Question

Two solutions containing weak acids were titrated with 0.100 M NaOH. If the first solution contained 100 ml of 0.100 M acetic acid (Ka= 1.8 × 10-5) and the second solution contained 100 mL of 0.100 M boric acid (Ka= 5.4 × 10-10), which of the following will be true.

(A) The pH in the buffering region will be lower for the second solution.

(B) The pH at the equivalence point will be higher for the second solution.

(C) The second solution will require more HCl to reach the equivalence point.

(D) The first solution will require more HCl to reach the equivalence point.

Answer 1

The correct statement is option (B): the pH at the equivalence point will be higher for the solution containing 0.100 M boric acid, because boric acid has a higher pKa value than acetic acid, resulting in a more basic pH at the equivalence point. Option B is correct.

Step-by-step explanation:

You are asking which of the given statements will be true when comparing the titration of 0.100 M acetic acid with 0.100 M NaOH and the titration of 0.100 M boric acid with 0.100 M NaOH. In looking at the properties of weak acids during titration with a strong base such as NaOH, the pH at various points in the titration process is affected by the pKa of acid, which reflects its acid strength.

The buffering region of the titration curve is typically around the pKa of the acid, and because boric acid has a higher pKa than acetic acid, the pH in the buffering region will be higher for boric acid, making option (A) incorrect. Options (C) and (D) are incorrect because the volumes of the acids are the same and the molarities of the acids and NaOH are also the same, meaning the volume of NaOH needed to reach the equivalence point will be equivalent for both solutions.

Therefore, option (B) is correct: the pH at the equivalence point will be higher for the second solution containing boric acid, since boric acid has a much higher pKa value compared to acetic acid, and a weaker acid results in a more basic pH at the equivalence point.