Redox I: Ionic Half-Equations and Redox Equations
Quick Notes
- Half-equations show oxidation or reduction processes separately.
- Oxidation half-equation: electrons are lost (appear on the right).
- Reduction half-equation: electrons are gained (appear on the left).
- In aqueous solutions, particularly under acidic conditions, some redox reactions involve H⁺, OH- and H₂O to balance atoms and charges
- To write a redox equation:
- Write oxidation and reduction half-equations.
- Balance atoms and charges (especially electrons).
- Combine half-equations, ensuring electrons cancel.
- You must be able to:
- Write and balance ionic half-equations.
- Combine half-equations into a full balanced redox equation.
Full Notes
What Are Ionic Half-Equations?
Half-equations allow us to show oxidation and reduction steps separately. Each redox reaction can be broken down into 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
For Oxidation: electrons are lost and written on the right.
Example Oxidation of magnesium
Mg → Mg²⁺ + 2e⁻
For Reduction: electrons are gained and written on the left.
Example Reduction of chlorine
Cl₂ + 2e⁻ → 2Cl⁻
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:
- 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⁻
Combined: Mg + Cl₂ → MgCl₂
Electrons cancel out (2e⁻ lost = 2e⁻ gained).
More Complex Example
Example Copper reacting with silver ions
Half-equations:
- Oxidation: Cu → Cu²⁺ + 2e⁻
- Reduction: Ag⁺ + e⁻ → Ag
Balance electrons:
- Multiply silver equation by 2: 2Ag⁺ + 2e⁻ → 2Ag
Combined: 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. These help:
- 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⁻) and oxygen-containing species in acidic redox reactions.
Example Acidic reduction of MnO₄⁻ to Mn²⁺
- Balance Mn atoms (already 1 each side).
- Balance O with H₂O: MnO₄⁻ → Mn²⁺ + 4H₂O.
- Balance H with H⁺: 8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O.
- Balance charge with e⁻: add 5e⁻ to left side.
Final equation: 8H⁺ + MnO₄⁻ + 5e⁻ → Mn²⁺ + 4H₂O
Summary
- Half-equations show oxidation (electrons lost) and reduction (electrons gained) separately.
- Oxidation: e⁻ on right, Reduction: e⁻ on left.
- Combine half-equations by balancing electrons and adding them together.
- In acidic solutions, use H⁺ and H₂O to balance atoms and charges.
- Always ensure electrons cancel to give a full balanced redox equation.