Solubility Product Constant and Common Ion Effect
Abstract
The solubility product constant (KSP) is a powerful tool for determining the solubility of an ionic
compound. The purpose of this experiment was to calculate the K SP of Ca(OH)2 and determine the
effect of the common ion effect on this. The common ion effect further decreases the solubility of a
compound as the equilibrium will be pushed towards the reactants. Through titration of dissolved
Ca(OH)2 solutions, the KSP was calculated successfully. Overall, the results obtained showed that
Ca(OH)2 is weakly soluble and the addition of Ca 2+ hinders this further.
Introduction
When an ionic compound dissolves in water, the dissolution of its aqueous cations and anions will
result in a reversible state of equilibrium. Establishment of an equilibrium means the rate of a
forward reaction will be equal to the rate of a backward reaction – giving no net change to the
concentration of products or reactants. Equilibrium constants tell us the favoured direction of a
reaction once equilibrium has been reached, the solubility product constant is very similar to this.
The solubility product constant (KSP) is the equilibrium constant for an ionic compound dissolving,
indicating the solubility of the compound. The higher the K SP, the more soluble the compound is
within water. The equation for calculating K SP is indicated below, it can be regarded as the product of
the two reactant ions, however the stoichiometric ratio of products must be taken into account.
Generally, if ionic compound composed of an anion and cation that both have a charge of -1 and +1,
respectively, then the KSP would simply involve calculating the concentration of both resulting ions
and multiplying them. For more complex stoichiometric ratios, the equation is still practical and
easy-to-use, requiring the concentrations of ions to be raised to the power of the ratio.
AB (s) ⇌ aAz+ (aq) + bBy- (aq)
KSP = [Az+]a × [By-]b
Following the dissolution of an ionic compound, a titration using a reagent of a known molarity can
be carried out to measure the concentration of one of the resultant ions. Subsequently, the
Abstract
The solubility product constant (KSP) is a powerful tool for determining the solubility of an ionic
compound. The purpose of this experiment was to calculate the K SP of Ca(OH)2 and determine the
effect of the common ion effect on this. The common ion effect further decreases the solubility of a
compound as the equilibrium will be pushed towards the reactants. Through titration of dissolved
Ca(OH)2 solutions, the KSP was calculated successfully. Overall, the results obtained showed that
Ca(OH)2 is weakly soluble and the addition of Ca 2+ hinders this further.
Introduction
When an ionic compound dissolves in water, the dissolution of its aqueous cations and anions will
result in a reversible state of equilibrium. Establishment of an equilibrium means the rate of a
forward reaction will be equal to the rate of a backward reaction – giving no net change to the
concentration of products or reactants. Equilibrium constants tell us the favoured direction of a
reaction once equilibrium has been reached, the solubility product constant is very similar to this.
The solubility product constant (KSP) is the equilibrium constant for an ionic compound dissolving,
indicating the solubility of the compound. The higher the K SP, the more soluble the compound is
within water. The equation for calculating K SP is indicated below, it can be regarded as the product of
the two reactant ions, however the stoichiometric ratio of products must be taken into account.
Generally, if ionic compound composed of an anion and cation that both have a charge of -1 and +1,
respectively, then the KSP would simply involve calculating the concentration of both resulting ions
and multiplying them. For more complex stoichiometric ratios, the equation is still practical and
easy-to-use, requiring the concentrations of ions to be raised to the power of the ratio.
AB (s) ⇌ aAz+ (aq) + bBy- (aq)
KSP = [Az+]a × [By-]b
Following the dissolution of an ionic compound, a titration using a reagent of a known molarity can
be carried out to measure the concentration of one of the resultant ions. Subsequently, the
