A student prepares a solution by combining 100 mL of 0.30 M HNO2(aq) and 100 mL of 0.30 M KNO2(aq). — Acids and Bases Chemistry Question

STEPS:

1. Identify the nature of the solution: Combining a weak acid ($\text{HNO}_2$) with its conjugate base (provided by the soluble ionic salt $\text{KNO}_2$, which dissociates completely into $\text{K}^+$ and $\text{NO}_2^-$) creates an acidic buffer solution.
2. Understand how a buffer operates: A buffer resists changes in pH because it contains both a weak acid component to neutralize added base, and a weak conjugate base component to neutralize added acid.
3. Identify the chemical stress introduced: Adding a few drops of the strong acid $\text{HCl}(aq)$ introduces a supply of free hydrogen ions ($\text{H}^+$) into the aqueous system:
$$\text{HCl}(aq) \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)$$
4. Determine the neutralization pathway: To prevent the pH from dropping significantly, the added $\text{H}^+$ ions must be absorbed. The basic component of our buffer, the nitrite ion ($\text{NO}_2^-$), will react with these incoming $\text{H}^+$ ions to form the weak, poorly-dissociating nitrous acid:
$$\mathbf{\text{H}^+(aq) + \text{NO}_2^-(aq) \rightarrow \text{HNO}_2(aq)}$$
Because the strong acid's $\text{H}^+$ ions are bound up into weak acid molecules, the overall concentration of free hydronium ions in solution is kept nearly constant, which stabilizes the pH. This identifies Option D as the correct representation.

*

WHY_OTHERS_WRONG:

• Option A is incorrect: Both $\text{K}^+$ and $\text{Cl}^-$ are spectator ions in this buffer system. Potassium chloride ($\text{KCl}$) is highly soluble in water and will not precipitate as a solid under these dilute conditions, nor does this physical dissolution process have any effect on the hydronium ion concentration.
• Option B is incorrect: This reaction represents the forward ionization (dissociation) of the weak acid $\text{HNO}_2$. Adding $\text{HCl}$ actually increases the concentration of $\text{H}^+$ in the system, which would shift this dissociation equilibrium to the *left* (by Le Chatelier's principle) rather than promoting the forward ionization.
• Option C is incorrect: Although this is the net ionic equation for a classic strong acid-strong base neutralization, there is no significant concentration of free hydroxide ions ($\text{OH}^-$) present in an acidic buffer of $\text{HNO}_2/\text{KNO}_2$. The task of neutralizing the added acid falls directly on the weak conjugate base $\text{NO}_2^-$.