Grade 9

Metals and the activity series

Use displacement reactions and hydrogen as a reference to compare how readily metals form ions.

1 Lower secondary 2 Upper secondary 3 Undergraduate 4 Advanced 5 Research

IntuitionA ladder of readiness to react

A strip of zinc placed in blue copper sulfate slowly becomes coated with copper. Zinc atoms give electrons to copper ions, so the blue colour fades.

This displacement is a competition: the metal that loses electrons more readily can replace the ion of a less reactive metal in solution.

The first view orders metals from K to Au; switch modes to compare Zn/CuSO₄, Cu/AgNO₃ and the non-reaction of Cu/ZnSO₄.

The Zn/Cu pair illustrates the same electron-transfer tendency in a separated galvanic cell.

SchoolHow to read the series

Definition: Reactivity series

An empirical ordering of metals by their tendency to form cations. Metals higher in the usual school series are stronger reducing agents; hydrogen is a reference, not a metal.

Typical school activity series (decreasing reactivity left to right)
Position	Metal / reference	Common observation
1	K, Na, Ca	react with cold water
2	Mg, Al, Zn, Fe	react with acids; rate varies
3	Pb, (H), Cu, Ag, Au	below H: no H₂ from dilute non-oxidising acids

A metal above another in the series can often displace its ions: $Zn + CuSO₄$ proceeds, whereas Cu + ZnSO₄ does not. The spectator ions are omitted in the net ionic equation.

\ce{Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s)}

\ce{Zn + CuSO4 -> ZnSO4 + Cu}

Example: Will copper displace zinc from zinc sulfate?

Predict whether a clean copper strip reacts with aqueous ZnSO₄.

Solution

No. Zn is above Cu and is more readily oxidised; Cu cannot reduce Zn²⁺ to Zn under these conditions.

Example: Hydrogen from magnesium and acid

Write the net ionic equation for Mg(s) in dilute hydrochloric acid.

Solution

Mg lies above H and displaces it: $Mg + 2H⁺ → Mg²⁺ + H₂$ . Chloride ions are spectators.

UndergraduateElectrode potentials explain the ordering

Standard reduction potentials are quantitative equilibrium measures, not a memorised magic ladder. A more negative E°(M²⁺/M) usually means oxidation of M is more favourable relative to the standard hydrogen electrode; actual cell voltage also depends on activities and conditions.

Selected standard reduction potentials at 25 °C
Half-reaction	E° / V
K⁺ + e⁻ → K	−2.93
Zn²⁺ + 2e⁻ → Zn	−0.76
Fe²⁺ + 2e⁻ → Fe	−0.44
2H⁺ + 2e⁻ → H₂	0.00
Cu²⁺ + 2e⁻ → Cu	+0.34

E^\circ_{\rm cell}=E^\circ_{\rm cathode}-E^\circ_{\rm anode}

Example: Predict Zn/Cu cell voltage

Use E°(Cu²⁺/Cu)=+0.34 V and E°(Zn²⁺/Zn)=−0.76 V.

Solution

Zn is oxidised at the anode and Cu²⁺ reduced at the cathode. E°cell = 0.34 − (−0.76) = +1.10 V. A positive value indicates spontaneity in the written direction under standard conditions.

For non-standard concentrations at 25 °C, the Nernst equation shifts the potential: concentration can reverse or weaken a prediction from the simple series. K and Na do not behave as freely depositing metals in water because water is reduced preferentially.

References

Chemistry: The Central Science · Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Catherine J. Murphy, Patrick M. Woodward, Matthew W. Stoltzfus, 2018