Two halves of one reaction
A redox reaction is one in which oxidation and reduction take place at the same time. Historically oxidation meant gaining oxygen or losing hydrogen, but the modern definition is based on electron transfer: oxidation is the loss of electrons and reduction is the gain of electrons (mnemonic OIL RIG).
Consider the displacement of copper by zinc:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Zinc atoms lose electrons (oxidation): Zn → Zn2+ + 2e−. Copper(II) ions gain electrons (reduction): Cu2+ + 2e− → Cu. The species that accepts electrons is the oxidising agent (here Cu2+); the species that donates electrons is the reducing agent (here Zn).
Oxidation number
The oxidation number is a signed value that tracks electron control. Key rules: uncombined elements are 0; Group 1 is +1; Group 2 is +2; oxygen is usually −2 (−1 in peroxides); hydrogen is +1 (−1 in metal hydrides); the sum equals the ion charge. An increase in oxidation number = oxidation; a decrease = reduction.
Redox reactions are everywhere: respiration, combustion, photosynthesis and the bleaching of dyes all rely on electron transfer. To balance a redox equation, balance each half-equation for atoms first, then balance the charge by adding electrons, and finally scale the two halves so that the number of electrons lost equals the number gained before adding them together.
Worked example
Find the oxidation number of manganese in KMnO4. K is +1, each O is −2 (total −8). Let Mn = x. Then +1 + x + (−8) = 0, so x = +7. In the reaction MnO4− → Mn2+, Mn drops from +7 to +2, a gain of 5 electrons, so it is reduced and acts as an oxidising agent.
Remember
- Oxidation = loss of electrons = rise in oxidation number.
- Reduction = gain of electrons = fall in oxidation number.
- The oxidising agent is itself reduced; the reducing agent is itself oxidised.