Chemistry 30 · Redox
Oxidation and reduction
Redox is bookkeeping for electrons. Assign oxidation numbers, see which ones moved, split the reaction into half-reactions, and a table of reduction potentials will tell you whether the reaction happens at all.
- 1. Oxidation numbers
- 2. Who does what to whom
- 3. Half-reactions and balancing
- 4. Predicting whether it happens
Oxidation numbers
The idea: an oxidation number is a bookkeeping charge — what an atom's charge would be if every bond were fully ionic. Comparing before and after tells you which atoms lost or gained electrons.
| Rule | Example |
|---|---|
| An element by itself is 0 | Na, O₂, Cl₂, Fe all 0 |
| A monatomic ion equals its charge | Na⁺ is +1, S²⁻ is −2 |
| Oxygen is −2 (−1 in peroxides) | −2 in H₂O; −1 in H₂O₂ |
| Hydrogen is +1 (−1 with metals) | +1 in HCl; −1 in NaH |
| The numbers add to the overall charge | 0 for a compound, the ion charge for an ion |
Worked: sulfur in SO₄²⁻
Four oxygens at −2 total −8. The ion is −2 overall, so sulfur must make up the difference: x + (−8) = −2, giving x = +6.
The same method handles anything: in Cr₂O₇²⁻, seven oxygens give −14, the ion is −2, so two chromiums share +12 — each is +6.
Reading the change: oxidation number up means electrons were lost — oxidation. Down means electrons were gained — reduction. OIL RIG: Oxidation Is Loss, Reduction Is Gain.
Who does what to whom
The idea: the agents are named for what they do to their partner, not for what happens to them — which is why they always sound backwards at first.
| Substance | Electrons | Oxidation number | Called |
|---|---|---|---|
| Is oxidized | Loses them | Increases | The reducing agent |
| Is reduced | Gains them | Decreases | The oxidizing agent |
In 2Mg(s) + O₂(g) → 2MgO(s), magnesium goes 0 → +2 (oxidized, so it is the reducing agent) and oxygen goes 0 → −2 (reduced, so it is the oxidizing agent). Every redox reaction has one of each: electrons have to come from somewhere and go somewhere.
Watch out: not every reaction is redox. In neutralization and precipitation, ions swap partners but no oxidation number changes. Check the numbers before reaching for half-reactions.
Half-reactions and balancing
The idea: split the reaction into the losing half and the gaining half, balance each one completely, then scale them so the electrons cancel.
- Balance the main atom in the half-reaction.
- Balance oxygen with H₂O, adding a water for each oxygen needed.
- Balance hydrogen with H⁺.
- Balance the charge with electrons, added to whichever side makes both sides match.
- Scale both halves so the electrons lost equal the electrons gained, then add and cancel.
- In basic solution, finish by adding OH⁻ to both sides to neutralize every H⁺, and simplify the water.
Worked: permanganate in acid
MnO₄⁻ → Mn²⁺
- Manganese is already balanced.
- Four oxygens on the left, so add four waters on the right: MnO₄⁻ → Mn²⁺ + 4H₂O.
- Eight hydrogens on the right, so add 8H⁺ on the left.
- Charge is +7 on the left and +2 on the right, so add 5 electrons to the left.
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. The 5 electrons match manganese dropping from +7 to +2, which is the quickest way to check the answer.
Watch out: electrons never appear in a final balanced equation. If any are left, the two halves were not scaled to match.
Predicting whether it happens
The idea: a table of standard reduction potentials ranks every half-reaction. The strongest oxidizing agent present takes electrons from the strongest reducing agent present — if the numbers allow it.
- List every species in the mixture, including water if it is an aqueous solution.
- Find the strongest oxidizing agent — the one highest on the left of the table, with the most positive reduction potential.
- Find the strongest reducing agent — the one lowest on the right.
- If the oxidizing agent sits above the reducing agent, the reaction is spontaneous.
- Check with the numbers: E°cell = E°(cathode) − E°(anode). A positive value means spontaneous.
Worked: zinc and copper
Cu²⁺ + 2e⁻ → Cu has E° = +0.34 V; Zn²⁺ + 2e⁻ → Zn has E° = −0.76 V.
Copper's half-reaction is higher, so Cu²⁺ is reduced and zinc metal is oxidized:
E°cell = 0.34 − (−0.76) = +1.10 V, so a strip of zinc dropped into copper sulfate really does plate with copper.
Reverse the pairing — copper metal in a zinc solution — and E°cell would be −1.10 V, which is the table's way of saying nothing happens.
Never multiply E° by anything. Doubling a half-reaction doubles the electrons but not the voltage: potential is energy per charge, so it does not scale with amount — unlike ΔH, which does.