Lesson Plan

• Describe oxidation and reduction as electron loss and gain.
• Define a reducing agent and explain why it is oxidised.
• Identify the reducing agent in a redox reaction by tracking oxidation‑number changes.
• Write the oxidation half‑reaction for a given reducing agent.
• Apply the half‑reaction method to balance simple redox equations.

• Projector and screen
• Whiteboard and markers
• Printed worksheet with oxidation‑number tables
• Reaction cards (e.g., Zn + CuSO₄, Fe + CuSO₄)
• Calculator
• Diagram of electron flow (handout)

Begin with a quick demonstration of a rusting nail to spark curiosity about why metals change. Ask students what they already know about electrons moving in chemical reactions. Explain that today they will be able to define a reducing agent, spot it in a reaction, and write its oxidation half‑reaction.

1. Do‑now (5'): Short quiz on assigning oxidation numbers.
2. Mini‑lecture (10'): Definitions of oxidation, reduction, oxidising and reducing agents; illustrate with Zn → Zn²⁺ + 2e⁻.
3. Guided practice (15'): Whole‑class walk‑through of Zn + CuSO₄ reaction, identify Zn as the reducing agent.
4. Pair activity (10'): Students analyse Fe + CuSO₄, determine the reducing agent and write its oxidation half‑reaction.
5. Whole‑class discussion (5'): Share answers, clarify common misconceptions.
6. Exit ticket (5'): Write the definition of a reducing agent and give one real‑world example.

Summarise that a reducing agent donates electrons and is itself oxidised, reinforcing the electron‑flow diagram. Collect exit tickets and remind students to complete the homework worksheet, which asks them to balance two redox equations using the half‑reaction method.