Titration and Buffers (5A) Practice Test

A 50.0 mL sample of 0.100 M acetic acid ($\mathrm{HA}$) is titrated with 0.100 M NaOH. The reaction is $\mathrm{HA + OH^- \rightarrow A^- + H_2O}$, and $pK_a(\mathrm{HA})=4.76$ (25°C). At the equivalence point, essentially all initial $\mathrm{HA}$ has been converted to $\mathrm{A^-}$, and the pH is determined primarily by base hydrolysis: $\mathrm{A^- + H_2O \rightleftharpoons HA + OH^-}$. For acetate, $K_b=K_w/K_a$ with $K_w=1.0\times10^{-14}$.

Which statement best explains why the pH at the equivalence point is expected to be greater than 7?

Constants provided: $pK_a=4.76$; $K_w=1.0\times10^{-14}$.

A weak base B is titrated with strong acid HCl at 25°C. The reaction is:

$\mathrm{B + H^+ \rightarrow BH^+}$

A 50.0 mL sample of 0.100 M B is titrated with 0.100 M HCl. The conjugate acid has $pK_a(\mathrm{BH^+}) = 8.30$. Assume ideal behavior.

At a point during the titration where both B and $\mathrm{BH^+}$ are present in substantial amounts, the solution shows minimal pH change upon addition of a small amount of HCl.

Which observation is most consistent with the solution being in its maximum buffering region during this titration?