1. From the explosive to the tamest: the reactivity series of metals

Metals generally react with acids. We know this, but we often miss out the key word: generally. This sweeping statement is a gross generalisation. Metals have varying tendency to react. Some react explosively, some moderately, and others not at all. The reactivity series arranges metals from the most reactive to the least.

2. How to memorise the reactivity series without memorising: the Periodic Table

The reactivity series is written into the Periodic Table!

Highly Reactive Group I and II Metals

Group I alkali metals and Group II alkaline earth metals are the most reactive. They have the greatest tendency to lose electrons to form cations.

For the four metals we need to know, those from Group I are more reactive than those from Group II. Within each group, the metal lower down the group is more reactive.

Moderatively Reactive Metals

Transition metals are generally less reactive than the main group metals. Lead in Group IV is also moderately reactive.

The Trio of Unreactive Metals

And the least reactive metals we need to know are copper, silver, and gold. The trio sit in the same column within the transition metal hood.

3. The reactivity series tells us if a reaction will happen and under what conditions

Metal Reacts with
Cold water
Hot steam
Hydrochloric acid
Hot steam
Hydrochloric acid
No reaction

reactive metal + water ⟶ metal hydroxide + hydrogen
reactive metal + steam ⟶ metal oxide + hydrogen
reactive metal + hydrochloric acid ⟶ metal chloride + hydrogen

Firstly, reactivity affects what a metal can react with.

  • Highly reactive metals are trigger happy, reacting with acids and even water at room temperature.
  • Moderately reactive metals require a higher temperature to react with water.
  • The unreactive metals do not react at all, even when we use strong acid or high temperature.

Secondly, reactivity affects the rate of reaction. More reactive metals react more vigorously, increasing the rate of effervescence of hydrogen gas.

4. The reactivity series tells us how much a metal wants to be a cation

A more reactive metal donates electrons more readily to form a cation. The converse is true. A less reactive metal is a hoarder — it would rather have its valence electrons.

In other words, a highly reactive metal prefers to be a cation while its less reactive friend prefers to be a free element. This provides the driving force for metal displacement reaction.

Metal displacement occurs when a more reactive metal forms a cation and displaces a less reactive metal from its compounds.

5. A more reactive metal can displace a less reactive one out of its salt solution

Formation of reddish-brown copper, displaced by more reactive iron

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

When we drop iron into copper sulfate solution, a dramatic precipitation of reddish-brown copper occurs. At the same time, we form iron(II) sulfate, whereby iron has taken the place of copper. We say that iron has displaced copper from copper(II) oxide.

Fe(s) → Fe2+(aq) + 2e

This happens because iron is more reactive than copper. It has a greater propensity to lose electrons and become a cation. We can see the loss of electrons in the oxidation half-equation above.

Cu2+(aq) + 2e→ Cu(s)

On the other hand, copper is less reactive. It is more comfortable being a pure element, somewhat retaining its valence electrons. We illustrate the gain of electrons in the reduction half-equation above.

6. Metal displacement reaction can also happen outside of water

The highly exothermic reaction of aluminium displacing iron from iron(III) oxide

2Al + Fe2O3 → Al2O3 + 2Fe

Reactive metals can also displace less reactive ones from solid metal oxides. For example, because aluminium is more reactive than iron, it can displace iron from iron(III) oxide. This sparks a spectacular reaction that is highly exothermic, forming molten iron and aluminium oxide.

C + 2CuO → CO2 + 2Cu

Interestingly, carbon can behave like a metal. Its reactivity is slightly above that of zinc. This means that it can displace zinc and other less reactive metals from their oxides. Metalworkers exploit this displacement reaction to extract less reactive metals from their oxides. Sneaky!