Equilibrium in a system having more than one phase is called heterogeneous equilibrium. The equilibrium between water vapour and liquid water in a closed container is an example of heterogeneous equilibrium.
H2O(l) ⇔ H2O(g)
In this example, there is a gas phase and a liquid phase. In the same way, equilibrium between a solid and its saturated solution.
Ca(OH)2 (s) + (aq) ⇔ Ca2+(aq) + 2OH-(aq) is a heterogeneous equilibrium
Heterogeneous equilibria often involve pure solids or liquids. We can simplify equilibrium expressions for the heterogeneous equilibria involving a pure liquid or a pure solid, as the molar concentration of a pure solid or liquid is constant (i.e., independent of the amount present).
In other words if a substance ‘X’ is involved, then [X(s)] and [X(l)] are constant, whatever the amount of ‘X’ is taken. Contrary to this, [X(g)] and [X(aq)] will vary as the amount of X in a given volume varies.
Let us take thermal dissociation of calcium carbonate which is an interesting and important example of heterogeneous chemical equilibrium.
CaCO3(s) ⇔(∆) CaO (s) + CO2 (g)
On the basis of the stoichiometric equation, we can write,
KC = [CaO(s)][CO2(g)] / [CaCO3(s)]
Since [CaCO3(s)] and [CaO(s)] are both constant, therefore modified equilibrium constant for the thermal decomposition of calcium carbonate will be
K´c = [CO2(g)] or Kp = PCO2
This shows that at a particular temperature, there is a constant concentration or pressure of CO2 in equilibrium with CaO(s) and CaCO3(s).
Experimentally it has been found that at 1100 K, the pressure of CO2 in equilibrium with CaCO3(s) and CaO(s), is 2.0 ×105 Pa. Therefore, equilibrium constant at 1100K for the above reaction is:
Kp = PCO2 = 2×105Pa/105Pa = 2.00
Similarly, in the equilibrium between nickel, carbon monoxide and nickel carbonyl (used in the purification of nickel),
Ni (s) + 4 CO (g) ⇔ Ni(CO)4 (g)
the equilibrium constant is written as
KC = [Ni(CO)4]/[CO]4
It must be remembered that for the existence of heterogeneous equilibrium pure solids or liquids must also be present (however small the amount may be) at equilibrium, but their concentrations or partial pressures do not appear in the expression of the equilibrium constant. In the reaction,
Ag2O(s) + 2HNO3(aq) ⇔ 2AgNO3(aq) +H2O(l)
KC = [AgNO3]2/[HNO3]2