1.1 Expanded Definitions of Oxidation and Reduction

Traditional Definitions:

• Oxidation: Gain of oxygen.
• Reduction: Loss of oxygen.

Expanded Definitions:

• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.

Memory Aid: OIL RIG

• OIL RIG stands for:
  • Oxidation is the Is Loss of electrons.
  • Reduction is the Is Gain of electrons.

1.2 Analyzing Redox Reactions Using Ionic Equations

Example: Oxidation of Copper 2Cu(s)+O₂(g)→2CuO(s)

Ionic Perspective:

• Copper (Cu):
  • Oxidation: Cu → Cu²⁺ + 2e⁻
• Oxygen (O₂):
  • Reduction: O₂ + 4e⁻ → 2O²⁻

Summary:

• Copper atoms lose electrons (are oxidized) to form Cu²⁺ ions.
• Oxygen molecules gain electrons (are reduced) to form O²⁻ ions.

1.3 Redefining Redox Reactions

• With the expanded definitions, redox reactions encompass a broader range of reactions beyond those involving oxygen, including:

• Metal Displacement Reactions: Where a more reactive metal displaces a less reactive metal from its compound.

• Example: Reaction Between Zinc and Copper(II) Sulfate
• Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

• Ionic Equation:
• Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Analysis:

• Zinc (Zn):
  • Oxidation: Zn → Zn²⁺ + 2e⁻
• Copper (Cu²⁺):
  • Reduction: Cu²⁺ + 2e⁻ → Cu

Conclusion:

• Zinc is oxidized (loses electrons).
• Copper(II) ions are reduced (gain electrons).
• This is a redox reaction as it involves both oxidation and reduction.

1.4 Metal Displacement Reactions as Redox Reactions

Definition:

• Displacement Reaction: A reaction where a more reactive metal displaces a less reactive metal from its compound.

• Example: Chlorine Displacing Iodine from Potassium Iodide
  Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)
• Ionic Equation:
  Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) +I₂(aq)

Analysis:

• Chlorine (Cl₂):
  • Reduction: Cl₂ + 2e⁻ → 2Cl⁻
• Iodide Ions (I⁻):
  • Oxidation: 2I⁻ → I₂ + 2e⁻²

Conclusion:

• Chlorine is reduced (gains electrons).
• Iodide ions are oxidized (lose electrons).
• This is a redox reaction as it involves both oxidation and reduction.

1.5 Summary of Electron Transfer in Redox Reactions

• Oxidation: Always involves the loss of electrons.
• Reduction: Always involves the gain of electrons.
• Redox Reactions: Involve simultaneous oxidation and reduction processes.

2. Oxidation and Reduction in Electrolysis and Fuel Cells

2.1 Electrolysis

Definition:

• Electrolysis: A chemical process that uses electricity to drive a non-spontaneous reaction, resulting in the decomposition of compounds.

Key Concepts:

• Electrodes:
  • Anode: Positive electrode where oxidation occurs.
  • Cathode: Negative electrode where reduction occurs.
• Ion Movement:
  • Negative Ions (Anions): Move to the anode to lose electrons (oxidation).
  • Positive Ions (Cations): Move to the cathode to gain electrons (reduction).

Example: Electrolysis of Concentrated Sodium Chloride (Brine)

• Overall Reaction:
  2NaCl(aq)→ 2Na(s) + CI₂(g)
• At the Anode (Oxidation):
  2Cl⁻(aq)→ CI₂(g) + 2e⁻²
• At the Cathode (Reduction):
  2H⁺(aq) + 2e⁻ → H2(g)

Process:

• Chloride ions (Cl⁻) are oxidized to chlorine gas at the anode.
• Hydrogen ions (H⁺) are reduced to hydrogen gas at the cathode.

Memory Aid: AN OIL RIG CAT

• AN: Anode for Oxidation (Loss of electrons).
• CAT: Cathode for Reduction (Gain of electrons).

