Acids and Bases Question 12

\begin{array}{c c c} \text{Solution} & \text{Acid} & K_a \\ \hline 1 & \ce{CH3CO2H} & 1.75\times 10^{-5} \\ 2 & \ce{CF3CO2H} & 1.0\times 10^0 \\ \end{array}

Acid-dissociation constants of two acids are listed in the table above. A 20. mL sample of a $0.10 M$ solution of each acid is titrated to the equivalence point with 20. mL of $0.10 M \ce{NaOH}$ . Which of the following is a true statement about the pH of the solutions at the equivalence point?

Solution 1 has a higher pH at the equivalence point because

\ce{CH3CO2H}

is the stronger acid.

Solution 1 has a higher pH at the equivalence point because

\ce{CH3CO2H}

has the stronger conjugate base.

Solution 1 has a lower pH at the equivalence point because

\ce{CH3CO2H}

is the stronger acid.

Solution 1 has a lower pH at the equivalence point because

\ce{CH3CO2H}

has the stronger conjugate base.

Approach

Since the $K_a$ value for $\ce{CH3CO2H}$ is smaller than that of solution 2, $\ce{CH3CO2H}$ is the weaker acid. Weak acids typically have strong conjugate bases. Therefore, the pH will be high since the solution will be basic at the equivalence point.