Titration calculations for weak acids and strong bases involve different approaches, depending on the primary reactant.

Initially, 50 mL of 0.1 M acetic acid has a pH of 2.87, calculated with the K_a and an ICE table.

Post titration with 0.1 M NaOH, the solution forms a buffer.

Adding 10 mL of NaOH creates 0.001 moles of acetate, leaving 0.004 moles of acetic acid. The resultant pH is 4.14, calculated using the Henderson-Hasselbalch equation.

Halfway through, the pH equals pK_a due to equal acetic acid and acetate ion concentration.

At the equivalence point, the addition of 50 mL of NaOH converts all acetic acid to acetate, resulting in a pH transition to basic. Using an ICE table and K_b for acetate ions, the pH is found to be 8.72.

Any further addition of NaOH will dictate the pH, as it's a stronger base than acetate. For instance, adding 70 mL of NaOH results in a final pH of 12.22.