An acetate buffer solution is prepared by combining 50. mL of 0.20 \(M\) acetic acid, \(\ce{HC2H3O2(aq)}\), and 50. mL of 0.20 \(M\) sodium acetate, \(\ce{NaC2H3O2(aq)}\). A 5.0 mL sample of 0.10 \(M\) \(\ce{NaOH(aq)}\) is added to the buffer solution. Which of the following is a correct pairing of the acetate species present in greater concentration and of the pH of the solution after the \(\ce{NaOH(aq)}\) is added? (The p\(K_a\) of acetic acid is 4.7.)

$$ \begin{array}{c|c|c} \hspace{3em}& \text{Acetate Species} & \text{pH} \\ \hline \text{(A)} & \ce{HC2H3O2} & \lt 4.7 \\ \text{(B)} & \ce{HC2H3O2} & \gt 4.7 \\ \text{(C)} & \ce{C2H3O2-} & \lt 4.7 \\ \text{(D)} & \ce{C2H3O2-} & \gt 4.7 \\ \end{array} $$

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The strong acid and sodium acetate will dissociate completely in water. Acetic acid is a weak acid (hence its p\(Ka\)), and will dissociate slightly. The following reaction will proceed to completion since hydroxide ions will react with any free hydronium ions that dissociate from acetic acid.

$$ \ce{OH- + HC2H3O2 -> H2O + C2H3O2-} $$

There will therefore be more \(\ce{C2H3O2-}\) ions in solution. The pH will be slightly above 4.7 due to the addition of the base.