Experiment Overview

These titrations are used to determine the concentration of weak acids or weak bases by reacting them with standardized strong titrants. These experiments also demonstrate how equilibrium and hydrolysis affect the pH at the equivalence point.

Weak Acid + Strong Base Titration

What This Experiment Demonstrates

• Neutralization of a weak acid by a strong base
• Effect of conjugate base hydrolysis on endpoint pH

Apparatus Required

• Burette
• Pipette
• Conical flask
• Volumetric flask
• White tile
• Retort stand with clamp

Chemicals Required

• Standardized strong base solution
• Weak acid solution of unknown concentration
• Phenolphthalein indicator
• Distilled water

Theory

A weak acid (\(\ce{HA}\)) reacts with hydroxide ions (\(\ce{OH^-}\)) to form its conjugate base (\(\ce{A^-}\)) and water.

\[\ce{CH3COOH + OH^- \rightarrow CH3COO^- + H2O}\]

At the equivalence point, the solution is effectively the same as the one obtained by dissolving a salt of the ion \(A^-\) in water. Because we're talking about a weak acid, the solution will contain both \(A^-\) and \(HA\). The conjugate base hydrolyzes with water to produce hydroxide ions, making the solution slightly basic.

\[\ce{CH3COO^- + H2O \rightleftharpoons CH3COOH + OH^-}\]

Since the equivalence point lies in the basic region, phenolphthalein is used as the indicator because it changes color around pH 8–10.

Method

1. Using a pipette, transfer a known volume of the weak acid solution into a conical flask.
2. Add 2–3 drops of phenolphthalein indicator.
3. Fill the burette with the standardized strong base solution.
4. Titrate while continuously swirling the flask.
5. Continue until a faint permanent pink color appears.
6. Record the final burette reading and repeat to obtain concordant results.

Observations

• Initial solution: colorless
• During titration: remains colorless until near endpoint
• At endpoint: pale pink color persists for about 30 seconds

Calculations

Example: \(\ce{NaOH + CH3COOH}\)

• Balanced equation ratio: 1:1
• Moles of \(NaOH = M \times V\)
• Moles of weak acid = moles of \(NaOH\)
• Concentration of weak acid:

\[\text{M}_{\text{acid}} = \frac{\text{M}_{\text{base}} \times \text{V}_{\text{base}}}{\text{V}_{\text{acid}}}\]

We can show, using the Henderson-Hasselbach equation that, the semi-equivalence pH (the pH when half of the required base was added) is \(pH_{\text{semi}} = pK_a\)

Result

The concentration of the weak acid is determined by direct titration using a standardized strong base and phenolphthalein indicator.

Strong Acid + Weak Base Titration

What This Experiment Demonstrates

• Neutralization reaction between a strong acid and a weak base
• Effect of conjugate acid hydrolysis on endpoint pH

Apparatus Required

• Burette
• Pipette
• Conical flask
• Volumetric flask
• White tile
• Retort stand with clamp

Chemicals Required

• Standardized strong acid solution
• Weak base solution of unknown concentration
• Methyl orange indicator
• Distilled water

Theory

A weak base (\(B\)) reacts with hydronium ions from the strong acid to form its conjugate acid ((\BH^+\)).

\[\ce{NH3 + H3O^+ \rightarrow NH4^+ + H2O}\]

At the equivalence point, the solution is effectively the same as the one obtained by dissolving a salt of the ion \(BH^+\) in water. Because we're talking about a weak base, the solution will contain both \(BH^+\) and \(B\). This conjugate acid hydrolyzes to produce hydronium ions, making the solution slightly acidic.

\[\ce{NH4^+ + H2O \rightleftharpoons NH3 + H3O^+}\]

Because the equivalence point is acidic, methyl orange is used as the indicator. Phenolphthalein would give an inaccurate endpoint because its transition range is too basic.

Method

1. Using a pipette, transfer a known volume of the weak base solution into a conical flask.
2. Add 2–3 drops of methyl orange indicator.
3. Fill the burette with the standardized strong acid solution.
4. Titrate while continuously swirling the flask.
5. Continue until the color changes from yellow to faint orange/pink.
6. Record the final burette reading and repeat to obtain concordant results.

Observations

• Initial solution: yellow
• At endpoint: sharp change to pale orange/pink
• No permanent color change before the endpoint

Calculations

Example: \(\ce{HCl + NH3}\)

• Balanced equation ratio: 1:1
• Moles of \(\ce{HCl} = M \times V\)
• Moles of weak base = moles of \(\ce{HCl}\)
• Concentration of weak base:

\[\text{M}_{\text{base}} = \frac{\text{M}_{\text{acid}} \times \text{V}_{\text{acid}}}{\text{V}_{\text{base}}}\]

We can show, using the Henderson-Hasselbach equation that, the semi-equivalence pH (the pH when half of the required base was added) is \(pH_{\text{semi}} = pK_w - pK_b\) or \(pOH = pK_b\) or \(pH = pKa\) of the conjugate acid.

Result

The concentration of the weak base is determined by direct titration with a standardized strong acid using methyl orange indicator.