2.2 Fuel Cells

Definition:

• Fuel Cell: A device that converts chemical energy from a fuel into electricity through a redox reaction.

Key Concepts:

• Common Fuel Cell: Hydrogen-Oxygen Fuel Cell.
• Reactions:
  • At the Anode (Oxidation):
    2H₂(g) → 4H⁺(aq) + 4e⁻²
  • At the Cathode (Reduction):
    2O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)
  • Overall Reaction:
    2H₂(g) + O₂(g) → 2H₂O(l) + Energy

Process:

• Hydrogen is oxidized to protons and electrons at the anode.
• Oxygen is reduced by accepting electrons and protons to form water at the cathode.
• Energy is released and can be used as electrical power.

Advantages:

• High Efficiency: More efficient than traditional combustion.
• Clean Energy: Produces only water as a by-product.

2.3 Summary of Electrolysis and Fuel Cells

• Electrolysis: Uses electrical energy to cause oxidation and reduction, decomposing compounds.
• Fuel Cells: Utilize redox reactions to produce electrical energy from chemical fuels.
• Both Processes:
  • Involve separation of oxidation and reduction at different electrodes.
  • Follow the principles of electron loss (oxidation) and gain (reduction).

3. Key Vocabulary

• Redox Reaction: A reaction involving both oxidation and reduction processes.
• Oxidizing Agent: A substance that causes another to oxidize by accepting electrons; it is itself reduced.
• Reducing Agent: A substance that causes another to reduce by donating electrons; it is itself oxidized.
• Displacement Reaction: A reaction where a more reactive element displaces a less reactive element from its compound.
• Anode: The electrode where oxidation occurs in electrolysis.
• Cathode: The electrode where reduction occurs in electrolysis.
• Electrolysis: A process that uses electricity to drive a non-spontaneous chemical reaction.
• Fuel Cell: A device that converts chemical energy into electrical energy through redox reactions.

5. Additional Key Concepts

5.1 Balancing Redox Equations

Methods:

• Half-Reaction Method: Separate the oxidation and reduction processes into half-reactions, balance each for mass and charge, and then combine them.
• Ion-Electron Method: Focus on balancing the number of electrons lost and gained in the reaction.

Steps:

1. Write the unbalanced equation.
2. Separate into oxidation and reduction half-reactions.
3. Balance all atoms except hydrogen and oxygen.
4. Balance oxygen atoms by adding H₂O.
5. Balance hydrogen atoms by adding H⁺ (in acidic solutions) or OH⁻ (in basic solutions).
6. Balance the charge by adding electrons (e⁻).
7. Multiply the half-reactions by appropriate coefficients to equalize electrons.
8. Add the half-reactions together and simplify.

5.2 Oxidizing and Reducing Agents in Detail

Reducing Agents:

• Function: Donate electrons to another substance, causing it to reduce.
• Common Examples:
  • Hydrogen (H₂): Used in hydrogenation reactions.
  • Carbon (C): Used in the reduction of metal ores.
  • Carbon Monoxide (CO): Acts as a reducing agent in the blast furnace.

Oxidizing Agents:

• Function: Accept electrons from another substance, causing it to oxidize.
• Common Examples:
  • Oxygen (O₂): In combustion and respiration.
  • Hydrogen Peroxide (H₂O₂): Used in disinfection and bleaching.
  • Potassium Permanganate (KMnO₄): Used as an oxidizing agent in titrations.
  • Potassium Dichromate (K₂Cr₂O₇): Used in cleaning glassware.

5.3 Applications of Redox Reactions

• Metallurgy: Extraction and purification of metals through reduction of metal oxides.
• Energy Production: Combustion engines, batteries, and fuel cells rely on redox reactions.
• Biological Systems: Cellular respiration and photosynthesis involve redox processes.
• Environmental Processes: Decomposition of pollutants and corrosion control utilize redox reactions.