Lithium (Li)

Lithium has the lowest density among metals (about 0.534 g cm⁻³) because its atoms are small and pack into a relatively open metallic lattice. Its valence configuration is \(1s^2\,2s^1\), placing it in Group 1 (alkali metals).

Lithium reacts exothermically with water to form lithium hydroxide and hydrogen gas:

\(\mathrm{2\,Li(s) + 2\,H_2O(l) \rightarrow 2\,LiOH(aq) + H_2(g)}\)

In moist air, lithium forms a thin oxide/nitride/hydroxide layer that dulls its surface. It is commonly stored under mineral oil or in an inert atmosphere to prevent oxidation.

Lithium gives a crimson red flame. Thermal excitation promotes valence electrons to higher energy levels; when they relax, photons in the red region are emitted, producing characteristic emission lines observable in flame tests and atomic spectra.

Due to the small \(\mathrm{Li^+}\) ion, lithium compounds show high lattice enthalpies and significant polarization of large anions (Fajans' rules). This increases covalent character, lowering solubility compared with Na/K analogues. For example, Li₂CO₃ is sparingly soluble, whereas Na₂CO₃ is readily soluble.

Lithium (Group 1, Period 2) and magnesium (Group 2, Period 3) exhibit a diagonal relationship with similarities such as:

• Formation of less soluble carbonates and phosphates.
• Formation of thermally stable nitrides on heating in nitrogen.
• Pronounced covalent character in many compounds due to higher polarization power.

Key reasons:

• Very low atomic mass → high specific capacity (mAh g⁻¹).
• High standard reduction potential \(E^\circ(\mathrm{Li^+/Li}) \approx -3.04\,\mathrm{V}\), enabling high-voltage cells.
• Intercalation of \(\mathrm{Li^+}\) into layered/olivine/spinel hosts allows reversible charge storage.

Overall cell reactions are host-dependent, but the driving force is the large free energy change for \(\mathrm{Li^+}\) insertion/extraction.

Li-ion: Uses a carbonaceous or other host anode storing \(\mathrm{Li^+}\) (no bulk Li metal). Safer and more cycle-stable; widely used in electronics and EVs.

Li-metal: Uses metallic Li as the anode, giving higher energy density but challenges with dendrite growth, safety, and cycling. Solid-state electrolytes are an active research area to mitigate dendrites.

Organolithiums (e.g., \(\mathrm{n\text{-}BuLi}\), \(\mathrm{PhLi}\)) are strong bases and nucleophiles formed by metal–halogen exchange or direct metalation:

\(\mathrm{2\,Li + R\!\!\;X \rightarrow R\!\!\;Li + LiX}\)

They are used to generate carbanions, perform deprotonations (\(pK_a\) bases), and construct C–C bonds in advanced organic synthesis. They are air/moisture sensitive and must be handled under inert conditions.

Lithium has two stable isotopes: \(^6\!\mathrm{Li}\) and \(^7\!\mathrm{Li}\). \(^6\!\mathrm{Li}\) is used in tritium breeding via neutron capture:

\(\mathrm{^6Li + n \rightarrow ^4He + ^3H\;(T) + 4.78\,MeV}\)

Such reactions are of interest in fusion-energy research and neutron detection technologies.

Lithium reacts with water, acids, and humid air. For fires, do not use water or CO₂ extinguishers. Use a Class D (dry powder, e.g., graphite, copper powder, or sodium chloride) extinguisher.

• Store under dry mineral oil or inert gas.
• Avoid contact with moisture; use dry tools and gloves.
• Dispose of lithium waste in accordance with hazardous materials protocols.