Chemical Equilibrium

Acid-Base Titrations Involving Weak Acids and Bases

Acid-base titrations are a fundamental analytical technique in chemistry used to determine the concentration of an unknown acid or base solution. When weak acids or bases are involved, the process is influenced by chemical equilibrium, resulting in unique titration curves and calculations compared to strong acid-strong base titrations.

• Titration: A laboratory method where a solution of known concentration (the titrant) is gradually added to a solution of unknown concentration until the reaction reaches completion.

• Endpoint: The pH at which the acid-base indicator changes color, signaling that the titration is complete.

• Equivalence Point: The point at which the amount (in moles) of acid equals the amount of base, according to the reaction stoichiometry.

• Indicator: A dye that changes color at a specific pH range, used to visually signal the endpoint of a titration.

Key Concepts in Weak Acid-Base Titrations

• Weak Acids and Bases: These do not fully dissociate in water, resulting in an equilibrium between the undissociated and dissociated forms.

• pH Measurement: The pH measures the concentration of hydronium ions (), but for weak acids, the total acid concentration includes both dissociated and undissociated forms.

• Titration Curve: A plot of pH versus volume of titrant added, which shows characteristic regions depending on the strength of the acid and base involved.

Titration of a Weak Acid with a Strong Base

When titrating a weak acid (e.g., acetic acid, CH3COOH) with a strong base (e.g., NaOH), the titration curve differs significantly from that of a strong acid-strong base titration. The process can be divided into four key regions:

1. Initial pH: The starting pH is higher than that of a strong acid because the weak acid does not fully dissociate. The equilibrium expression for a weak acid is: The acid dissociation constant () is:

2. Acidic Region: As the strong base is added, the undissociated acid reacts with the base. The pH rises gradually. The point where (i.e., ) is called the half-equivalence point.

   • At the half-equivalence point:

3. Equivalence Point: The amount of base added equals the amount of acid originally present. For a weak acid-strong base titration, the pH at the equivalence point is greater than 7 due to the formation of a weak conjugate base in solution.

4. Alkaline Region: After the equivalence point, excess strong base is present, and the pH rises sharply.

Example: Titration of Acetic Acid with Sodium Hydroxide

• 25.00 mL of 0.200 M CH3COOH is titrated with 0.200 M NaOH.

• The titration curve will show a gradual increase in pH, a buffer region around the half-equivalence point, and a sharp rise after the equivalence point.

Comparison: Weak Acid-Strong Base vs. Strong Acid-Strong Base Titrations

• Initial pH: Higher for weak acids due to incomplete dissociation.

• Equivalence Point pH: Greater than 7 for weak acid-strong base; exactly 7 for strong acid-strong base.

• Curve Shape: The weak acid-strong base titration curve has a more gradual slope before the equivalence point and a less steep rise at the equivalence point.

Titration of a Weak Base with a Strong Acid

The process is analogous but opposite to the weak acid-strong base titration. For example, titrating 25.00 mL of 0.200 M ammonia (NH3) with 0.200 M HCl:

• The initial pH is basic due to the weak base.

• The pH decreases as acid is added, with a buffer region around the half-equivalence point.

• The equivalence point pH is less than 7 due to the formation of a weak conjugate acid.

Summary Table: Key Features of Titration Curves

Additional info: In practice, the choice of indicator is important and should match the expected pH at the equivalence point for accurate endpoint detection. Calculations often involve the Henderson-Hasselbalch equation for buffer regions: