Chemistry

Problem

A student prepares an aqueous solution containing

0.100\,\text{mol}

of ammonia

(NH_3)

in a total volume of

1.00\,\text{L}

. The student then adds

0.030\,\text{mol}

of hydrochloric acid

(HCl)

directly to the solution and allows the system to reach equilibrium at

25^\circ\text{C}

. No additional substances are present. (a) Explain, in terms of the chemical species present before and after the addition of the strong acid, why the resulting mixture is classified as a buffer solution. (b) Describe qualitatively how the relative amounts of

NH_3

and

NH_4^+

control the pH of the buffer, referencing the positions of the relevant equilibrium reactions. (c) Predict and justify the effect of diluting the final mixture to

10.0\,\text{L}

on (i) the numerical value of the pH and (ii) the overall buffer capacity toward added strong acid.

Intuitive explanation

This is a buffer chemistry problem where a weak base reacts with a strong acid. When you add HCl to ammonia solution, the acid neutralizes some of the base in a 1:1 reaction, leaving you with a mixture of unreacted ammonia and its conjugate acid (ammonium ion). The key insight is recognizing that after the acid–base reaction completes, you'll have both $NH_3$ and $NH_4^+$ present—a classic buffer system—and you can use the Henderson-Hasselbalch equation or the $K_b$ expression to find the pH or equilibrium concentrations.

Intuitive explanation

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