Group II

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Group II

Reactivity of Group II elements increases down the group.

This can be explained by the increase in ease at losing two outer electrons as we descend the group. The loss of electrons becomes easier due to the decreasing ionisation energy required.

Ionisation energy decreases due to extra shielding from inner shells and increase in distance from the nucleus.

Group II elements are less reactive than the corresponding Group I element in the same period, since two rather than one electron need to be lost in order to gain noble gas configuration.

Reaction with Water

Reactivity increases down the group. This is clearly seen if we observe the reactions of magnesium and calcium in water.

a) Virtually no reaction occurs between magnesium and cold water. However, in a reaction with steam it forms magnesium oxide and hydrogen.

Mg(s) + H2O(g) → MgO(s) + H2(g)

b) Calcium is more reactive. It reacts with cold water to produce an alkaline solution of calcium hydroxide and hydrogen gas is released.

Ca(s) + H2O(l) → Ca(OH)2(aq) + H2(g)

Reaction with Air and Water

All group II elements react in air to form an oxide layer. Again the reactivity increases down the group, all forming a white oxide.

2Mg(s) + O2(g) → 2MgO(s)


The thermal stability of Group II carbonates increases as we descend the group. Beryllium carbonate decomposing most easily.

MCO3(s) → MO(s) + CO2(g)

Where M is a Group II element.

This trend is explained in terms of the Group II metal ions ability to polarise the anion, the carbonate ion. Small highly charged positive ions distort the electron cloud of the anion.

The larger the anion the easier the distortion, as seen with the carbonate ion. Hence polarising ability of the M2+ ion decreases down the group.

The greater the distortion caused by the polarising ion the less stable the compound is to heat. This means that beryllium carbonate decomposes at a lower temperature to the rest of the group.


The thermal stability of the nitrates follows the same trend as that of the carbonates, with thermal stability increasing with proton number.

The reason, once more, is that the polarising power of the M2+ decreases as ionic radius increases.

The solubility of these sulphates decreases as we descend the group, with barium sulphate being insoluble in water.

Two factors are involved in dissolving:

  1. Lattice energy
  2. Enthalpy of hydration.

Due to the large size of the sulphate anion there is little difference between the lattice energies for these compounds. However, due to the change in ionic radius (i.e. charge density), there is significant difference in terms of their ability to hydrate.

The greater the charge density the easier it is for the cation to hydrate and hence dissolve in water due to greater attraction with the polar water molecules.

Magnesium sulphate dissolves in water whereas barium sulphate does not.

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