3.3
Titration of a Weak Acid with a Strong Base
In titrating a weak acid with a strong base, different calculation methods are applied at various stages. Initially, the pH of a weak acid like acetic…
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 Ka 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 pKa 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 Kb 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.
Q1: How do you calculate the pH of a weak acid before adding any 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.
Q2: What equation is used to find pH when a buffer forms during titration?
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.
Q3: Why does pH equal pKa at the halfway point of a weak acid titration?
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.
Q4: How is pH calculated at the equivalence point of a weak acid-strong base titration?
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.
Q5: What determines pH when excess strong base is added beyond the equivalence point?
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.
Q6: How do moles of acid and conjugate base change during weak acid titration?
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.
Q7: Why is the equivalence point pH basic in weak acid-strong base titrations?
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.