The reactivity series ranks metals by how readily they react. More reactive metals displace less reactive metals from their compounds and react with water.

The reactivity series of metals

The reactivity series of metals is a chart listing metals in order of decreasing reactivity. In general, the more reactive a metal is:

• the more vigorously it reacts with other substances
• the more easily it loses electrons to form positive ions (cations)

We can examine the reactivity of metals by observing their reactions with oxygen, water, steam and whether it displaces other metals in displacement reactions.

This table summarises the reactions of some metals in the reactivity series. Hydrogen is included for comparison.

Metal	Reaction with oxygen (when heated and at room temperature)	Reaction with water	Reaction with steam	Reactivity
Potassium (K)	Reacts vigorous when heated. Lilac flame and white solid formed. Tarnishes when freshly cut at room temperature	Reacts vigorously. Floats on the surface. Moves and fizzes. Burns with a lilac flame. Heat is released. Crackle as it disappears. Colourless solution remains.	Reaction too dangerous to be attempted.	Most reactive
Sodium (Na)	Reacts vigorously when heated. Yellow/orange flame and white solid formed. Tarnishes when freshly cut at room temperature	Reacts vigorously. Floats on the surface. Moves and fizzes. Melts to form a silvery ball. Heat is released. Crackle as it disappears. Colourless solution remains.	Reaction too dangerous to be attempted.	⬆
Calcium (Ca)	Reacts vigorously with strong heating. Brick red flame and white solid formed. Slowly forms a surface oxide at room temperature	Reacts readily. Fizzes. Grey solid rises then sinks. Heat is released. Grey solid disappears. Solution appears milky.	Reaction too dangerous to be attempted.	⬆
Magnesium (Mg)	Reacts readily with strong heating. White light and white solid formed. Slowly forms a surface oxide at room temperature	Very slow reaction. A few bubbles of gas produced.	Reacts on strong heating. White light and white solid formed.	⬆
Aluminium (Al)	Reacts readily with strong heating as a powder. White solid formed. Slowly forms a surface oxide at room temperature	No reaction	Reacts as a powder on strong heating. White solid formed.	⬆
Zinc (Zn)	Reacts steadily when heated forming a yellow solid which changes to white on cooling. Slowly forms a surface oxide at room temperature	No reaction	Reacts as a powder on strong heating. Yellow solid forms which changes to white on cooling.	⬆
Iron (Fe)	Reacts readily when heated as iron filings. Orange sparks and black solid formed. Slowly forms a surface oxide at room temperature	No reaction	Reacts as a powder on very strong heating. Black solid formed.	⬆
Copper(Cu)	Reacts on heating to form a black solid. Slowly forms a surface oxide at room temperature	No reaction	No reaction	Least reactive

Reactions of metals with oxygen in air

Many metals react with oxygen to form metal oxides. Potassium and sodium are soft metal which are easily cut exposing a shiny surface which changes to dull rapidly. The change from shiny to dull is called tarnishing.

For example:

Potassium burns with a lilac flame when heated in air.

potassium + oxygen ➞ potassium oxide

4K(s) + O₂(g) ➞ 2K₂O(s)

Magnesium reacts readily in air burning with a white light.

magnesium + oxygen ➞ magnesium oxide

2Mg(s) + O₂(g) → 2MgO(s)

Reactions of metals with water

When a metal reacts with water, a metal hydroxide and hydrogen are formed.

For example:

Sodium reacts vigorously with water.

sodium + water ➞ sodium hydroxide + hydrogen

2Na(s) + 2H₂O(l) ➞ 2NaOH(aq) + H₂(g)

Calcium reacts readily with water.

calcium + water ➞ calcium hydroxide + hydrogen

Ca(s) + 2H₂O(l) ➞ Ca(OH)₂(aq) + H₂(g)

Calcium hydroxide is slightly soluble in water so once the solution is saturated, it starts to become milky as solid calcium hydroxide appear.

The apparatus below is used to react calcium with water and collect the gas. This would not be used with sodium or potassium as they float and the reaction is too vigorous.

Reactions of metals with steam

Magnesium reacts very slowly with water. However, it reacts vigorously with steam:

magnesium + steam → magnesium oxide + hydrogen

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

Metals which react with steam form the solid metal oxide and hydrogen gas.

In general, the more reactive the metal, the more rapid the reaction. Aluminium is unusual, because it is a reactive metal that does not react with water. Its surface forms a protective layer of aluminium oxide that keeps water away from the metal below.

The apparatus used to react a metal with steam and collect the gas produced is shown below.

Displacement in solutions

More reactive metals displace less reactive metals from their compounds. For example, magnesium is more reactive than copper. It displaces copper from copper(II) sulfate solution.

magnesium + copper(II) sulfate → magnesium sulfate + copper

Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)

In this displacement reaction:

• the copper coats the magnesium
• the solution’s blue colour fades as blue copper(II) sulfate is replaced by colourless magnesium sulfate solution.

Working out a reactivity series

You can work out a reactivity series for many substances by carrying out displacement reactions. The table below shows the results of reacting different metals with different salt solutions.

	Magnesium sulfate solution	Copper(II) sulfate solution	Iron(II) sulfate solution	Number of reactions
Magnesium	Not done	Brown coating	Black coating	2
Copper	No visible reaction	Not done	No visible reaction	0
Iron	Brown coating	No visible reaction	Not done	1

Magnesium is the most reactive metal from these three as it undergoes two reactions, iron is next and copper is the least reactive.

Displacement reactions as redox reactions

Here is the balanced ionic equation for the reaction between magnesium and copper(II) sulfate. The equation below shows the ions present on both sides of the equation:

Mg(s) + Cu²⁺(aq) + SO₄^2-(aq) → Mg²⁺(aq) + SO₄^2-(aq) + Cu(s)

Sulfate ions, SO₄^2- , appear on both sides of the equation because they do not take part in the reaction. The ionic equation for the reaction is written without sulfate ions:

Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s)

Extraction of metals

Ores

Some metals, typically the unreactive ones like gold, are found in the Earth’s crust as pure elements. But most metals are found combined with other elements to form compounds. These compounds are mixed in with other types of rock.

Extraction methods

Different metals are extracted from their ores using different methods depending on their position in the reactivity series.

Electrolysis of molten compounds is used to extract the most reactive metals. In principle, all metals can be extracted using electrolysis but it is expensive.

If a metal is less reactive than carbon, it is cheaper to extract it by heating with carbon.

The very least reactive metals, such as gold, occur as pure elements, but may be contaminated with other elements. They are cleaned through various chemical reactions.

Phytomining (Higher Tier only)

Plants absorb metal ions through their roots in a process called Phytomining. It removes toxic metals from contaminated soil – around old mines for example.

In the future, when supplies of higher grade ores have run out, metals might be extracted by burning the plants to produce ash. The ash would contain a higher concentration of metal than the soil.

1. Plants are used to absorb metal compounds such as copper(II) compounds

2. The plants are harvested, then burned to produce ash, which contains the metal compounds

3. An acid is added to the ash to produce a solution containing dissolved metal compounds (leachate)

4. Copper can be obtained from these solutions by displacement using scrap iron

Phytomining is slow, but it:

• reduces the need to obtain new ore by mining
• conserves limited supplies of more valuable ores with higher metal content