Introduction
Redox titrations involve oxidation-reduction reactions between two solutions — one acting as the titrant and the other as the analyte. These titrations are essential for determining the concentration of unknown substances and are commonly tested in NECTA practical and theory exams.

🔹 Common Redox Titrants

• Potassium permanganate (KMnO₄) – self-indicating, purple

• Potassium dichromate (K₂Cr₂O₇) – requires external indicator

• Iodine (I₂) – usually titrated with thiosulfate

• Potassium iodate (KIO₃) – generates iodine in acidic medium

🔹 Common Analytes

• Iron(II) salts (e.g., FeSO₄)

• Oxalic acid (H₂C₂O₄)

• Sodium thiosulfate (Na₂S₂O₃)

• Hydrogen peroxide (H₂O₂)

🔹 Redox Indicators

These are substances that change color due to changes in oxidation state.

🔹 Basic Setup for a Redox Titration

1. Titrant (known concentration) in burette

2. Analyte (unknown concentration) in conical flask

3. Add indicator if not self-indicating

4. Titrate until the end point (color change)

🧪 Example 1: KMnO₄ vs Fe²⁺

• Burette: KMnO₄ (purple)

• Flask: Fe²⁺ in H₂SO₄

• Reaction:

  MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

• End point: Disappearance of colorless → pale pink (1st drop in excess)

🧪 Example 2: Iodine vs Thiosulfate

• Burette: Na₂S₂O₃

• Flask: Iodine (I₂) + starch indicator

• Reaction:

  I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻

• End point: Blue-black color → colorless

🧠 NECTA Tip

Always:

• Use fresh KMnO₄ (it decomposes easily)

• Acidify with H₂SO₄, not HCl (Cl⁻ interferes with oxidation)

• Rinse burette with KMnO₄ before titrating

• Record burette readings to 2 decimal places (e.g. 25.30 mL)

✅ Summary

• Redox titrations are based on electron transfer

• KMnO₄ is a common self-indicating oxidizing agent

• Starch is used for iodine titrations

• End points must be sharp and well-observed

• Equations must be balanced properly for correct stoichiometry