Alkali Metals (Group 1)

Atomic radius increases from Li (152 pm) → Na (186 pm) → K (227 pm) → Rb (248 pm) → Cs (265 pm). Each element adds a new principal quantum shell, pushing the outermost electron far from the nucleus. The ionic radii (Li⁺ 76 pm, Na⁺ 102 pm, K⁺ 138 pm, Rb⁺ 152 pm, Cs⁺ 167 pm) are all substantially smaller than the parent atoms because removal of the lone valence electron eliminates the outer shell entirely, reducing electron–electron repulsion and allowing the remaining electrons to contract toward the nucleus.

First ionisation energy (IE₁) decreases down the group: Li 520, Na 496, K 419, Rb 403, Cs 376 kJ mol⁻¹. Two factors drive this trend: (i) increasing atomic radius places the valence electron farther from the nucleus so the Coulombic attraction is weaker (E ∝ 1/r), and (ii) greater shielding by inner-shell electrons reduces the effective nuclear charge (Z_eff) experienced by the valence electron. Second ionisation energies (IE₂) are always enormously larger — for Na, IE₂ ≈ 4562 kJ mol⁻¹ — because removing a second electron requires breaking into a filled noble-gas core.

When an alkali metal ion dissolves in water, ion–dipole forces attract water molecules, releasing hydration enthalpy (ΔH_hyd): Li⁺ −519, Na⁺ −406, K⁺ −322, Rb⁺ −301, Cs⁺ −276 kJ mol⁻¹. The trend reverses the radius trend: smaller ions have higher charge density (charge/volume) and polarise water more strongly. Li⁺ carries a primary hydration sphere of approximately 4–6 water molecules held very tightly, making its effective hydrated radius the largest of the group — a striking inversion that explains why Li⁺ moves more slowly in solution than K⁺ despite its smaller bare ionic size.

When alkali metal atoms are introduced into a gas flame (~1200 °C), thermal energy promotes the valence electron from the ground state (e.g., 3s for Na) to an excited state (3p). When the electron falls back, it emits a photon of energy ΔE = hν = hc/λ. For sodium, the 3p→3s transition produces the intense yellow doublet at 589.0/589.6 nm (the sodium D-lines). Lithium emits crimson at 671 nm, potassium lilac/violet at 766–770 nm, rubidium red-violet at 780 nm, and caesium blue at 455 nm. These unique fingerprints make flame tests a rapid qualitative test for Group 1 ions.

All alkali metals reduce water, releasing hydrogen gas and forming a metal hydroxide. The general equation is: \ce{2M(s) + 2H_2O(l) -> 2MOH(aq) + H_2(g)}. Reactivity increases sharply down the group. Lithium fizzes steadily; sodium melts into a silvery ball and skids; potassium ignites the H₂ with a lilac flame; rubidium and caesium react explosively. The increasing reactivity reflects decreasing ionisation energy: less energy is needed to release the valence electron, so the reaction releases more net energy per mole.

The nature of the oxide product changes down the group — a subtle but important trend. Lithium forms primarily the normal oxide (\ce{4Li + O_2 -> 2Li_2O}). Sodium forms mainly the peroxide (\ce{2Na + O_2 -> Na_2O_2}) with some normal oxide. Potassium, rubidium, and caesium form superoxides (\ce{K + O_2 -> KO_2}). The larger, more polarisable ions stabilise the larger, more diffuse peroxide (O₂²⁻) and superoxide (O₂⁻) anions through lattice energy considerations — smaller Li⁺ favours the compact O²⁻ ion which maximises lattice energy.

• \ce{Na} — Sodium metal (solid): Group 1 alkali metal acting as reducing agent; loses one electron to form Na⁺
• \ce{H_2O} — Water (liquid): Acts as oxidising agent; H⁺ is reduced to H₂ gas (oxidation state 0)
• \ce{NaOH} — Sodium hydroxide (aqueous): Strong alkali formed; fully dissociates in solution to Na⁺ and OH⁻ ions
• \ce{H_2} — Hydrogen gas (gas): Flammable by-product; can ignite spontaneously in reactions of K, Rb, Cs with water

Problem. A student performs flame tests on three unknown Group 1 solutions and observes: Solution A — crimson flame; Solution B — yellow flame; Solution C — lilac flame. (i) Identify each metal ion. (ii) Explain why sodium gives such an intense yellow colour compared with the others. (iii) Arrange Li⁺, Na⁺, and K⁺ in order of decreasing hydration enthalpy and explain the trend.

Answer. A = Li⁺ (crimson, 671 nm); B = Na⁺ (yellow, 589 nm); C = K⁺ (lilac, 766 nm). Sodium's yellow is intense due to high oscillator strength and low excitation energy of the 3p state. Hydration enthalpy order: Li⁺ > Na⁺ > K⁺, because smaller ions have higher charge density and stronger ion–dipole interactions.

• Alkaline Earth Metals (Group 2) — Group 2 elements share the s-block with Group 1 and display analogous but attenuated trends in radius, IE, and reactivity due to the addition of a second valence electron and higher nuclear charge.
• Ionic Bonding and Lattice Energy — Lattice energy, governed by the Born–Landé equation, determines why Li₂O is stable while Cs forms CsO₂; understanding it requires knowledge of ionic radii and charge, central to this topic.
• Enthalpies of Solution and Hydration — The Born–Haber cycle for dissolving alkali metal salts directly uses hydration enthalpies of M⁺ ions, linking this topic to thermodynamic cycles and solubility trends.
• Atomic Emission Spectra — Flame coloration is a direct consequence of electronic transitions; understanding quantised energy levels and emission lines deepens the explanation of why each alkali metal shows a unique colour.
• Standard Electrode Potentials — Despite having the lowest IE₁, lithium has the most negative standard electrode potential (−3.04 V) due to its enormous hydration enthalpy, illustrating how thermodynamic cycles govern electrochemical behaviour.