Half-Reaction Method
Balance redox reactions by separating oxidation and reduction half-reactions, then combining them.
The Half-Reaction Method balances redox reactions by treating oxidation and reduction separately, then combining them.
Remember: In acidic solutions use H⁺ and H₂O. In basic solutions, balance as if acidic, then add OH⁻ to neutralize all H⁺ ions, creating H₂O.
A redox (reduction-oxidation) reaction involves the transfer of electrons between species. One reactant loses electrons (oxidation) while another gains them (reduction). Balancing these equations means finding whole-number coefficients so that both atoms and charge are conserved on both sides.
The most reliable technique is the half-reaction method: split the overall reaction into an oxidation half-reaction and a reduction half-reaction, balance each separately (atoms, then charge), then combine them so electrons cancel out completely.
Step 1: Split into Half-Reactions
Assign oxidation numbers and separate the species being oxidized from the species being reduced.
Step 2: Balance Atoms Other Than O and H
Balance every element except oxygen and hydrogen first.
Step 3: Balance Oxygen with H₂O
Add H₂O to the side that needs oxygen.
Step 4: Balance Hydrogen with H⁺
Add H⁺ ions to balance hydrogen atoms (acidic solution).
Step 5: Balance Charge with Electrons
Add e⁻ to the more positive side so the net charge matches on both sides of each half-reaction.
Step 6: Equalize and Add Electrons
Multiply each half-reaction so the number of electrons lost equals the number gained, then add the two half-reactions together and cancel electrons.
Basic Solution: One Extra Step
Balance as if acidic, then add OH⁻ to both sides to neutralize every H⁺ (forming H₂O), and simplify.
Problem:
Balance the reaction between iron(II) ions and permanganate ion in acidic solution: Fe²⁺ + MnO₄⁻ → Fe³⁺ + Mn²⁺
Step 1: Write the half-reactions
Oxidation: Fe²⁺ → Fe³⁺
Reduction: MnO₄⁻ → Mn²⁺
Step 2: Balance O with H₂O, then H with H⁺
MnO₄⁻ → Mn²⁺ + 4H₂O
8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O
Step 3: Balance charge with electrons
Fe²⁺ → Fe³⁺ + e⁻
5e⁻ + 8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O
Step 4: Equalize electrons (×5 on oxidation) and add
5Fe²⁺ → 5Fe³⁺ + 5e⁻
5e⁻ + 8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O
Balanced Equation:
5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
Both atoms and charge (+17 on each side) are conserved.
Balancing atoms but forgetting charge
A correctly atom-balanced equation can still have unequal net charge on each side.
Mixing up H⁺ and OH⁻
Use H⁺ for acidic solution; only convert to OH⁻/H₂O at the final step for basic solution.
Not equalizing electrons before adding half-reactions
The electrons lost in oxidation must exactly equal the electrons gained in reduction.