Revision Notes: Redox Reactions | Chemistry Class 11 - NEET PDF Download

Introduction

Redox reactions involve simultaneous oxidation and reduction processes, playing a key role in chemistry. They are essential in physical and biological phenomena, with applications in pharmaceuticals, industry, metallurgy, agriculture, and environmental issues like the Hydrogen Economy and Ozone Hole.

Classical Idea of Redox Reactions

• Oxidation: Originally, addition of oxygen (e.g., 2Mg(s) + O₂(g) → 2MgO(s)), later expanded to removal of hydrogen (e.g., 2H₂S(g) + O₂(g) → 2S(s) + 2H₂O(l)) or electropositive elements.
• Reduction: Initially, removal of oxygen (e.g., 2HgO(s) → 2Hg(l) + O₂(g)), later included addition of hydrogen (e.g., CH₂=CH₂(g) + H₂(g) → H₃C-CH₃(g)) or electropositive elements.
• Key Insight: Oxidation and reduction occur together, forming "redox" reactions.

Electron Transfer Perspective

• Oxidation: Loss of electrons (e.g., 2Na(s) → 2Na⁺(g) + 2e^-).
• Reduction: Gain of electrons (e.g., Cl₂(g) + 2e^- → 2Cl^-(g)).
• Half Reactions: Redox splits into oxidation and reduction halves (e.g., 2Na(s) + Cl₂(g) → 2NaCl(s)).
• Agents: Oxidizing agent accepts electrons (e.g., Cl₂), reducing agent donates electrons (e.g., Na).
• Competitive Reactions: Zn reduces Cu²⁺ (Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)), Cu reduces Ag⁺, showing order Zn > Cu > Ag.

Oxidation Number

• Definition: Charge assigned assuming electron transfer to the more electronegative atom (e.g., H₂O: H = +1, O = -2).
• Rules:
  • Elements in free state: 0 (e.g., H₂, O₂).
  • Monoatomic ions: Charge equals oxidation number (e.g., Na⁺ = +1).
  • Oxygen: -2 (except -1 in peroxides like H₂O₂, +2 in OF₂).
  • Hydrogen: +1 (except -1 in metal hydrides like NaH).
  • Fluorine: -1 always.
  • Sum in compounds = 0, in ions = charge (e.g., CO₃^2- = -2).
• Stock Notation: Roman numerals for oxidation state (e.g., Fe(II)O, Fe(III)₂O₃).
• Redox Identification: Increase in oxidation number = oxidation, decrease = reduction (e.g., 2Cu₂O(s) + Cu₂S(s) → 6Cu(s) + SO₂(g)).

Types of Redox Reactions

• Combination: A + B → C, elements combine (e.g., C(s) + O₂(g) → CO₂(g)).
• Decomposition: AB → A + B, compound breaks (e.g., 2H₂O(l) → 2H₂(g) + O₂(g)).
• Displacement: X + YZ → XZ + Y (e.g., Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)).
• Disproportionation: Element oxidizes and reduces (e.g., 2Cu⁺(aq) → Cu(s) + Cu²⁺(aq)).

Balancing Redox Reactions

• Oxidation Number Method:
  • Write formulas, assign oxidation numbers.
  • Equalize oxidation number changes.
  • Add H⁺/OH^- and H₂O to balance charges and atoms.
  • Example: Cr₂O₇^2- + 3SO₃^2- + 8H⁺ → 2Cr³⁺ + 3SO₄^2- + 4H₂O.
• Half Reaction Method:
  • Split into oxidation and reduction halves.
  • Balance atoms, add H₂O/H⁺, then electrons.
  • Equalize electrons, combine halves.
  • Example: 6Fe²⁺ + Cr₂O₇^2- + 14H⁺ → 6Fe³⁺ + 2Cr³⁺ + 7H₂O.

Redox Titrations

• Self-Indicator: MnO₄^- turns pink at endpoint.
• External Indicator: Cr₂O₇^2- with diphenylamine (blue).
• Iodine Method: Cu²⁺ + I^- → I₂, titrated with S₂O₃^2-, starch turns blue.

Electrode Processes

• Daniell Cell: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), electrons flow via wire, ions via salt bridge.
• Redox Couple: Oxidized/reduced forms (e.g., Zn²⁺/Zn).
• Electrode Potential: Tendency to gain/lose electrons, standard E° at 298 K, 1 atm, 1 M (e.g., H⁺/H₂ = 0 V).
• Table - Standard Electrode Potentials:

Summary

Redox reactions involve oxidation (electron loss) and reduction (electron gain), identified via classical (O/H changes), electron transfer, or oxidation number methods. They are classified as combination, decomposition, displacement, or disproportionation. Balancing uses oxidation number or half-reaction methods. Applications include titrations and electrochemical cells (e.g., Daniell cell), with electrode potentials indicating reactivity.