A solution is prepared by adding 100 mL of 1.0 M HC2H3O2(aq) to 100 mL of 1.0 M NaC2H3O2(aq) . The s — Acids and Bases Chemistry Question

STEPS:

1. Identify the chemical species in the starting mixture: The solution is prepared by mixing a weak acid, acetic acid ($\text{HC}_2\text{H}_3\text{O}_2$), with sodium acetate ($\text{NaC}_2\text{H}_3\text{O}_2$). Sodium acetate is a soluble salt that dissociates completely into sodium ions ($\text{Na}^+$) and acetate ions ($\text{C}_2\text{H}_3\text{O}_2^-$) in water. This combination of a weak acid and its conjugate base forms a buffer solution that resists changes in pH.
2. Understand the effect of adding a strong acid: When hydrochloric acid ($\text{HCl}$), a strong acid, is added to the solution, it dissociates completely in water to produce hydronium ions ($\text{H}_3\text{O}^+$) and chloride ions ($\text{Cl}^-$).
3. Determine how the buffer counteracts the added acid: To prevent the pH from decreasing significantly, the buffer system must consume the added hydronium ions ($\text{H}_3\text{O}^+$). In a buffer, the weak conjugate base is responsible for neutralizing added acid. Therefore, the acetate ions ($\text{C}_2\text{H}_3\text{O}_2^-$) present in the solution will react with the added $\text{H}_3\text{O}^+$.
4. Write the net ionic equation for the neutralization: The acetate ion acts as a Brønsted^-Lowry base, accepting a proton from the hydronium ion to form water and undissociated acetic acid molecules:
$$\text{H}_3\text{O}^+(aq) + \text{C}_2\text{H}_3\text{O}_2^-(aq) \rightarrow \text{HC}_2\text{H}_3\text{O}_2(aq) + \text{H}_2\text{O}(l)$$
Because this reaction converts a strong acid ($\text{H}_3\text{O}^+$) into a weak acid ($\text{HC}_2\text{H}_3\text{O}_2$), the concentration of free hydronium ions in solution remains nearly constant, which accounts for why the pH does not detectably change. This identifies Option C as the correct equation.

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WHY_OTHERS_WRONG:

• A is incorrect: This equation represents the neutralization of a strong acid by a strong base ($\text{OH}^-$). While this reaction is chemically correct, the buffer solution has an acidic pH of $4.73$. In an acidic solution, the concentration of free hydroxide ions ($\text{OH}^-$) is extremely low ($\approx 5.4 \times 10^{-10}\text{ M}$), meaning there is virtually no hydroxide available to neutralize the added hydronium ions.
• B is incorrect: Hydrochloric acid is a strong acid that remains completely ionized in dilute aqueous solutions. Hydronium and chloride ions do not combine in water to form gaseous $\text{HCl}$.
• D is incorrect: Acetic acid ($\text{HC}_2\text{H}_3\text{O}_2$) is an acid and cannot act as a base to neutralize added hydronium ions ($\text{H}_3\text{O}^+$) under these conditions.