Redox Titration

Redox titration is an important analytical technique used in chemistry to determine the concentration of a substance through oxidation–reduction (redox) reactions. In such reactions, one substance undergoes oxidation by losing electrons, while another undergoes reduction by gaining electrons. These two processes occur simultaneously and form the basis of redox titrations.

Oxidation: Oxidation is the loss of electrons or increase in oxidation number of a substance.

Example: Zn → Zn²⁺ + 2e− ( Zinc loses electrons, so it undergoes oxidation.)

Reduction: Reduction is the gain of electrons or decrease in oxidation number of a substance.

Example: Cu²⁺ + 2e− → Cu (Copper ions gain electrons, so reduction occurs.)

Principle of Redox Titration

The principle of redox titration involves the quantitative measurement of this electron transfer to determine the concentration of a solution.

• Oxidation-Reduction Reactions: Redox titrations involve reactions where one substance is oxidized (loses electrons) and another is reduced (gains electrons).
• Titration Endpoint: The endpoint is determined by a slight change in the system, like a colour change or appearance of a precipitate, indicating that the reaction is complete.
• Equivalence Point: At this point, the amount of titrant added is equal to the amount of analyte present, marking the completion of the reaction.
• Choice of Indicator: Indicators are selected based on their ability to change colour near the equivalence point.
• Stoichiometry: Balanced chemical equations help calculate the amount of titrant needed for the reaction.

Redox Titration Indicators

In redox titrations, indicators are substances used to detect the endpoint of the titration. They help in identifying the completion of the reaction by producing a visible change, usually in colour.

• Redox indicators show a colour change due to change in oxidation state.
• The colour change occurs near the endpoint of the titration.
• Some indicators are external indicators, while some are internal indicators.
• In certain titrations, no separate indicator is required because the titrant acts as a self-indicator.
• The selection of a suitable indicator is important for obtaining accurate results.

Examples:

Potassium Permanganate (KMnO₄) as Self-Indicator

• KMnO₄ is purple in colour
• At the endpoint, a permanent light pink colour appears
• No external indicator is required

Procedure of Redox Titration

The titration between potassium permanganate (KMnO₄) and oxalic acid is a common redox titration. In this reaction, KMnO₄ acts as a strong oxidizing agent, while oxalic acid acts as a reducing agent. The reaction is carried out in an acidic medium and KMnO₄ acts as a self-indicator.

Step 1: Prepare a standard solution of potassium permanganate (KMnO₄) and fill it in the burette.

Step 2: Pipette out a measured volume of oxalic acid solution into a clean conical flask.

Step 3: Add dilute sulphuric acid (H₂SO₄) to the oxalic acid solution to make the medium acidic.

MnO ^-₄ ​+ 8H⁺ + 5e− → Mn ²⁺ + 4H_2​O

MnO ^-₄ + 8H⁺ + 5e− →Mn ²⁺ + 4H₂O

Step 4: Heat the conical flask gently to about 60–70°C, as the reaction is slow at room temperature.

C_2​O₄²⁻ ​→ 2CO_{2 ​}+ 2e−

C₂O₄²⁻ → 2CO₂ + 2e−

Step 5: Begin the titration by adding KMnO₄ solution slowly from the burette into the conical flask with continuous stirring.

Step 6: Observe that the pink colour of KMnO₄ disappears initially due to reduction.

Step 7: Continue adding KMnO₄ dropwise until a light pink colour starts to persist.

Step 8: Stop the titration when a permanent light pink colour persists for about 30 seconds, indicating the endpoint.

2KMnO₄ + 5H₂C2​O₄ ​+ 3H₂SO₄ ​→ K₂​SO₄ ​+ 2MnSO₄ ​+ 10CO₂ ​+ 8H₂O​

Step 9: Note the final burette reading and repeat the titration to obtain accurate results.

Related Articles:

• Applications of Titration
• Acid-Base Titration