Problem 152 - A Complete Resource Book in Chemistry for JEE Main

Question

When \(\mathrm{NH}_{4} \mathrm{Cl}\) is added to an aqueous solution of \(\mathrm{NH}_{4} \mathrm{OH}\), (a) Conc. of \(\left[\mathrm{OH}^{-}\right]\)ions decreases. (b) Conc. of \(\left[\mathrm{OH}^{-}\right]\)ions increases. (c) Conc. of \(\left[\mathrm{NH}_{4}^{+}\right]\)ions as well as conc. \(\left[\mathrm{OH}^{-}\right]\)ions increase. (d) Conc. of \(\left[\mathrm{NH}_{4}^{+}\right]\)ions decreases.

Step-by-Step Solution

Answer

Adding \( \mathrm{NH}_{4} \mathrm{Cl} \) decreases the \( \left[\mathrm{OH}^{-}\right] \) concentration.

Step 1: Understanding Initial Conditions

Initially, we have an aqueous solution of \( \mathrm{NH}_{4} \mathrm{OH} \). This is a weak base, which means it partially dissociates in water to produce \( \mathrm{NH}_{4}^{+} \) and \( \mathrm{OH}^{-} \) ions.

Step 2: Adding Ammonium Chloride

When \( \mathrm{NH}_{4} \mathrm{Cl} \) is added to the solution, it dissociates completely in water to give \( \mathrm{NH}_{4}^{+} \) and \( \mathrm{Cl}^{-} \) ions. The increase in \( \mathrm{NH}_{4}^{+} \) ions is due to this dissociation.

Step 3: Le Chatelier's Principle Application

According to Le Chatelier’s Principle, increasing \( \mathrm{NH}_{4}^{+} \) ions will shift the equilibrium of the dissociation of \( \mathrm{NH}_{4} \mathrm{OH} \) towards the left (towards the undissociated \( \mathrm{NH}_{4} \mathrm{OH} \)), thus decreasing the concentration of \( \mathrm{OH}^{-} \) ions.

Step 4: Conclusion Deduction

The addition of \( \mathrm{NH}_{4} \mathrm{Cl} \) decreases the concentration of \( \mathrm{OH}^{-} \) ions due to common ion effect and increases the concentration of \( \mathrm{NH}_{4}^{+} \) ions. Therefore, option (a) is correct.

Key Concepts

Common Ion Effect

The common ion effect refers to the shift in equilibrium that occurs when a solution already containing an ion from a weak electrolyte is supplemented with an external source of the same ion. In our case, the solution contains \( \mathrm{NH}_4\mathrm{OH} \) which dissociates partially to give \( \mathrm{NH}_4^+ \) and \( \mathrm{OH}^- \) ions. When \( \mathrm{NH}_4\mathrm{Cl} \) is added, it introduces more \( \mathrm{NH}_4^+ \) ions into the solution.
This injection of additional \( \mathrm{NH}_4^+ \) ions from \( \mathrm{NH}_4\mathrm{Cl} \) exerts pressure on the equilibrium state of the original \( \mathrm{NH}_4\mathrm{OH} \) dissociation.
The extra \( \mathrm{NH}_4^+ \) ions cause the equilibrium to shift towards the left, forming undissociated \( \mathrm{NH}_4\mathrm{OH} \) and reducing the concentration of \( \mathrm{OH}^- \) ions in the solution.

Weak Bases

Weak bases are substances that do not fully dissociate in water. Instead, they reach an equilibrium state that allows only a partial release of ions.
\( \mathrm{NH}_4\mathrm{OH} \) serves as a classic example of a weak base. It partially dissociates into \( \mathrm{NH}_4^+ \) and \( \mathrm{OH}^-\) ions.
Unlike strong bases that dissociate completely, weak bases exist in a dynamic balance between their undissociated form and their ions.

• Partial Dissociation: \( \mathrm{NH}_4\mathrm{OH} \leftrightarrows \mathrm{NH}_4^+ + \mathrm{OH}^- \)
• Equilibrium State: Adjusts based on changes to the system like addition of more ions.

When extra ions of \( \mathrm{NH}_4^+ \) are added via \( \mathrm{NH}_4\mathrm{Cl} \), the equilibrium position shifts, reinforcing the presence of undissociated molecule and limiting the presence of \( \mathrm{OH}^-\) ions.

Equilibrium Shift

An equilibrium shift occurs when a change in the conditions affects the balance of a chemical system. According to Le Chatelier's Principle, the system will adjust itself to counteract the imposed change.
In our example, the equilibrium of the weak base \( \mathrm{NH}_4\mathrm{OH} \) was disturbed by adding \( \mathrm{NH}_4\mathrm{Cl} \).
This extra \( \mathrm{NH}_4^+ \) from \( \mathrm{NH}_4\mathrm{Cl} \) pushes the reaction to the left, reducing the concentration of \( \mathrm{OH}^- \) ions, which is an outcome of the equilibrium's attempt to restore balance. The equilibrium reaction can be expressed as:

• \( \mathrm{NH}_4\mathrm{OH} \leftrightarrows \mathrm{NH}_4^+ + \mathrm{OH}^- \)
• Shift: Caused by increased \( \mathrm{NH}_4^+ \), moves left.

This shift results in fewer \( \mathrm{OH}^- \) ions, illustrating how equilibrium naturally adjusts to changes, ensuring that even minor disturbances can significantly influence outcomes in chemical reactions.