General Characteristics of Alkali Metals

Occurrence. Alkali metals are so reactive that none of them is found free in nature -- they occur only combined, as compounds. Rubidium and caesium occur in only minute quantities, associated with the minerals of the other alkali metals. Francium is intensely radioactive and does not occur appreciably in nature; its longest-lived isotope has a half-life of only 21 minutes. Their relative abundance in the earth's crust and principal mineral sources are given in Table 5.1 -- lithium from spodumene, sodium from rock salt (NaCl), potassium from sylvite (KCl); rubidium has no convenient ore of its own and is recovered as a by-product of lithium processing. Figure 5.1 shows how the reactive metals Li, Na and K are kept stored under oil, away from air and moisture.

Electronic configuration. Every alkali metal has the general valence-shell configuration ns¹, where n is the period number (Table 5.2) -- a single, loosely-held electron outside a closed noble-gas core.

Common oxidation state. All the alkali metals are highly electropositive. They readily lose this single valence electron to form a monovalent cation, M⁺, and consequently show only one oxidation state, +1, in all their compounds.

Atomic and ionic radii. Being the very first element of each period, every alkali metal has the largest atomic and ionic radius in its own period. Down the group, an extra electron shell is added at each step, so both the atomic radius and the ionic radius steadily increase from Li to Cs (Table 5.3). As expected, each M⁺ ion is smaller than its parent atom, having lost its only valence-shell electron.

Ionisation enthalpy. Alkali metals have the lowest first ionisation enthalpy of any element in their respective period (Table 5.3). Down the group the ionisation enthalpy decreases, because the atomic size increases and, along with it, the number of inner shells (and hence the screening/shielding effect on the valence electron) also increases -- both effects make the lone valence electron progressively easier to remove. Their second ionisation enthalpies, by contrast, are extremely high: once the single valence electron is removed, the resulting M⁺ ion already has a stable, noble-gas-like electronic configuration, so pulling out a second electron from that stable core requires a very large amount of energy.

Hydration enthalpy. Lithium salts are markedly more soluble than the corresponding salts of the other group 1 metals. For example, LiClO₄ is up to 12 times more soluble than NaClO₄, while KClO₄, RbClO₄ and CsClO₄ have solubilities only about 10⁻³ times that of LiClO₄. This exceptional solubility of Li⁺ salts traces to the very small size of the Li⁺ ion, which lets it be strongly solvated (hydrated) by water molecules -- the hydration enthalpy of Li⁺ is the largest in magnitude among the alkali metal cations and falls steadily down the group (Figure 5.2): Li⁺ (-519) > Na⁺ (-406) > K⁺ (-322) > Rb⁺ (-293) > Cs⁺ (-264 kJ mol⁻¹).

Electronegativity. Alkali metals have a comparatively small electronegativity value relative to other elements of their period (Table 5.3), and it falls further down the group. Because of this low electronegativity, when they react with other (more electronegative) elements they typically form ionic compounds -- for example, reacting with the halogens gives ionic halides. …