Reactivity Series — Edexcel Linear - OL (Chemistry) Revision Notes & Topical Questions

The reactivity series lists metals from most to least reactive: potassium, sodium, lithium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold. Reactivity is the tendency of a metal to lose electrons and form positive ions. Position is judged by how vigorously a metal reacts with oxygen, water, acids and by displacement.

Equations

• Ca + 2H₂O → Ca(OH)₂ + H₂;
  • Mg + H₂O(g) → MgO + H₂;
  • Zn + 2HCl → ZnCl₂ + H₂.

Collecting hydrogen

• test with a lit splint, which gives a squeaky pop.
• Rate of bubbling or temperature rise provides a way to rank metals when the acid and its concentration are kept the same.

A more reactive metal displaces a less reactive metal from solutions of its salts, and from its oxide on heating.

Results can be arranged in a grid.

The more reactive the metal, the more stable its compounds. Carbonates and hydroxides of reactive metals (sodium, potassium) resist heating; those of less reactive metals (copper carbonate → copper oxide + carbon dioxide) decompose readily. Metal nitrates decompose similarly, with copper nitrate giving brown nitrogen dioxide. This ties reactivity to extraction method in the later lesson.

To rank metals by their reaction with acid, use equal-sized pieces (or equal masses) of each metal, the same volume and concentration of dilute hydrochloric acid at the same starting temperature, and measure either the time taken to produce a fixed volume of hydrogen or the temperature rise.

Safetywear eye protection, use small quantities of the very reactive metals, never put potassium or sodium into acid, and keep flames away from hydrogen.

Oxidation and reduction always happen together, so a reaction involving both is a redox reaction. Three definitions are used:

A useful memory aid is OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain. The electron definition is the most powerful because it works for reactions where oxygen and hydrogen are not involved, such as the formation of ionic compounds.

An oxidising agent oxidises another species by taking electrons from it. The oxidising agent is itself reduced, so it gains electrons. A reducing agent gives electrons to another species and is itself oxidised, so it loses electrons.

Common oxidising agents

• oxygen, chlorine, acidified potassium manganate(VII) (purple to colourless when reduced), potassium dichromate(VI) (orange to green).
• Common reducing agents: hydrogen, carbon, carbon monoxide and reactive metals.

In displacement Zn + CuSO₄ → ZnSO₄ + Cu, the sulfate ions do not change and are spectator ions.

In halogen displacement, Cl₂ + 2Br^- → 2Cl^- + Br₂: chlorine is reduced (gains electrons) and bromide is oxidised (loses electrons). Higher reactivity of a metal means it is a stronger reducing agent; higher reactivity of a halogen means it is a stronger oxidising agent.

In electrolysis, oxidation occurs at the anode (positive electrode, where negative ions lose electrons) and reduction at the cathode (negative electrode, where positive ions gain electrons). Remember "An Ox, Red Cat".

Redox links many topics. The table summarises where each change happens.

When asked to "explain in terms of electrons", always say which species loses electrons and which gains them, never just state that oxygen is added or removed.