Heterogeneous Equilibria

Full Notes

What is Heterogeneous Equilibrium?

In many chemical reactions, reactants and products exist in different phases – such systems are called heterogeneous equilibria. This contrasts with homogeneous equilibria, where all substances are in the same phase.

Examples:

• Solid-gas systems
• Solid-liquid systems
• Liquid-gas systems

A classic example:
CaCO₃(s) ⇌ CaO(s) + CO₂(g)

In this reaction, calcium carbonate and calcium oxide are solids, while carbon dioxide is a gas.

Equilibrium Constant Expression

In heterogeneous systems, we write the equilibrium constant expression using only those species whose concentrations can change – that is, gases and solutes.

Concentrations of pure solids and pure liquids are constant and are excluded from the expression.

Example: Decomposition of Calcium Carbonate

CaCO₃(s) ⇌ CaO(s) + CO₂(g)

K_c = [CO₂]

K_p = P_CO2

Solid CaCO₃ and CaO are omitted from the equilibrium expression.

Why Are Solids and Liquids Omitted?

From a thermodynamic perspective:

• The molar concentration of a pure solid or liquid is fixed because it's defined by its density and molar mass.
• It does not change with the amount present and hence has no effect on the position of equilibrium.

Example: Ammonium Chloride Sublimation

NH₄Cl(s) ⇌ NH₃(g) + HCl(g)

K_c = [NH₃][HCl]

K_p = P_NH3 × P_HCl

NH₄Cl(s) is omitted as it is a solid.

Additional Notes

• When writing equilibrium expressions:
  • Include gases and aqueous solutions (their concentrations/partial pressures vary).
  • Exclude solids and pure liquids.
• Temperature still affects K just as it does in homogeneous systems.

Summary

• Heterogeneous equilibria involve different phases.
• Omit pure solids and liquids from K expressions.
• Only gaseous and aqueous species appear in K_c and K_p.
• Temperature changes affect the value of K.