to 5 Li
atoms and the M2(PS4)3 framework of AgTi2(PS4)3 illus-
trated in Figure 18 has been shown93,94 to accommodate
up to 10 Li atoms; but, reduction of the PS4 groups as well
as the Ti4þ to Ti0 destabilizes the framework. Moreover,
the host must be stable in the electrolyte; LiV2(PS4)3
dissolves in the electrolyte LiPF6 in EC/DEC.
Element Hosts.Any element as host intended for use as
an anode gives a voltage less than the voltage of the
LUMO of a carbonate electrolyte. Therefore, a passivat-
ing SEI layermust protect such an anode against chemical
reaction with the electrolyte. The most successful elemen-
tal host is carbon. Graphite has a layered structure with
three strong bonds in the graphite planes and half-filled pz
orbitals perpendicular to the planes that can interact with
the Li 2s orbitals. This bonding arrangement limits the
volume expansion on Li insertion, but also the number of
Li atoms that can be accommodated. Although graphite
can only accept one Li per six C atoms, the capacity is still
large, the volume expansion is manageable, and its vol-
tage changes little with the Li content. The situation is
similar with disordered carbon. Moreover, the EC com-
ponent and VC additive of a DEC or DMC electrolyte
creates a passivating SEI layer on carbon that is healed
quickly when cracked by the volume expansion of the
carbonmass onLi insertion.Nevertheless, the rate of Liþ-
ion transfer across the SEI layer and/or the rate of healing
of the SEI layer limits the safe rate of recharging of a
carbon anode. Safety concerns associated with plating of
Li on the surface of the anode and subsequent dendrite
formation require a design of large-scale power batteries
using carbon anodes that protects against failure of an
individual cell.
More Li can be inserted into Si than into C, which
makes it potentially an anode of exceptionally high
capacity, but a volume expansion of over 300% on full
charge is not manageable. Although the fabrication
of nanowires has been touted as a solution to this
problem,95,96 this anode suffers from continuing forma-
tion of an SEI layer that does not protect against inter-
action with the electrolyte. Similar problems are encoun-
tered with Li replacements for H in metal hydrides.
Displacement Reactions. Displacement of an element
from a compound or an alloy can be reversible,97,98 as was
first demonstrated with the displacement of iron from
Fe3O4 on insertion of more than one Li per formula
unit.99 Exploitation of this phenomenon promised a large
Figure 18. (a) Structural units and (b) projection in the a-b plane for AgTi2(PS4)3. (c) Discharge and charge curves for AgTi2(PS4)3 over five cycles at (a)
1.5-3.5 V at 0.1 mA/cm2. Reprinted with permission from refs 93 (Copyright 2008 American Chemical Society) and 94 (Copyright 2008 Elsevier).

602 Chem. Mater., Vol. 22, No. 3, 2010 Goodenough and Kim
capacity since a displacement may involve reducing
the displaced atom by more than one electron.100,101
These reactions are of particular interest for anodes with
main-group atoms; they offer voltages less than the
voltage of the LUMO of a carbonate electrolyte. How-
ever, this strategy suffers from large volume changes in a
charge/discharge cycle and the need for a stable passivat-
ing layer that can healmore quickly than reactionwith the
bulk electrolyte or the excess Liþ ions at the surface. This
approach is only promising if the LUMO of the electro-
lyte can be raised above the μAof the anode.Nevertheless,
amorphous Si or displacement reactions that are buffered
with graphite can improve the capacity of a carbon anode.
Summary
The principal challenges facing the development of
batteries for electric vehicles are cost, safety, cell energy
density (voltage  capacity), rate of charge/discharge,
and service life. Automation of manufacturing, material
selection, and service life are the keys to lower costs; the
availability of Li need not be a problem. Long service life
requires elimination of unwanted chemical reactions
between the electrodes and the electrolyte as well as
retention over many charge/discharge cycles of the elec-
tronic contact between the active particles of an electrode
and the current collector. The latter requirement restricts
the volume change versus state of charge that can be
tolerated in an electrode unless the active electrode par-
ticles are tethered to a current collector. This attachment
may bemade directly by the growth of nanowires of active
material on a current-collector substrate; it may also be
done by bonding the active particles to a conductive
polymer having a conduction band that overlaps the μA
or μC of the electrode. The former rate is being researched
with Si nanowires,95 the latter has been accomplished with
polypyrrole bonded to carbon-coated LiFePO4.
102,103
Safety is related to the flammability of the electrolyte,
the rate of charge and/or discharge, and the engineering
of the battery pack. A hybrid ionic liquid and organic
liquid as the electrolyte solvent can be made nonflam-
mable without too great a compromise of the electrolyte
σLi. Both the LiFePO4 cathode and the Li4Ti5O12 anode
have demonstrated safe and rapid charge and discharge
over many cycles where the μC and μA, respectively, are
located within the window of the electrolyte, which
removes the requirement of a passivating layer. However,
location of the μC and μA within the electrolyte window is
not sufficient if the electrode is operating at its voltage
limit with a redox couple pinned to the top of an anion p
band or overlapping a cation 4s band. An electrolyte with
a larger window, especially one with a higher LUMO,
would be helpful; but EC and/or VC in the electrolyte
solvent do allow use of a carbon anode with a limited
charging rate. Finally, engineering the battery-pack
design to allow failure of one cell without compromising
the entire stack can also build in safety.
A greater energy density requires both a larger voltage
and larger capacity. To achieve a greater energy density
with inexpensive electrodes having a long service life is a
challenge for the materials scientist. On the anode side, it
will be difficult to have a better capacity than that of
carbon, and the passivation layermakes it possible to take
advantage of a μA above the LUMO of the electrolyte,
albeit at the expense of a reduced charging rate. On the
cathode side, a material having a redox couple pinned at
the top of the O 2p band, but with two valence states
accessible without exceeding the intrinsic voltage limit of
the oxide, would offer the maximum voltage and capa-
city. The upper intrinsic voltage limit of such an oxide
can be raised by increasing the covalent component of
the M-O bond of a countercation to lower the energy
EV of the top of the O 2p bands; but this tuning of EV
must not also push the pinned redox couple further than
the change inEV. Comparison of the upper voltage limits
of the nickel oxides having Li in octahedral versus
tetrahedral sites illustrates a tuning of EV by 1.0 V.
However, to take advantage of the availability of two
redox couples on a single transition-metal atom, it is
necessary to have a host structure that can accept two Li
atoms per transition-metal cation. Can nature and/or
the imagination of the solid state chemist identify such
an electrode material with a μC matched to the electro-
lyte window? An example of the imaginative materials
design that may be needed is given by a recent report
from Sun et al.104
Acknowledgment. This work was supported by the office
of FreedomCAR and Vehicle Technologies of the U.S.
Department of Energy under contract no. DE-AC03-
76SF00098 and the Robert A.Welch Foundation of Houston,
TX (Grant No. F-1066).
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