Reactivity Series of Metals
Metals above carbon → extracted by electrolysis. Between carbon and hydrogen → reduced with carbon. Below hydrogen → found native or reduced easily.
Displacement Reactions
A more reactive metal displaces a less reactive metal from its compound. This is a redox reaction. The more reactive metal is oxidised, the less reactive metal ion is reduced.
Metal–Metal Displacement
Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Ionic equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Observation: Blue solution fades (Cu²⁺ consumed); reddish-brown solid (Cu) deposits on zinc.
But: Cu + ZnSO₄ → No reaction (Cu is below Zn).
Metal–Acid Displacement
Metals above hydrogen react with dilute acids (HCl, H₂SO₄) to produce a salt + hydrogen gas.
Example: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Ionic equation: Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g)
Metals below hydrogen (Cu, Ag, Au) do not react with dilute acids.
Metal Extraction Methods
| Position | Method | Example |
|---|---|---|
| Above carbon (K, Na, Ca, Mg, Al) | Electrolysis of molten compound | 2Al₂O₃(l) → 4Al(l) + 3O₂(g) |
| Between C and H (Zn, Fe, Sn, Pb) | Reduction with carbon/CO | Fe₂O₃ + 3CO → 2Fe + 3CO₂ |
| Below hydrogen (Cu, Ag, Au, Pt) | Found native / gentle heating | 2Ag₂O → 4Ag + O₂ |
Corrosion & Rusting
Corrosion is the electrochemical degradation of a metal by reaction with its environment. Rusting is the corrosion of iron. Both oxygen and water are required.
Rusting Mechanism (electrochemical)
Anode (oxidation): 2Fe(s) → 2Fe²⁺(aq) + 4e⁻
Cathode (reduction): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)
Fe²⁺ is further oxidised to form hydrated iron(III) oxide. Rust: Fe₂O₃·xH₂O.
Saltwater accelerates rusting by increasing the conductivity of the water (more electrolytes).
Preventing Corrosion
- Barrier protection: Paint, oil, grease, or plastic coating physically blocks O₂ and H₂O from reaching the iron surface.
- Sacrificial protection (galvanisation): Coat iron with a more reactive metal (e.g. Zinc). Zinc oxidises preferentially (E° = −0.76 V vs Fe at −0.44 V), protecting the iron even if the coating is scratched.
- Stainless steel: Alloying with chromium creates a self-healing oxide layer.
Predicting Reactions with E° (HL)
At HL, the reactivity series is quantified using standard electrode potentials. A more negative E° = stronger reducing agent (more reactive metal). Calculate E°cell = E°cathode − E°anode. If positive, the reaction is spontaneous.
Worked Example: Will Zn displace Cu²⁺?
E°(Cu²⁺/Cu) = +0.34 V (cathode. Reduced)
E°(Zn²⁺/Zn) = −0.76 V (anode. Oxidised)
E°cell = +0.34 − (−0.76) = +1.10 V → spontaneous ✓
⚠ Common Exam Mistakes
Never multiply E° values: When balancing half-equations (e.g. ×2), do NOT multiply E°. It is an intensive property.
Electrode polarities: In voltaic cells, anode = (−). In electrolytic cells, anode = (+). But in both: An Ox, Red Cat.
Aqueous electrolysis (HL): If a metal has a more negative E° than water (−0.83 V), water is reduced instead → H₂ evolved at the cathode.
Think About It
Why can iron be extracted from its ore using carbon, but aluminium cannot?
Carbon is above iron in the reactivity series, so carbon can reduce iron oxide. But aluminium is above carbon, so carbon cannot reduce aluminium oxide. Electrolysis is needed instead.