The Reactivity Series & Extraction of Metals

Why some metals are reactive

Metals lose electrons when they react. The more easily a metal forms a positive ion, the more reactive it is. By comparing how vigorously metals react with water, steam and dilute acids, we can rank them in order of reactivity — the reactivity series.

Key terms Reactivity series — a list of metals arranged in order of how readily they react (most reactive at the top).

Displacement — a more reactive metal pushing a less reactive metal out of its compound.

Native metal — a very unreactive metal (e.g. gold) found uncombined in the Earth.

The reactivity series

Carbon and hydrogen are non-metals but are included as useful reference points — they tell us how a metal can be extracted and how it reacts with acid.

Exam tip Learn the order with a mnemonic such as "Please Send Cats Monkeys And Zebras In Cages, Securely Guarded" — Potassium, Sodium, Calcium, Magnesium, Aluminium, Zinc, Iron, Copper, Silver, Gold (with Carbon and Hydrogen placed in their gaps).

Reactions with water and steam

The most reactive metals react with cold water to give a metal hydroxide and hydrogen:

$$2Na + 2H_2O \rightarrow 2NaOH + H_2$$

Potassium, sodium, calcium react with cold water (potassium fastest, fizzing and igniting).
Magnesium reacts only very slowly with cold water but readily with steam to give the oxide and hydrogen:

$$Mg + H_2O \rightarrow MgO + H_2$$

Zinc and iron react with steam when heated strongly, but copper, silver and gold do not react with water or steam at all.

Reactions with dilute acid

Metals above hydrogen react with dilute acids (e.g. hydrochloric or sulfuric) to give a salt and hydrogen. Metals below hydrogen (copper, silver, gold) do not react.

$$Mg + 2HCl \rightarrow MgCl_2 + H_2$$

The more reactive the metal, the faster the fizzing (rate of hydrogen given off). This is a common way to compare reactivity in the lab.

Watch out Potassium, sodium and calcium react dangerously fast (even explosively) with dilute acid, so this reaction is not used to rank them — their reaction with water is observed instead.

Displacement reactions

A more reactive metal displaces a less reactive metal from a solution of its salt. For example, iron displaces copper from copper(II) sulfate solution — the blue solution fades and a pink-brown copper coating forms on the iron:

$$Fe + CuSO_4 \rightarrow FeSO_4 + Cu$$

The ionic equation shows what actually changes:

$$Fe + Cu^{2+} \rightarrow Fe^{2+} + Cu$$

Iron atoms give up electrons (oxidised) and copper ions accept them (reduced). The sulfate ions are spectator ions — they appear unchanged on both sides, so they are left out.

Worked example Will zinc displace silver from silver nitrate solution?

Zinc is above silver in the series, so yes.

$Zn + 2AgNO_3 \rightarrow Zn(NO_3)_2 + 2Ag$

Ionic: $Zn + 2Ag^+ \rightarrow Zn^{2+} + 2Ag$

Using the series to choose an extraction method

A metal found combined in an ore must have its compound broken down to release the metal. How easily this can be done depends on the metal's reactivity — which decides the extraction method.

The rule: a metal more reactive than carbon cannot be displaced by carbon, so it must be extracted by electrolysis. A metal less reactive than carbon can be reduced by carbon, which is much cheaper.

Key terms Ore — a rock containing enough of a metal compound to be worth extracting.

Reduction — removal of oxygen (or gain of electrons).

Oxidation — addition of oxygen (or loss of electrons).

Extraction of iron in the blast furnace

Iron is less reactive than carbon, so it is reduced with carbon in a blast furnace. The raw materials are iron ore (haematite, Fe₂O₃), coke (carbon) and limestone (calcium carbonate), with hot air blown in.

The key reactions:

1. Coke burns in the hot air, giving out heat:

$$C + O_2 \rightarrow CO_2$$

1. More coke reduces this to carbon monoxide:

$$CO_2 + C \rightarrow 2CO$$

1. Carbon monoxide reduces the iron ore to molten iron (the main reaction):

$$Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$$

The limestone removes sandy impurities. It decomposes to calcium oxide, which reacts with silicon dioxide to form slag (calcium silicate), which floats on top of the iron and is run off separately:

$$CaCO_3 \rightarrow CaO + CO_2$$

$$CaO + SiO_2 \rightarrow CaSiO_3$$

Extraction of aluminium by electrolysis

Aluminium is above carbon, so carbon cannot reduce it — it is extracted by electrolysis of molten aluminium oxide (Al₂O₃) from the ore bauxite. The oxide is dissolved in molten cryolite to lower its melting point and save energy.

At the cathode (–) aluminium ions gain electrons (reduction):

$$Al^{3+} + 3e^- \rightarrow Al$$

At the anode (+) oxide ions lose electrons (oxidation), forming oxygen:

$$2O^{2-} \rightarrow O_2 + 4e^-$$

The hot oxygen burns away the carbon anodes (forming CO₂), so they must be replaced regularly. Electrolysis uses huge amounts of electricity, which is why aluminium is expensive to produce.

Oxidation and reduction (redox)

Oxidation and reduction always happen together — this is a redox reaction. Use OIL RIG:

Key terms OIL RIG — Oxidation Is Loss of electrons, Reduction Is Gain of electrons.

In terms of oxygen: oxidation = gaining oxygen, reduction = losing oxygen.

In $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$, the iron oxide loses oxygen (reduced) and carbon monoxide gains oxygen (oxidised).

Rusting of iron

Rust is hydrated iron(III) oxide. Iron rusts only when both water and oxygen are present:

$$iron + oxygen + water \rightarrow hydrated\ iron(III)\ oxide$$

A classic experiment uses three test tubes: iron in water + air (rusts), iron in boiled water with oil on top — no oxygen (no rust), and iron in dry air with a drying agent like calcium chloride — no water (no rust). Salt water speeds rusting up.

Preventing rusting works by keeping out water and oxygen, or by sacrificing a more reactive metal:

Barrier methods — painting, oiling/greasing, coating with plastic.
Galvanising — coating iron with zinc. The zinc forms a barrier and, being more reactive, corrodes first.
Sacrificial protection — attaching blocks of a more reactive metal (e.g. zinc or magnesium) to ships' hulls or pipelines. The reactive metal is oxidised instead of the iron and is replaced when used up.

Real world Steel ships have zinc blocks bolted to the hull below the waterline. The zinc dissolves away slowly, protecting the steel — this is sacrificial protection in action, the same principle that galvanising uses on a smaller scale.

Exam tip If asked why galvanising protects iron even when scratched, say: zinc is more reactive than iron, so the zinc is oxidised in preference — the exposed iron does not rust.