Lesson 1 of 11 · 7 min

Oxidation and reduction: the classical view

NCERT §7.1

A magnesium ribbon burns with a white flame and leaves a white ash. The metal has joined with oxygen, and that joining is what chemists first called oxidation.

The story this chapter follows: A strip of zinc turns red

Drop a zinc strip into blue copper sulphate solution and, within an hour, it is coated with red-brown copper while the blue fades. Nothing was burned and no oxygen was involved, yet chemists call this an oxidation. This chapter builds up that idea step by step, from oxygen to electrons to oxidation numbers, and ends with the same reaction powering a cell.

The lesson in notes

In short

Originally, oxidation meant adding oxygen to a substance, as when magnesium burns to MgO or sulphur burns to SO₂. Oxygen is not the only reagent that counts.

The idea was broadened to cover removal of hydrogen: in 2H₂S + O₂ → 2S + 2H₂O, hydrogen is taken away from H₂S, so H₂S is oxidised.

It was widened again to adding an electronegative element: Mg + F₂ → MgF₂, Mg + Cl₂ → MgCl₂ and Mg + S → MgS all count as oxidation of magnesium.

Removal of an electropositive element is oxidation too: when K₄[Fe(CN)₆] reacts with H₂O₂ to give K₃[Fe(CN)₆], one potassium is removed per complex unit.

Reduction is the reverse in each case: removing oxygen (2HgO → 2Hg + O₂), adding hydrogen (ethene + H₂ → ethane), removing an electronegative element (2FeCl₃ + H₂ → 2FeCl₂ + 2HCl) or adding an electropositive one (2HgCl₂ + SnCl₂ → Hg₂Cl₂ + SnCl₄, where mercury(II) chloride is reduced).

Oxidation and reduction always happen together; one substance cannot be oxidised unless another is reduced. Such paired reactions are called redox reactions.