- Calculate the oxidation state of an element in a compound or ion
- Identify oxidation, reduction, the oxidising agent and the reducing agent
- Balance redox equations using the electron-balance method
A rusting iron nail, methane burning on a gas stove, the current from a phone battery and even our own breathing are all the same kind of reaction: electrons move from one atom to another. Such reactions are called oxidation–reduction reactions, or redox for short. In this lesson you will learn to keep track of those electrons.
Oxidation state
The charge an atom would have if all its bonds were purely ionic — a bookkeeping charge. Here we write the sign before the number (Mn⁺⁷, O⁻²), whereas a real ion charge is written number first (Mn²⁺, SO₄²⁻).
- Atoms in simple substances have oxidation state 0: Na, O₂, Cl₂, Fe.
- In a neutral molecule the oxidation states add up to 0; in an ion they add up to the charge of the ion.
- Fluorine is always −1; oxygen is usually −2 (−1 in H₂O₂, +2 in OF₂).
- Hydrogen is usually +1, but −1 in metal hydrides (NaH, CaH₂).
- Group IA metals are always +1, group IIA metals +2 and aluminium +3.
Find the oxidation states of sulfur in H₂SO₄, nitrogen in NH₄⁺ and chromium in Cr₂O₇²⁻.
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H₂SO₄: 2 · (+1) + x + 4 · (−2) = 0 → x = +6.
NH₄⁺: x + 4 · (+1) = +1 → x = −3.
Cr₂O₇²⁻: 2x + 7 · (−2) = −2 → 2x = 12 → x = +6.
Oxidation and reduction
In Zn + CuSO₄ → ZnSO₄ + Cu, each zinc atom gives away two electrons and goes from Zn⁰ to Zn⁺² — this is oxidation. The copper ion takes those electrons: Cu⁺² → Cu⁰ — this is reduction. The particle that gives electrons is the reducing agent; the one that accepts them is the oxidising agent. Oxidation and reduction always happen together: any electrons one atom gains, another atom has lost.
| Process | Electrons | Oxidation state | The substance is the… |
|---|---|---|---|
| Oxidation | lost (−e⁻) | increases | reducing agent |
| Reduction | gained (+e⁻) | decreases | oxidising agent |
Balancing by electron balance
In redox equations, finding the coefficients by trial and error is often hard. Instead we rely on one key rule: the number of electrons lost by the reducing agent equals the number gained by the oxidising agent.
The key condition of the electron balance
- 1Write the oxidation states
Find the oxidation states in the skeleton equation and mark the ones that change.
- 2Write the electron half-equations
For each changing element, write how many electrons are lost or gained.
- 3Find the multipliers
Take the LCM of the electrons lost and gained and work out a multiplier for each half-equation.
- 4Place coefficients and finish
Put the multipliers in front of the key substances, then balance metals and acid residues, and hydrogen and oxygen last.
In a blast furnace, iron(III) oxide is reduced by carbon monoxide: Fe₂O₃ + CO → Fe + CO₂. Balance the equation and find the mass of iron obtained from 16 kg of Fe₂O₃.
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C⁺² − 2e⁻ → C⁺⁴ | ×3 — oxidation, CO is the reducing agent.
Fe₂O₃ + 3CO → 2Fe + 3CO₂
M(Fe₂O₃) = 2 · 56 + 3 · 16 = 160 g/mol; n = 16 000 g / 160 g/mol = 100 mol.
n(Fe) = 2 · 100 = 200 mol → m(Fe) = 200 · 56 = 11 200 g = 11.2 kg.
Copper reacts with dilute nitric acid: Cu + HNO₃ → Cu(NO₃)₂ + NO + H₂O. Balance it by the electron-balance method.
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N⁺⁵ + 3e⁻ → N⁺² | ×2 — HNO₃ is the oxidising agent.
LCM(2, 3) = 6 → 3Cu and 2NO.
3Cu(NO₃)₂ holds 6 more nitrogen atoms whose oxidation state does not change → HNO₃: 2 + 6 = 8.
8 H atoms → 4H₂O. Oxygen check: left 8 · 3 = 24, right 18 + 2 + 4 = 24.
3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO↑ + 4H₂O
Redox all around us
Iron is oxidised in moist air; a simplified equation is 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ — that is rust. In our cells glucose is oxidised by oxygen to release energy: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O. In metallurgy, metals are reduced from their ores. Batteries send the electrons of a redox reaction through a wire — more on that in the electrochemistry lesson.
Key points
- An oxidation state is a bookkeeping charge: they add up to 0 in a molecule and to the ion charge in an ion.
- Oxidation is loss of electrons (oxidation state rises); reduction is gain of electrons (it falls).
- The reducing agent gives electrons and is oxidised; the oxidising agent accepts them and is reduced (OIL RIG).
- Electron balance: electrons lost = electrons gained; the LCM gives the multipliers.
- Corrosion, respiration, combustion, metallurgy and batteries are all based on redox.
Check yourself
10 questions. Every correct answer earns XP.