Redox

Full Notes

Oxidising and Reducing Agents

Oxidising agent: accepts electrons and becomes reduced.

Reducing agent: donates electrons and becomes oxidised.

Use changes in oxidation number to identify what is being oxidised and reduced.

What Are Ionic Half-Equations?

Half-equations allow us to show oxidation and reduction steps separately.

• Each redox reaction has two half-equations:
  • One showing oxidation (loss of electrons)
  • One showing reduction (gain of electrons)
• Electrons (e⁻) are included to show the movement of charge.

Writing Half-Equations

• Oxidation → electrons are lost, written on the right:

  Mg → Mg²⁺ + 2e⁻

• Reduction → electrons are gained, written on the left:

  Cl₂ + 2e⁻ → 2Cl⁻

Matt’s exam tip

Always remember that half equations don’t occur on their own; they’re a tool to represent just the oxidation or reduction part of a redox reaction, focusing only on the species gaining or losing electrons and ignoring everything else.

When to Use H⁺ and H₂O in Half-Equations

In aqueous solutions, particularly under acidic conditions, when writing redox reactions we sometimes also include other ions such as H⁺ and H₂O. This enables us to balance atoms and charges.

• Add H₂O to balance oxygen atoms
• Add H⁺ to balance hydrogen atoms (can use OH⁻ in alkaline conditions)

Example: Acidic Reduction of Manganate(VII) Ions (MnO₄⁻ → Mn²⁺)

Unbalanced reaction: MnO₄⁻ → Mn²⁺

Balance Mn atoms (already 1 on each side)

Balance O using H₂O – add 4H₂O to right: MnO₄⁻ → Mn²⁺ + 4H₂O

Balance H using H⁺ – add 8H⁺ to left: 8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O

Balance charge with e⁻ – add 5e⁻ to left: 8H⁺ + MnO₄⁻ + 5e⁻ → Mn²⁺ + 4H₂O

Constructing Redox Equations

To form a full redox equation:

• Write both half-equations
• Balance electrons
• Add the two equations together so electrons cancel out

Example: Reaction of magnesium with chlorine

Half-equations:

Mg → Mg²⁺ + 2e⁻

Cl₂ + 2e⁻ → 2Cl⁻

Add together: Mg + Cl₂ → MgCl₂

Electrons cancel (2e⁻ lost = 2e⁻ gained)

Predicting Electron Transfer Reactions

Reactions occur if a stronger oxidising agent meets a stronger reducing agent.

Example: If Zn is mixed with Cu²⁺, a reaction will occur. Zn will be oxidised (lose 2e⁻) and Cu²⁺ will be reduced (gain 2e⁻), because Zn is a more powerful reducing agent than Cu.

Example: If Cu is mixed with Zn²⁺, no reaction will occur. Cu is unable to reduce Zn²⁺ as it is a weaker reducing agent than Zn.

Summary

• Oxidising agents gain electrons and are reduced, reducing agents lose electrons and are oxidised.
• Half-equations show individual oxidation and reduction steps.
• Balance half-equations with H⁺, H₂O, and e⁻ as needed.
• Redox equations are built by combining half-equations and cancelling electrons.
• A reaction is feasible if a stronger reducing agent reacts with a stronger oxidising agent.