The Intercalation Host

Lithium, Explosive Dendrites, and the Layered Oxide Cathode

Element number three on the periodic table—lithium, comprising a helium core of two inner electrons surrounded by a single $2s^1$ valence electron—is the lightest of all solid elements. It has an atomic mass of 6.94 grams per mole and a metallic density of $0.534\text{ g/cm}^3$. Because that lone $2s$ electron is shielded by the inner $1s^2$ shell, lithium surrenders its valence electron more readily in solution than any other metal, yielding a standard reduction potential of $-3.04\text{ volts}$ relative to the standard hydrogen electrode. Combined with Faraday's constant ($96,485\text{ coulombs per mole}$), one gram of metallic lithium stores 3,861 milliampere-hours of electrical charge—nearly five times the theoretical capacity of zinc or lead.

Yet for decades, electrochemists could not build a safe rechargeable battery around metallic lithium. In the mid-1970s at Exxon's corporate research laboratories in Linden, New Jersey, chemist M. Stanley Whittingham constructed the first rechargeable lithium cell by pairing a pure lithium metal foil anode with a titanium disulfide ($\text{TiS}_2$) cathode. During discharge, lithium atoms stripped off their electrons at the anode and migrated across a liquid organic electrolyte as tiny $\text{Li}^+$ ions ($0.076\text{ nanometers}$ in radius). At the cathode, those ions slid smoothly between the van der Waals layers of the $\text{TiS}_2$ crystal—a reversible host-guest process termed intercalation.

When Whittingham's Exxon cells were repeatedly recharged, however, metallic lithium did not plate back onto the anode foil in smooth, flat layers. Microscopic surface irregularities concentrated the electric field lines at tiny bumps on the foil. Incoming $\text{Li}^+$ ions preferentially deposited onto these high-field tips, growing sharp, branching metallic needles known as dendrites. Within dozens of charge cycles, lithium dendrites pierced the porous polymer separator between the electrodes, creating a dead internal short circuit. Because liquid organic carbonate electrolytes are highly flammable and metallic lithium melts at $180.5^\circ\text{C}$, the shorted cells vented and burst into flame.

Eliminating the explosive metallic anode required doubling the voltage of the cathode so that both electrodes could operate as intercalation hosts without sacrificing energy density. In 1980, working at the Inorganic Chemistry Laboratory at the University of Oxford, solid-state physicist John B. Goodenough recognized that Whittingham's sulfide cathode ($\text{TiS}_2$) was limited to 2.2 volts because sulfur's $3p$ orbitals sit relatively high in energy. Replacing sulfur ($S^{2-}$) with smaller, more electronegative oxygen anions ($O^{2-}$) lowered the cathode's electronic energy bands and enabled access to the deep $\text{Co}^{3+}/\text{Co}^{4+}$ redox couple:

Goodenough and his postdoctoral researchers Koichi Mizushima and Philip Jones synthesized lithium cobalt oxide ($\text{LiCoO}_2$), a crystal adopting the layered alpha-$\text{NaFeO}_2$ rock-salt structure in which two-dimensional sheets of edge-sharing $\text{CoO}_6$ octahedra alternate with planar galleries of mobile $\text{Li}^+$ ions. Extracting up to half of the lithium ions ($x = 0.5$) raised the cell voltage from 2.2 volts to 4.0 volts while expanding the crystal lattice volume by less than three percent.

Five years later in Kawasaki, Japanese chemist Akira Yoshino at Asahi Kasei completed the modern rocking-chair cell architecture by replacing the dangerous metallic lithium anode with graphitic carbon derived from petroleum coke. During charging, lithium ions slip between the planar hexagonal carbon sheets to form $\text{LiC}_6$ at just 0.1 volts above metallic lithium, storing one lithium ion for every six carbon atoms without plating free metal. Even inside this dual-intercalation cell, lithium's tiny $0.076\text{-nanometer}$ ionic radius imposes a residual kinetic bottleneck. Its intense local electric field binds polar solvent molecules into a bulky solvation shell ($0.4\text{ nanometers}$ across), and stripping away that solvent cage at the graphite interface requires overcoming a $50\text{ kJ/mol}$ activation barrier that governs fast-charging limits at freezing temperatures.