Metals: reactivity, extraction and electrolysis

The reactivity series

The reactivity series ranks metals by how vigorously they react (e.g. with water, oxygen, dilute acids). From most to least reactive:

Potassium (K) > Sodium (Na) > Calcium (Ca) > Magnesium (Mg) > Aluminium (Al) > (Carbon — non-metal, included for extraction) > Zinc (Zn) > Iron (Fe) > (Hydrogen) > Copper (Cu) > Silver (Ag) > Gold (Au)

Reactions of metals with water:

• K, Na, Ca react vigorously with cold water → metal hydroxide + hydrogen
• Mg reacts very slowly with cold water; quickly with steam → MgO + H₂
• Metals below Mg do not react with water

Reactions with dilute acid:

• Metals above hydrogen in the series displace hydrogen from acids
• Metals below hydrogen (Cu, Ag, Au) do not react with dilute acids

Oxidation and reduction (REDOX)

Oxidation = loss of electrons (OIL) Reduction = gain of electrons (RIG) Together: REDOX reactions

OIL RIG (or LEO says GER):

• Oxidation Is Loss of electrons
• Reduction Is Gain of electrons

In ionic terms:

• A metal losing electrons to form an ion is oxidised: Mg → Mg²⁺ + 2e⁻
• A non-metal gaining electrons is reduced: Cl₂ + 2e⁻ → 2Cl⁻

Oxidation states: the apparent charge an atom has in a compound. Rules:

• Uncombined elements: 0
• Simple ions: same as charge (Na⁺ = +1, O²⁻ = −2)
• H is usually +1 (except metal hydrides: −1)
• O is usually −2 (except peroxides: −1)
• Sum in a neutral compound = 0; in a polyatomic ion = ion charge

Displacement reactions

A more reactive metal displaces a less reactive metal from a solution of its salt: Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)

This is a REDOX reaction:

• Mg is oxidised: Mg → Mg²⁺ + 2e⁻
• Cu²⁺ is reduced: Cu²⁺ + 2e⁻ → Cu

The half-equations can be combined to give the overall ionic equation: Mg + Cu²⁺ → Mg²⁺ + Cu

Extraction of metals

The method of extraction depends on a metal's position in the reactivity series:

Blast furnace (iron extraction): Iron ore (Fe₂O₃), coke C and limestone (CaCO₃) are heated together.

• Coke burns: C + O₂ → CO₂; CO₂ + C → 2CO (carbon monoxide, the reducing agent)
• Reduction of iron ore: Fe₂O₃ + 3CO → 2Fe + 3CO₂
• Limestone removes acidic silica impurities: CaCO₃ → CaO + CO₂; CaO + SiO₂ → CaSiO₃ (slag)

Aluminium extraction (electrolysis of molten Al₂O₃): Al₂O₃ has a very high melting point (~2072 °C). It is dissolved in cryolite (Na₃AlF₆) to lower the melting point to ~950 °C. Carbon electrodes are used.

• Cathode (−): Al³⁺ + 3e⁻ → Al (reduction)
• Anode (+): 2O²⁻ → O₂ + 4e⁻ (oxidation); O₂ reacts with the carbon anode, burning it away → anodes must be replaced regularly.

Electrolysis of solutions

Electrolysis uses electrical energy to decompose an ionic compound. Cations migrate to the cathode (−); anions migrate to the anode (+).

Electrolysis of copper sulfate with copper electrodes (WJEC required practical): copper is deposited at the cathode and dissolved from the anode → copper is purified / refined.

Electrolysis of dilute H₂SO₄:

• Cathode: 2H⁺ + 2e⁻ → H₂ (hydrogen gas)
• Anode: 2H₂O → O₂ + 4H⁺ + 4e⁻ (oxygen gas)
• Volume ratio H₂:O₂ = 2:1 (reflecting the formula of water)

Common examiner traps

1. OIL RIG confusion: oxidation = LOSS of electrons (not gain); reduction = GAIN of electrons.
2. Electrolysis ions go to wrong electrode: cations (+) go to cathode (−) — opposites attract.
3. Why aluminium requires electrolysis: it is too reactive to be reduced by carbon — carbon cannot reduce Al₂O₃.
4. Anode replacement: the carbon anode in aluminium smelting burns away in the oxygen produced. Students forget this detail.
5. Half-equations must be balanced for both mass and charge. Check electrons cancel when adding half-equations.