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Intercalation Chemistry and Battery Materials
M. Stanley Whittingham
Chemistry Dept and Institute for Materials Research,
State University of New York at Binghamton
Binghamton, NY 13902-6016
It was recognized, about 30 years ago, that most
electrode reactions in reversible batteries are intercalation
reactions. In these reactions lithium or hydrogen are
inserted into a host matrix with essential retention of the
crystal structure. Surprisingly very little was known about
the reactions of lithium with oxides and sulfides, although
the heavier alkali metal reactions were being extensively
studied by the Rouxel and Hagnenmuller/Delmas groups
in France. Exxon workers found the unique properties of
lithium titanium disulfide, which has become the
prototypical ideal intercalation cathode [1, 2], showing
essentially complete reversibility in lithium reactions at
very high rates, 5-10 mA/cm2 even at high cathode
loadings. It met or exceeded the criteria for an advanced
cathode:
1. Reacts with 1 lithium per transition metal in a
totally reversible manner over many cycles.
2. Maintains its structural integrity during reaction.
2
3. Reacts rapidly,
.
4. Is low cost, environmentally harmless and
preferably is a good electronic conductor.
Exxon marketed a battery based on a LiAl anode
and a titanium disulfide cathode, and demonstrated that
such rechargeable lithium batteries could be built in the
discharged state [1]. However, the potential generated in a
lithium titanium disulfide cell is only around 2 volts,
insufficient for many applications. This is compounded
when used in combination with an intercalation material,
such as aluminum or graphite, as the anode.
Goodenough and co-workers found that lithium
cobalt oxide solved [3] the voltage challenge, and is the
basis of the popular SONY Li-ION cell. However, the
energy density of that cell is no higher than titanium
disulfide because only around half a lithium can be
reversibly cycled. In this case the oxide layers take up a 3block structure (3R) rather than the 1-block 1T of TiS2
and LiTiS2. CoO2 itself is isostructural with TiS2, as are
probably all the lithium free dioxides formed by deintercalation of the layered LiMO2 phases. Cobalt is too
expensive for large scale applications, so many solid state
ionic devices are still limited by the cathode material.
Thus, the search is on for new oxide materials that
combine the positive attributes of titanium disulfide and
cobalt oxide. Such oxides must show rapid reaction to
give high current densities or equilibration within a
reasonable time period. In the case of lithium batteries
this dictates an open crystalline lattice that can
accommodate lithium ions. The oxide of choice should
meet the criteria listed above.
Manganese and vanadium oxides are the most
likely candidates. We, and others, have been
investigating, in a systematic manner, the synthesis of a
range of manganese and vanadium oxides to ascertain the
role of structure on electrochemical behavior.
Layered manganates, e.g. KyMnO2•nH2O have
been synthesized by the hydrothermal decomposition of
alkali permanganate solutions [4]. These compounds have
the same structures as the titanium and cobalt compounds
discussed above. They show reversibility in lithium cells,
but the capacity fades on cycling with the potassium
compound showing the best behavior, sodium
intermediate and lithium the worst. This is due to the
formation of spinel-like phases such as LiMn2O4. These
structures only differ in the distribution of the manganese
and lithium ions in the essentially cubic close-packed
oxygen lattice [4, 5]. However, in the potassium
compound, the layers are staggered giving trigonal
prismatic coordination around the potassium ions, which
should be unfavorable for manganese migration. These
potassium ions act as pillars in the lattice.
The layered manganates can be stabilized by either
including pillars to prevent the oxygen ions from stacking
in ccp arrangement or by modifying the electronic
structure, by partial substitution of a part of the
manganese by iron, nickel or cobalt, so that they behave
more like LiCoO2 or LiTiS2 where the layer phases are
more stable than the spinel phases. This substitution
increases the electrical conductivity of the manganese
oxide [6], presumably lessening the effect of the JahnTeller distortion. Cycling of cobalt-substituted layered
manganese oxide has been found to be enhanced on
cobalt substitution, by for example Armstrong et al [7].
Pacific Lithium showed quite dramatically that stable
layer compounds can be formed by mixing chromium and
manganese, for example in Li1.2Mn0.4Cr0.4O2. High
capacities are obtained but the capacity is associated with
cycling of chromium between the 3+ and 6+ oxidation
states. Subsequently significant work has been ongoing in
Japan, the USA and Canada on related substituted
materials such as LiMn0.4Ni0.4Co0.2O2, which show
potentially interesting behavior.
Recently a number of other compounds have been
studied. For example, vanadium oxides with structures
related to the δ-V2O5 lattice are showing large capacities,
> 200 mAh/g. Iron phosphates, pioneered by the
Goodenough group, are showing possibilities as low cost
alternatives for cobalt oxide where large capacities are not
critical. Some of these future trends and directions will be
discussed, including related anode reactions.
The work was supported by the Department of
Energy through the Office of Transportation Technologies
and by National Science Foundation under grant
DMR9810198.
References:
1. M. S. Whittingham, U.S. Patent 4009052 and U.K.
Patent 1468416, (1973)
2. M. S. Whittingham, Science, 192 (1976) 1126.
3. K. Mitzushima, P. C. Jones, P. J. Wiseman, and J. B.
Goodenough, Mat. Res. Bull., 17 (1980) 785.
4. R. Chen and M. S. Whittingham, J. Electrochem. Soc.,
144 (1997) L64.
5. M. S. Whittingham, The Relationship between
Structure and Cell Properties of the Cathode for
Lithium Batteries, in Lithium Batteries, O. Yamamoto
and M. Wakihara, Editor. 1998, Kodansha: Tokyo.
6. P. Sharma, G. Moore, F. Zhang, P. Y. Zavalij, and M.
S. Whittingham, Electrochem. Soc. Letters, 2 (1999)
in press.
7. A. R. Armstrong, R. Gitzendanner, A. D. Robertson,
and P. G. Bruce, Chem. Commun., (1998) 1833.