Lecture: Redox Reactions

Introduction

• Chemistry involves the transformation of matter through various reactions.
• Redox reactions are critical in various fields including pharmaceuticals, industry, and agriculture.
• Applications of redox reactions include:
  • Energy production from fuels
  • Electrochemical processes for metal extraction
  • Corrosion of metals
  • Environmental issues like Hydrogen Economy and Ozone Hole

Classical Idea of Redox Reactions

• Oxidation: Originally meant the addition of oxygen to a substance.
  • Examples: 2 Mg + O₂ → 2 MgO, S + O₂ → SO₂
• Reduction: Initially defined as the removal of oxygen.
  • Examples: 2 HgO → 2 Hg + O₂
• Oxidation also includes removal of hydrogen or electropositive elements.
• Redox reactions: Simultaneous occurrence of oxidation and reduction.

Redox Reactions in Terms of Electron Transfer

• Oxidation: Loss of electrons
• Reduction: Gain of electrons
• Oxidising Agent: Acceptor of electrons
• Reducing Agent: Donor of electrons
• Example: 2Na + Cl₂ → 2NaCl

Competitive Electron Transfer Reactions

• Experiment: Zinc in copper nitrate leads to displacement reaction.
• Zinc oxidized to Zn²⁺, Copper ions reduced to metallic copper.
• Similar experiments with copper and silver nitrate, and cobalt with nickel sulphate.
• Activity Series: Order of metals by electron releasing tendency.

Oxidation Number

• Helps track electron shifts in reactions.
• Rules for determining oxidation numbers:
  1. Free elements have oxidation number 0.
  2. For ions, the oxidation number equals the charge.
  3. Oxygen usually has an oxidation number of -2, exceptions in peroxides and with fluorine.
  4. Hydrogen is +1, except in metal hydrides.
  5. Halogens are usually -1, except in compounds with oxygen.
  6. Sum of oxidation numbers in a compound is zero; in polyatomic ions, it equals the charge.

Types of Redox Reactions

1. Combination Reactions: A + B → C
   • Example: C + O₂ → CO₂
2. Decomposition Reactions: Breakdown of compounds
   • Example: 2H₂O → 2H₂ + O₂
3. Displacement Reactions: Metal or non-metal displacement
   • Example: Zn + CuSO₄ → ZnSO₄ + Cu
4. Disproportionation Reactions: Element is both oxidized and reduced.
   • Example: 2H₂O₂ → 2H₂O + O₂

Balancing Redox Reactions

• Oxidation Number Method: Compare changes in oxidation numbers.
• Half Reaction Method: Split into oxidation and reduction half-reactions.

Redox Reactions in Titrations

• Used to determine the concentration of oxidants/reductants.
• Indicators may show color change to signal end points.

Redox Reactions and Electrode Processes

• Example: Daniell Cell
• Electrode Potential: Potential difference when species undergo redox reactions.
• Standard Electrode Potentials (E°) help determine the strength of reducing/oxidizing agents.

Summary

• Redox reactions involve simultaneous oxidation and reduction.
• Conceptualized through classical, electronic, and oxidation number views.
• Applications across various fields, crucial for understanding chemical processes.

Exercises

1. Assign oxidation numbers to elements in given species.
2. Justify given reactions as redox reactions.
3. Balance given redox reactions using methods discussed.
4. Predict product feasibility using standard electrode potentials.