(i) Electronic configuration : The valence electronic configuration of atoms of the group II A elements is ns2, where ‘n’ is the period number.
(ii) Atomic and ionic sizes : The size of the atom increases gradually from Be to Ra. Their ions are also large and size of the ion increases from Be2+ to Ra2+.
(iii) Ionisation enthalpy : The 1st and 2nd ionisation energies of these metals decrease from Be to Ba as size increases.
(iv) Melting and boiling points : Due to the presence of two electrons in the valence shell and stronger bonding in solid state, they have higher melting and boiling points than corresponding alkali metals.
(v) Metallic character : Due to low ionisation energy values, these metals are highly electropositive and readily form M2+ ions.
(vi) Flame colouration : Except Be and Mg, other members impart characteristic colours when their salts are introduced in the flame.
(vii) Hydration energy : The M2+ ions of alkaline earth metals are extensively hydrated to form [M(H2O)x]2+ ions and during hydration a huge amount of energy called hydration energy is released.
M2++xH2O → [M(H2O)x]2+ + Energy
(viii) Reactivity towards air or oxygen : They react with air or oxygen slowly on heating. Be, Mg and Ca form normal oxides (MO) while Sr and Ba form superoxides (SrO2,BaO2). BeO is amphoteric, MgO is weakly basic and others are distinctly basic.
(ix) Reactivity towards water : They have lesser reactivity towards water. Be does not react even with boiling water. Mg forms Mg(OH)2 liberating H2 gas with boiling water.
(x) Reactivity towards halogens : All the metals of the group combine with various halogens at appropriate temperature forming halides of the formula MX2.
(xi) Tendency to form complexes : Due to their smaller ionic sizes and greater charge densities Be, Mg metals have highest tendency to form complexes.
BeF2+2F− → [BeF4]2−
(xii) Reducing character : Except Be all other metals of this group are reducing agents.