Common Ion Effect

When some other electrolyte (that includes an ion that is also included in the first electrolyte — in other words, a common ion) is added, the common ion impact suppresses the ionisation of the first electrolyte. It is believed to result from Le Chatelier’s principle (or the Equilibrium Law).

State Common Ion Effect?

The common ion effect could be stated as follows: in a solution with the several species trying to associate together via a chemical equilibrium process, increasing the concentration of one of the ions dissociated in the solution by adding another species usually contains the same ion will boost the level of ion association.

The common ion effect occurs after gaseous hydrogen chloride is passed across a sodium chloride solution, causing the NaCl to precipitate due to excess chloride ions present in the solution (brought on by the dissociation of HCl).

Effect on Solubility

Adding a Common Ion

When a salt solution is added to a common ion (of the salt), the salt precipitates, so because ions are now being used to make the salt precipitate, their concentrations will drop. The concentration of ions may continue to drop until solubility equilibrium is attained.

To use the common ion concept as an example, adding zinc ions to a zinc hydroxide solution would result in zinc as the common ion.

Adding solution containing a Common Ion

Whenever you add a solution containing a common ion, the effect upon solubility is the same as if you merely added the common ions. The dissociation of the incoming solution causes an increase in salt precipitation and just a decrease in ion concentration. The solubility equilibrium is generally unaffected by the other ion in the incoming solution.

Using the common ion definition, adding sodium hydroxide to the zinc hydroxide solution will result in hydroxide as the common ion.

Common Ion Effect in Buffer Solutions

Whenever a conjugate ion is added to a buffer solution, the pH of the solution changes due to the common ion effect.

Consider an acid buffer solution that has been added and mixed with a strong electrolyte (i.e., conjugate base salt); the electrolyte will dissociate, creating both common ions plus ions that partly ionise the acid. They created common ions that will neutralise the acid’s ionisation. The equilibrium will shift in favour of the reactants according to Le Chatelier’s Principle (to the left). So the acid’s ionisation (and dissociation) reduces the pH rises.

Real-World Example of Common Ion Effect

When water is associated with natural things (such as rocks), it may include limestone and chalk once it is located underground. To reduce the hardness of water or the levels of magnesium and calcium, soluble sodium carbonate could be added. Calcium carbonate is formed. As a result, that is not particularly soluble.

Additional sodium carbonate is subsequently added, with carbonate becoming the common ion in this case. Calcium carbonate precipitates from the water concentration in the solution of this addition.

Common Ion Effect Example

As an example, consider a calcium sulphate solution. Calcium sulphate is in equilibrium with calcium ions and sulphate ions in a saturated solution. A small proportion of the calcium sulphate will dissociate into ions; however, the majority will stay as molecules.

CaSO4 (s) — Ca2+ (aq) + SO2-4 (aq)  Ksp = 2.4 × 10-5

To this saturated solution, add little calcium chloride. In this instance, the calcium ion concentration in the solution would rise. The product of [Ca2+] times [SO42-] will increase, eventually surpassing the Ksp. Le Chatelier’s principle states that the equilibrium here will shift to the left due to the extra calcium ion. More calcium sulphate will precipitate from the solution till the ion product equals the Ksp once more. Whenever equilibrium is restored, the concentrations of the calcium and sulphate ions will no longer be equal – the calcium ion concentration will be substantially more significant – but the product of the concentrations will remain similar to the Ksp.

This is an outstanding example of the common ion effect. When calcium chloride is added to a saturated calcium sulphate solution, it leads more calcium sulphate to precipitate out of the solution, diminishing its adequate solubility. You could achieve a comparable effect by adding a substance or solution containing sulphate ions, including sodium sulphate.

Conclusion

When a common ion is added to two solutes, it causes precipitation or reduces ionisation, known as the common ion effect. You’ve probably heard of Le Chatelier’s principle, which states that unless an external force acts upon a reaction, it will stay in equilibrium until the force is eliminated. At that point, it will shift to accommodate the force and re-establish equilibrium. This means that you can introduce an external force, including temperature, concentration, and pressure, to cause a reaction to change. A reaction can be shifted using the common ion effect by adding an ion common to both solutes, increasing the concentration of the ion in the solution, therefore, adjusting the reaction’s equilibrium.