Common Ion Effect and Its Application for JEE

The common ion effect defines the effect on equilibrium that happens with the addition of a common ion (an ion that is already present in the solution) into a solution.

When in an electrolyte, another electrolyte is added that contains ions which are already present in the first electrolyte then, the common ion effect suppresses the ionisation of the first electrolyte. In this topic, we’ll study the common ion effect, its application and its effect on solubility and the pH of the electrolyte.

Common Ion Effect and Solubility

The common ion effect can be stated as follows: in a solution with several species associating with each other via a chemical equilibrium process, increasing the concentration of one of the ions dissociated in the solution by adding another species containing the same ion will increase the degree of ion association.

The addition of a common ion results in a decrease in solubility, this is because when a common ion is added, the reaction shifts towards the left to reduce the product formation. This happens in accordance with Le Chatlier’s principle which states that if an equilibrium is disrupted, the reaction will shift in order to restore the equilibrium. This shift of equilibrium to the left causes precipitation.

On passing gaseous hydrogen chloride through a sodium chloride solution, the excess chloride ions present in the solution cause the NaCl to precipitate, this occurs due to the common ion effect.

$\begin{align} &\mathrm{NaCl} \rightarrow \mathrm{Na}^{+}+\mathrm{Cl}^{-} \\ &\mathrm{HCl} \rightarrow \mathrm{H}^{+}+\mathrm{Cl}^{-} \end{align}$

Solubility Product and Common Ion Effect

The solubility product (Ksp) is the product of the concentrations of the ions, with each concentration being raised to a power corresponding to the ion's coefficient in the balanced equation for solubility equilibrium.

For reaction $\mathrm{A}_{\mathrm{a}} \mathrm{B}_{\mathrm{b}}(\mathrm{s}) \rightleftharpoons \mathrm{aA}^{b+}(\mathrm{aq})+\mathrm{bB}^{a-}(\mathrm{aq})$

$\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{A}^{b+}\right]^\mathrm{a}\left[\mathrm{B}^{a-}\right]^\mathrm{b}$

Here, A+b  and B-a  are cation and anions, while a and b are coefficients of A+b  and B-a 

Let’s understand this better with a question

Ques: Calculate the solubility product of barium sulphate if its solubility is $1.2 \times 10^{-5} \mathrm{~mol} ~\mathrm{dm}^{-3}$

Ans: 1 mole Barium sulphate dissociates to give 1 mole barium and 1 mole sulphate ions

$\mathrm{BaSO}_{4}(\mathrm{~s}) \rightleftharpoons \mathrm{Ba}^{2+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})$

Therefore,

$ \begin{align} &{\left[\mathrm{Ba}^{2+}\right]=1.2 \times 10^{-5}} \\ &{\left[\mathrm{SO}_{4}^{2-}\right]=1.2 \times 10^{-5}} \end{align}$

Let’s use solubility product formula

$\begin{align} &\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Ba}^{2+}\right]\left[\mathrm{SO}_{4}^{2-}\right] \\ &=\left(1.2 \times 10^{-5}\right)\left(1.2 \times 10^{-5}\right)=1.44 \times 10^{-10} \mathrm{~mol}^{2} \mathrm{dm}^{-6} \end{align}$

The common ion effect doesn’t change the solubility product (Ksp). This is because (Ksp) is a constant related to the free energy difference between the reactants and the products and it stays constant as long as the temperature does not change.

$\Delta \mathrm{G}^{\circ}=-2.303 \mathrm{RT} \log \mathrm{K}_{\mathrm{sp}}$

Effect on pH

The pH of a buffer solution changes when the conjugate ion of a buffer solution (solution comprising a weak base and its conjugate acid, or weak acid and its conjugate base) is introduced to it due to the common ion effect.

According to Le Chatelier's principle, sodium acetate's acetate ions aid in the suppression of acetic acid ionisation, pushing the equilibrium to the left. Because acetic acid dissociation is reduced, the pH of the solution increases.

NaCH3COO- → Na+(aq) + CH3COO-(aq)

CH3COOH ⇄ H+(aq) + CH3COO-(aq)

This results in higher pH (less acidic) of the common ion solution comprising acetic acid and sodium acetate as compared to an acetic acid solution.

Application of Common Ion Effect

• The common ion effect can be utilised to acquire drinking water from chalk or limestone-containing groundwater. The hardness of the water is reduced by adding sodium carbonate to it.

• The common ion effect can be seen in the salting-out process, which is utilised in the making of soaps because soaps are sodium salts of carboxylic acids with a long aliphatic chain (fatty acids). The soaps are precipitated out by reducing the solubility of the soap solution with sodium chloride.

• For gravimetric measurement, the common ion effect is employed to completely precipitate one of the ions as a sparingly soluble salt with a very low solubility product value. For example, Silver ions get precipitated as silver chloride, Barium ions as Barium sulphate, and Ferric ions as Ferric chloride or Ferric sulphate.

• The common ion effect is employed not only for quantitative analysis but also for the purification of substances. This concept can be used to purify sodium chloride with impurities like sodium sulphate and magnesium sulphate.

Conclusion

The common ion effect defines the effect on equilibrium on the addition of a common ion (an ion that is already present in the solution) into a solution. Addition of a common ion results in shifting of equilibrium towards the left resulting in a decrease in solubility and precipitation. However, it does not change the solubility product as it stays constant at a constant temperature. The common ion effect is employed in the purification of substances, treatment of water, salting out in soap formation, and quantitative analysis.