Redox reaction class 11 notes: Electron Transfer in Chemistry
ElectrochemistryExploring oxidation-reduction reactions that power batteries, biological systems, and industrial processes
What Are Redox Reactions?
Redox (reduction-oxidation) reactions involve the transfer of electrons between chemical species. These reactions are fundamental to energy production, corrosion, metabolism, and countless industrial processes.
Oxidation
- Loss of electrons
- Increase in oxidation number
- Example: Zn → Zn²⁺ + 2e⁻
Reduction
- Gain of electrons
- Decrease in oxidation number
- Example: Cu²⁺ + 2e⁻ → Cu
Remembering Redox
OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons)
LEO GER: Lose Electrons Oxidation, Gain Electrons Reduction
Oxidation Numbers: The Electron Accounting System
Rules for Assigning Oxidation Numbers
- Free elements: 0 (e.g., Na, O₂)
- Monatomic ions: Equal to charge (Na⁺ = +1)
- Oxygen: Usually -2 (except peroxides: -1)
- Hydrogen: +1 with nonmetals, -1 with metals
- Sum equals charge of compound/ion
Example Calculation
KMnO₄: K(+1) + Mn(x) + 4O(-2) = 0
1 + x – 8 = 0 → x = +7
Practice Problems
Determine Mn oxidation number in:
K₂MnO₄
+6 (2(+1) + x + 4(-2) = 0 → x = +6)
Find S oxidation number in:
SO₄²⁻
+6 (x + 4(-2) = -2 → x = +6)
Balancing Redox Reactions
Example: Acidic Solution
MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺
- Oxidation: Fe²⁺ → Fe³⁺ + e⁻
- Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
- Combine: 5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
Example: Basic Solution
ClO⁻ + CrO₂⁻ → Cl⁻ + CrO₄²⁻
- Oxidation: CrO₂⁻ + 4OH⁻ → CrO₄²⁻ + 2H₂O + 3e⁻
- Reduction: ClO⁻ + H₂O + 2e⁻ → Cl⁻ + 2OH⁻
- Combine: 3ClO⁻ + 2CrO₂⁻ + 2OH⁻ → 3Cl⁻ + 2CrO₄²⁻ + H₂O
Electrochemical Cells
Galvanic (Voltaic) Cells
- Convert chemical energy → electrical energy
- Spontaneous redox reactions
- Anode (oxidation) and cathode (reduction)
- Examples: Batteries, fuel cells
Zn-Cu Galvanic Cell:
Zn|Zn²⁺(1M)||Cu²⁺(1M)|Cu
E°cell = +1.10 V
Electrolytic Cells
- Convert electrical energy → chemical energy
- Non-spontaneous reactions (require power)
- Used for electroplating, metal refining
- Example: Water electrolysis
Water Electrolysis:
2H₂O → 2H₂ + O₂
E° = -1.23 V (requires energy input)
Standard reduction potentials for common half-reactions (25°C, 1M, 1atm)
Real-World Applications
Energy Storage
- Lithium-ion batteries
- Lead-acid car batteries
- Fuel cells
Metallurgy
- Aluminum production (Hall-Héroult)
- Metal refining (copper, zinc)
- Electroplating
Biological Systems
- Cellular respiration
- Photosynthesis
- Nerve impulse transmission
Redox Reaction Calculator
Conclusion
Redox reactions form the backbone of electrochemical technology and biological energy systems. From the batteries powering our devices to the metabolic processes sustaining life, electron transfer reactions are ubiquitous in nature and technology. Mastering redox concepts enables innovations in energy storage, materials science, and environmental remediation.