oxidation and reduction worksheet with answers

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04-02-2025

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This worksheet provides a comprehensive guide to understanding oxidation and reduction (redox) reactions. It's designed to help you master the fundamental concepts, practice identifying oxidizing and reducing agents, and confidently tackle redox balancing. Each question includes a detailed answer to help solidify your learning.

Section 1: Identifying Oxidation and Reduction

Instructions: For each reaction, identify what is being oxidized and what is being reduced. State which species is the oxidizing agent and which is the reducing agent.

1. 2Fe(s) + 3Cl₂(g) → 2FeCl₃(s)

Answer:

• Oxidation: Fe(s) → Fe³⁺(s) + 3e⁻ (Iron loses electrons, its oxidation state increases from 0 to +3)
• Reduction: Cl₂(g) + 2e⁻ → 2Cl⁻(s) (Chlorine gains electrons, its oxidation state decreases from 0 to -1)
• Oxidizing Agent: Cl₂ (Chlorine accepts electrons, causing the oxidation of iron.)
• Reducing Agent: Fe (Iron donates electrons, causing the reduction of chlorine.)

2. Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Answer:

• Oxidation: Zn(s) → Zn²⁺(aq) + 2e⁻ (Zinc loses electrons, oxidation state increases from 0 to +2)
• Reduction: Cu²⁺(aq) + 2e⁻ → Cu(s) (Copper gains electrons, oxidation state decreases from +2 to 0)
• Oxidizing Agent: Cu²⁺ (Copper(II) ion accepts electrons, causing the oxidation of zinc.)
• Reducing Agent: Zn (Zinc donates electrons, causing the reduction of copper(II) ion.)

3. 2Mg(s) + O₂(g) → 2MgO(s)

Answer:

• Oxidation: 2Mg(s) → 2Mg²⁺(s) + 4e⁻ (Magnesium loses electrons, oxidation state increases from 0 to +2)
• Reduction: O₂(g) + 4e⁻ → 2O²⁻(s) (Oxygen gains electrons, oxidation state decreases from 0 to -2)
• Oxidizing Agent: O₂ (Oxygen accepts electrons, causing the oxidation of magnesium.)
• Reducing Agent: Mg (Magnesium donates electrons, causing the reduction of oxygen.)

Section 2: Balancing Redox Reactions (using the half-reaction method)

Instructions: Balance the following redox reactions in acidic solution using the half-reaction method.

4. MnO₄⁻(aq) + Fe²⁺(aq) → Mn²⁺(aq) + Fe³⁺(aq)

Answer:

1. Half-reactions:

   • Oxidation: Fe²⁺(aq) → Fe³⁺(aq) + e⁻
   • Reduction: MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ → Mn²⁺(aq) + 4H₂O(l)

2. Balance electrons: Multiply the oxidation half-reaction by 5.

3. Add half-reactions: 5Fe²⁺(aq) + MnO₄⁻(aq) + 8H⁺(aq) → 5Fe³⁺(aq) + Mn²⁺(aq) + 4H₂O(l)

5. Cr₂O₇²⁻(aq) + SO₃²⁻(aq) → Cr³⁺(aq) + SO₄²⁻(aq) (in acidic solution)

Answer:

1. Half-reactions:

   • Oxidation: SO₃²⁻(aq) + H₂O(l) → SO₄²⁻(aq) + 2H⁺(aq) + 2e⁻
   • Reduction: Cr₂O₇²⁻(aq) + 14H⁺(aq) + 6e⁻ → 2Cr³⁺(aq) + 7H₂O(l)

2. Balance electrons: Multiply the oxidation half-reaction by 3.

3. Add half-reactions: 3SO₃²⁻(aq) + Cr₂O₇²⁻(aq) + 8H⁺(aq) → 3SO₄²⁻(aq) + 2Cr³⁺(aq) + 4H₂O(l)

This worksheet provides a foundational understanding of redox reactions. Remember to practice more examples to solidify your understanding. Further exploration into more complex redox reactions and their applications in various fields will enhance your knowledge.