Science Class 11 - Redox Reactions Notes

Redox Reactions

Welcome to the chapter on Redox Reactions for Class 11. In this chapter, you will learn what redox reactions are, how to identify oxidation and reduction, and the importance of redox reactions in chemistry and daily life. By the end of this chapter, you will be able to balance redox equations and understand their applications.

Introduction

Redox reactions are chemical reactions in which oxidation and reduction occur simultaneously. These reactions are fundamental to many processes in chemistry, biology, and industry.

Key Concepts

• Oxidation: Loss of electrons or increase in oxidation number.
• Reduction: Gain of electrons or decrease in oxidation number.
• Redox Reaction: A reaction in which both oxidation and reduction take place.
• Oxidizing Agent: Substance that causes oxidation (itself gets reduced).
• Reducing Agent: Substance that causes reduction (itself gets oxidized).

Oxidation and Reduction

• Oxidation:
  • Loss of electrons
  • Addition of oxygen
  • Increase in oxidation number
• Reduction:
  • Gain of electrons
  • Removal of oxygen
  • Decrease in oxidation number

Examples of Redox Reactions

• Rusting of Iron:
  4Fe + 3O₂ → 2Fe₂O₃
• Reaction of Hydrogen and Oxygen:
  2H₂ + O₂ → 2H₂O
• Displacement Reaction:
  Zn + CuSO₄ → ZnSO₄ + Cu

Balancing Redox Reactions

Redox reactions can be balanced using the ion-electron method (half-reaction method) or the oxidation number method.

• Write the oxidation and reduction half-reactions separately.
• Balance atoms and charges in each half-reaction.
• Combine the half-reactions to get the balanced redox equation.

Applications of Redox Reactions

• Batteries and electrochemical cells
• Metallurgy (extraction of metals)
• Photosynthesis and respiration
• Bleaching and disinfection

Practice Questions

1. Define oxidation and reduction with examples.
2. Identify the oxidizing and reducing agents in the reaction: Zn + CuSO₄ → ZnSO₄ + Cu.
3. Balance the following redox reaction: Fe²⁺ + Cr₂O₇²⁻ + H⁺ → Fe³⁺ + Cr³⁺ + H₂O.
4. Explain the importance of redox reactions in daily life.
5. What is the difference between an oxidizing agent and a reducing agent?

Challenge Yourself

• Write the half-reactions for the reaction: 2Na + Cl₂ → 2NaCl.
• Give two examples of redox reactions in biological systems.

Did You Know?

• Redox reactions are responsible for the energy produced in our cells!
• The process of photosynthesis is a complex redox reaction.

Glossary

• Oxidation Number: A number assigned to an atom to show its degree of oxidation or reduction.
• Half-Reaction: The part of a redox reaction that involves only oxidation or only reduction.
• Electrochemical Cell: A device that generates electricity from a redox reaction.

Answers to Practice Questions

1. Oxidation: Loss of electrons (e.g., Zn → Zn²⁺ + 2e⁻).
   Reduction: Gain of electrons (e.g., Cu²⁺ + 2e⁻ → Cu).
2. Oxidizing agent: CuSO₄ (Cu²⁺), Reducing agent: Zn.
3. Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ → Fe³⁺ + 2Cr³⁺ + 7H₂O (balanced).
4. Redox reactions are important for respiration, photosynthesis, batteries, and many industrial processes.
5. An oxidizing agent gains electrons and is reduced; a reducing agent loses electrons and is oxidized.

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Understanding redox reactions is key to mastering chemistry and its real-world applications!