A displacement (or replacement) reaction is a type of redox reaction in which a free metal reacts with a compound and takes the place of another metal in that compound. When a metal is placed into the solution of a salt of another metal, the reaction proceeds only if the added metal loses electrons more readily than the metal already present in the salt — that is, it is more reactive.

This tendency of metals to lose electrons and form positive ions is summarised in the reactivity (activity) series, which ranks metals in decreasing order of chemical reactivity. A metal placed higher in this series can push a metal placed lower out of its salt solution because it has a greater tendency to be oxidised (lose electrons), while the ion of the less reactive metal is correspondingly reduced and deposited as free metal.

Example: Zn + CuSO₄ → ZnSO₄ + Cu. Zinc, which lies above copper in the reactivity series, loses electrons to form Zn²⁺ ions, while Cu²⁺ ions from the solution gain electrons and are deposited as metallic copper; the blue colour of the copper sulphate solution fades as reddish-brown copper is deposited on the zinc surface.

This is why the correct condition is that the displacing metal must be higher in the reactivity series than the metal in the salt.

The remaining suggested conditions do not hold up: the salt actually needs to be reasonably soluble (or at least dissociate to some extent) so that free metal ions are available in solution for the reaction to occur — insolubility would hinder rather than enable the reaction. If the added metal is identical to the metal already combined in the salt, there is no difference in reactivity between the two, so no net displacement can occur. Finally, such displacement reactions are typically exothermic rather than endothermic, since the more reactive metal releases energy as it is oxidised; an endothermic requirement plays no part in determining whether the reaction happens.