The dipositive oxidation state (M2+) is the predominant valence of Group 2 elements.
The alkaline earth metals form compounds which are predominantly ionic but less ionic than the corresponding compounds of alkali metals. This is due to increased nuclear charge and smaller size.
The oxides and other compounds of beryllium and magnesium are more covalent than those formed by the heavier and large sized members (Ca, Sr, Ba).
The general characteristics of some of the compounds of alkali earth metals are described below.
Oxides and Hydroxides: The alkaline earth metals burn in oxygen to form the monoxide, MO which, except for BeO, have rock-salt structure.
The BeO is essentially covalent in nature. The enthalpies of formation of these oxides are quite high and consequently they are very stable to heat.
BeO is amphoteric while oxides of other elements are ionic in nature. All these oxides except BeO are basic in nature and react with water to form sparingly soluble hydroxides.
MO +H2O → M(OH)2
The solubility, thermal stability and the basic character of these hydroxides increase with increasing atomic number from Mg(OH)2 to Ba(OH)2.
The alkaline earth metal hydroxides are, however, less basic and less stable than alkali metal hydroxides. Beryllium hydroxide is amphoteric in nature as it reacts with acid and alkali both.
Be(OH)2 + 2OH- → [Be(OH)4]2+
Be(OH)2 + 2HCl + 2H2O → [Be(OH)4]Cl4
Halides: Except for beryllium halides, all other halides of alkaline earth metals are ionic in nature. Beryllium halides are essentially covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state as shown below:
In the vapour phase BeCl2 tends to form a chloro-bridged dimer which dissociates into the linear monomer at high temperatures of the order of 1200 K.
The tendency to form halide hydrates gradually decreases (for example, MgCl2·8H2O, CaCl2·6H2O, SrCl2·6H2O and BaCl2·2H2O) down the group.
The dehydration of hydrated chlorides, bromides and iodides of Ca, Sr and Ba can be achieved on heating; however, the corresponding hydrated halides of Be and Mg on heating suffer hydrolysis.
The fluorides are relatively less soluble than the chlorides owing to their high lattice energies.
Salts of Oxoacids: The alkaline earth metals also form salts of oxoacids. Some of these are:
Carbonates: Carbonates of alkaline earth metals are insoluble in water and can be precipitated by addition of a sodium or ammonium carbonate solution to a solution of a soluble salt of these metals.
The solubility of carbonates in water decreases as the atomic number of the metal ion increases.
All the carbonates decompose on heating to give carbon dioxide and the oxide. Beryllium carbonate is unstable and can be kept only in the atmosphere of CO2. The thermal stability increases with increasing cationic size.
Sulphates: The sulphates of the alkaline earth metals are all white solids and stable to heat. BeSO4, and MgSO4 are readily soluble in water; the solubility decreases from CaSO4 to BaSO4.
The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water.
Nitrates: The nitrates are made by dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six molecules of water, whereas barium nitrate crystallises as the anhydrous salt.
This again shows a decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy. All of them decompose on heating to give the oxide like lithium nitrate.
2M (NO3)2 → 2MO + 4NO2 + O2 (M = Be, Mg, Ca, Sr, Ba)