Balancing Redox Reactions: Example
Let’s balance the reaction between permanganate (MnO₄⁻) and iodide (I⁻) ions in basic medium:
Step 1: Half-reactions
Oxidation: \( 2I^- (\text{aq}) \rightarrow I_2 (\text{s}) + 2e^- \)
Reduction: \( \text{MnO}_4^-(\text{aq}) + 2\text{H}_2\text{O} (l) + 3e^- \rightarrow \text{MnO}_2 (\text{s}) + 4\text{OH}^- (\text{aq}) \)
Step 2: Equalize electrons 🔋
Multiply oxidation by 3, reduction by 2:
\( 6I^- (\text{aq}) \rightarrow 3I_2 (\text{s}) + 6e^- \)
\( 2\text{MnO}_4^- (\text{aq}) + 4\text{H}_2\text{O} (l) + 6e^- \rightarrow 2\text{MnO}_2 (\text{s}) + 8\text{OH}^- (\text{aq}) \)
Step 3: Combine and verify ✅
Final balanced equation:
\( 6I^- (\text{aq}) + 2\text{MnO}_4^- (\text{aq}) + 4\text{H}_2\text{O} (l) \rightarrow 3I_2 (\text{s}) + 2\text{MnO}_2 (\text{s}) + 8\text{OH}^- (\text{aq}) \)
Atoms and charges balance on both sides!
Redox Titrations and Indicators
Just like pH indicators in acid-base titrations, redox titrations use special indicators:
1. Self-indicators 🌈
Some reagents change color by themselves! Example:
• Purple \( \text{MnO}_4^- \) turns colorless when reduced → pink tinge appears at endpoint (with \( Fe^{2+} \) or \( C_2O_4^{2-} \))
2. Added redox indicators 🧪
For reactions without color change (e.g., \( \text{Cr}_2\text{O}_7^{2-} \)):
• Diphenylamine indicator turns intense blue when oxidized after the endpoint
3. Starch-iodine method 🔵
Used for oxidizing agents that react with I⁻ ions (e.g., Cu²⁺):
• \( 2\text{Cu}^{2+} + 4I^- \rightarrow \text{Cu}_2\text{I}_2 (\text{s}) + \text{I}_2 \)
• Iodine (\( I_2 \)) + starch → deep blue color
• Blue color fades when \( I_2 \) reacts: \( \text{I}_2 + 2S_2O_3^{2-} \rightarrow 2I^- + S_4O_6^{2-} \)
Electrode Processes
Direct vs. indirect electron transfer:
• Direct: Zinc in CuSO₄ solution → electrons move directly from Zn to Cu²⁺ (heat released) 🔥
• Indirect: Daniell cell separates half-reactions:
Daniell cell setup ⚡
1. Anode (Zn): \( \text{Zn} \rightarrow \text{Zn}^{2+} + 2e^- \) (oxidation)
2. Cathode (Cu): \( \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu} \) (reduction)
3. Salt bridge: Completes circuit (ions flow without mixing solutions)
4. Wire: Electrons flow from Zn to Cu
Electrode potential 📊
• Standard Electrode Potential (E°): Measured at 298K, 1M concentration, 1 atm pressure
• Reference: Hydrogen electrode (\( 2\text{H}^+ + 2e^- \rightarrow \text{H}_2 \)) has E° = 0.00 V
• Negative E° → stronger reducing agent than H₂
• Positive E° → weaker reducing agent than H₂
Standard Electrode Potentials (Reduction)
| Reaction | E° (V) |
|---|---|
| \( \text{F}_2(\text{g}) + 2e^- \rightarrow 2\text{F}^- \) | +2.87 |
| \( \text{Cl}_2(\text{g}) + 2e^- \rightarrow 2\text{Cl}^- \) | +1.36 |
| \( \text{Ag}^+ + e^- \rightarrow \text{Ag}(\text{s}) \) | +0.80 |
| \( \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}(\text{s}) \) | +0.34 |
| \( 2\text{H}^+ + 2e^- \rightarrow \text{H}_2(\text{g}) \) | 0.00 |
| \( \text{Fe}^{2+} + 2e^- \rightarrow \text{Fe}(\text{s}) \) | -0.44 |
| \( \text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}(\text{s}) \) | -0.76 |
| \( \text{Al}^{3+} + 3e^- \rightarrow \text{Al}(\text{s}) \) | -1.66 |
| \( \text{Na}^+ + e^- \rightarrow \text{Na}(\text{s}) \) | -2.71 |
Important Concepts for NEET
- 🔋 Balancing redox reactions in acidic/basic media using half-reaction method
- 🧪 Redox titration indicators (self-indicator, starch-iodine, diphenylamine)
- ⚡ Daniell cell components and electron/ion flow directions
- 📉 Standard electrode potential (E°) interpretation (negative E° = stronger reducing agent)
- 🔄 Predicting redox spontaneity using E° values (higher reduction potential → preferred reduction)
