Redox Half-Equations and Self-Indicating Titrations

Full Notes

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:

ExampleMg → Mg²⁺ + 2e⁻

Reduction
Electrons are gained, written on the left:

ExampleCl₂ + 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.

Combining Half-Equations

To form a full redox equation from two half-equations:

• 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).

Example: Copper reacting with silver ions

Oxidation: Cu → Cu²⁺ + 2e⁻
Reduction: Ag⁺ + e⁻ → Ag

Multiply Ag⁺ equation by 2: 2Ag⁺ + 2e⁻ → 2Ag
Now combine: Cu + 2Ag⁺ → Cu²⁺ + 2Ag

When to Use H⁺ and H₂O in Half-Equations

In aqueous solutions, particularly under acidic conditions, some redox reactions involve ions like H⁺ and H₂O to balance atoms and charges.

• Balance oxygen atoms with H₂O
• Balance hydrogen atoms with H⁺
• Then balance charge with electrons

This is common in half-equations for:

• Transition metals (e.g. MnO₄⁻, Cr₂O₇²⁻)
• Oxygen-containing species
• Acidic redox reactions

Example: Acidic reduction of manganate(VII) ions

Unbalanced reaction: MnO₄⁻ → Mn²⁺

Balance O with H₂O: MnO₄⁻ → Mn²⁺ + 4H₂O

Balance H with H⁺: 8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O

Balance charge with e⁻: 8H⁺ + MnO₄⁻ + 5e⁻ → Mn²⁺ + 4H₂O

Note that OH^– ions can also be used to balance O and H in alkaline conditions.

Self-Indicating Redox Titrations?

Some redox reactions involve coloured species that change as they are oxidized or reduced. If being reacted together in a titration, these do not require an external indicator.

They are called self-indicating because the colour change of the reacting species indicates the endpoint.

Example: Titration of Fe²⁺ with MnO₄⁻

MnO₄⁻ ions are reduced to Mn²⁺, with a colour change from purple to colourless.

Endpoint = first permanent pink colour (excess MnO₄⁻).

Summary

• Half-equations show oxidation and reduction separately.
• Balance atoms with H₂O and H⁺, then charges with electrons.
• Combine half-equations so electrons cancel to form full redox equations.
• Self-indicating titrations use colour changes of the reactants themselves.