The reactivity series and properties of metals: WASSCE Chemistry

Gold stays shiny for thousands of years, iron rusts within weeks and potassium bursts into flame in water. Placing metals in order of reactivity explains all of this, and predicts how each metal reacts and how it must be extracted. This lesson builds the reactivity series from experiments and uses it.

Learning objectives

By the end of this lesson you should be able to:

state the general physical and chemical properties of metals
arrange metals in a reactivity series from their reactions with water, acids and oxygen
predict displacement reactions using the reactivity series
relate the stability of metal compounds to the reactivity of the metal
explain the position of carbon and hydrogen in the series

Properties of metals

Physical properties (most metals):

shiny (lustrous) when freshly cut
good conductors of heat and electricity
malleable (can be hammered into sheets) and ductile (can be drawn into wires)
high melting points and densities (sodium and potassium are exceptions)
sonorous (ring when struck)

Chemical properties:

form positive ions by losing electrons: M → M^n+ + ne⁻ (metals are reducing agents)
form basic oxides (some amphoteric)
the more reactive metals displace hydrogen from water or acids

Building the reactivity series

Metal	Cold water	Steam	Dilute HCl	Burning in air
K	Violent; H₂ ignites (lilac flame)	Explosive	Explosive	Burns brightly
Na	Vigorous; floats, melts, darts around	Explosive	Explosive	Burns with yellow flame
Ca	Steady; sinks, fizzes	Vigorous	Vigorous	Burns with brick-red flame
Mg	Very slow	Burns brightly	Rapid	Burns with dazzling white flame
Al	No reaction (oxide layer)	Slow (when hot)	Moderate, after the oxide layer goes	Burns as powder
Zn	No reaction	Reacts when red hot	Moderate	Burns when strongly heated
Fe	Rusts slowly	Reacts when red hot (reversible)	Slow	Burns as wool or powder
Pb	No reaction	No reaction	Very slow	Forms oxide on heating
Cu	No reaction	No reaction	No reaction	Forms black oxide on heating; does not burn
Ag, Au	No reaction	No reaction	No reaction	No reaction

From these results, the reactivity series (most reactive first):

K > Na > Ca > Mg > Al > (C) > Zn > Fe > Sn > Pb > (H) > Cu > Hg > Ag > Au

The reactivity series from potassium to gold, with carbon and hydrogen marked, and the method used to extract each group of metals

Figure 1. The reactivity series and how each metal is extracted. WASSCE Tutor diagram.

A useful order to memorise: "Please Send Charlie's Monkeys And Zebras In Large Hired Cages Making Sure Gorillas" for K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au.

Carbon and hydrogen

The non-metals carbon and hydrogen are included as reference points:

Carbon (between Al and Zn) can reduce the oxides of metals below it: zinc, iron, tin, lead, copper.
Hydrogen (between Pb and Cu): metals above it displace hydrogen from dilute acids; metals below it (Cu, Ag, Au) do not.

Displacement reactions

A more reactive metal displaces a less reactive metal from a solution of its salt or from its oxide.

Worked example

Iron nail in copper(II) sulphate solution.

\text{Fe(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{FeSO}_4\text{(aq)} + \text{Cu(s)}

\begin{gathered}\text{Fe(s)} + \text{CuSO}_4\text{(aq)} \\[6pt] \rightarrow \text{FeSO}_4\text{(aq)} + \text{Cu(s)}\end{gathered}

The iron becomes coated with reddish-brown copper, and the blue solution turns pale green (Fe²⁺). Iron is above copper, so it loses electrons more readily: Fe → Fe²⁺ + 2e⁻ and Cu²⁺ + 2e⁻ → Cu.

Copper placed in iron(II) sulphate solution shows no reaction, because copper is less reactive than iron.

The thermite reaction: aluminium displaces iron from iron(III) oxide so violently that molten iron is produced. It is used to weld railway lines:

\text{Fe}_2\text{O}_3\text{(s)} + 2\text{Al(s)} \rightarrow \text{Al}_2\text{O}_3\text{(s)} + 2\text{Fe(l)}

\begin{gathered}\text{Fe}_2\text{O}_3\text{(s)} + 2\text{Al(s)} \\[6pt] \rightarrow \text{Al}_2\text{O}_3\text{(s)} + 2\text{Fe(l)}\end{gathered}

Reactivity and the stability of compounds

The more reactive a metal, the more stable its compounds, and the harder they are to break down.

Compound	Reactive metals (K, Na)	Middle metals (Ca to Cu)	Least reactive (Ag, Au)
Oxide + heat	Stable	Stable (reduced by carbon if below Al)	Decomposes: 2Ag₂O → 4Ag + O₂
Hydroxide + heat	Stable	→ oxide + water: Cu(OH)₂ → CuO + H₂O	—
Carbonate + heat	Stable	→ oxide + CO₂	→ metal + CO₂ + O₂
Nitrate + heat	→ nitrite + O₂	→ oxide + NO₂ + O₂	→ metal + NO₂ + O₂

So potassium must be extracted by electrolysis, while silver oxide decomposes just by heating, and gold is found as the free metal.

In Sierra Leone

Artisanal miners in Sierra Leone find gold as tiny grains of the free metal in river gravels, because gold is so unreactive that it never combines with oxygen or sulphur. Iron, a more reactive metal, is never found free: at Tonkolili and Marampa it occurs as iron oxide ore, which must be chemically reduced to obtain the metal.

Common mistake

Aluminium is high in the series yet seems unreactive, for example in cooking pots. A thin, tough layer of aluminium oxide protects it. Remove the layer, by scratching it under mercury(II) chloride solution for instance, and aluminium reacts as vigorously as its position predicts.

Key points

Metals are lustrous, conducting, malleable and ductile; they form positive ions and basic oxides.
Reactivity series: K, Na, Ca, Mg, Al, (C), Zn, Fe, Sn, Pb, (H), Cu, Hg, Ag, Au.
Reactions with water, steam, acids and oxygen become less vigorous down the series.
A more reactive metal displaces a less reactive one (Fe + CuSO₄; thermite).
Carbon reduces oxides of metals below it; metals above hydrogen react with dilute acids.
The more reactive the metal, the more stable its compounds.

WASSCE examination focus

WAEC tests this

Expect to arrange metals in order from experimental observations, to predict and write equations for displacement reactions, to explain the positions of carbon and hydrogen, and to relate the effect of heat on compounds to the reactivity of the metal. The unusual behaviour of aluminium is often asked.

Quick check

Name a metal that reacts with steam but only very slowly with cold water.
What is seen when zinc is placed in copper(II) sulphate solution?
Why does copper not react with dilute hydrochloric acid?
Why is aluminium used for cooking pots even though it is a reactive metal?

Answers:

Magnesium (or iron or zinc at red heat).
Reddish-brown copper forms on the zinc and the blue colour fades.
Copper is below hydrogen in the reactivity series, so it cannot displace hydrogen from the acid.
It is protected by a thin, unreactive layer of aluminium oxide.

The reactivity series and properties of metals