Titration of a Weak Acid with a Strong Base

The initial pH of a weak acid is calculated using its dissociation constant (Ka) and an ICE table. For example, 50 mL of 0.1 M acetic acid has a pH of 2.87. The ICE table tracks initial concentrations, changes during dissociation, and equilibrium concentrations, allowing you to solve for hydrogen ion concentration and determine pH.

The Henderson-Hasselbalch equation calculates buffer pH during weak acid titration with strong base. When 10 mL of 0.1 M NaOH is added to acetic acid, a buffer forms with both acid and conjugate base present. The equation uses the pKa and the ratio of acetate to acetic acid concentrations to yield a pH of 4.14.

At the halfway point, the concentrations of weak acid and its conjugate base are equal. When equal molar amounts of acetic acid and acetate ion are present, the Henderson-Hasselbalch equation simplifies because the log of 1 equals zero, leaving pH equal to pKa. This occurs when half the original acid has been neutralized.

At the equivalence point, all weak acid converts to its conjugate base, creating a basic solution. An ICE table and the base dissociation constant (Kb) of the conjugate base determine pH. For acetic acid titrated with NaOH, 50 mL of base produces acetate ions, yielding a pH of 8.72 through Kb calculations.

Beyond the equivalence point, the pH is governed by the concentration of excess strong base, not the conjugate base. The strong base completely dominates the solution's acidity. Adding 70 mL of NaOH results in a final pH of 12.22, determined solely by the excess sodium hydroxide concentration.

As strong base is added, it neutralizes weak acid molecules, converting them to conjugate base. Adding 10 mL of 0.1 M NaOH creates 0.001 moles of acetate while leaving 0.004 moles of acetic acid unreacted. This changing ratio of acid to conjugate base directly affects pH calculations throughout the titration.

At the equivalence point, the conjugate base of the weak acid remains in solution and undergoes hydrolysis, accepting protons from water and producing hydroxide ions. This makes the solution basic. The acetate ion, conjugate base of acetic acid, hydrolyzes to form OH-, raising pH to 8.72 rather than the neutral pH seen in strong acid-strong base titrations.