The Reactivity Series and Displacement

When a metal reacts with oxygen, the oxygen atoms join to the metal atoms to make a metal oxide. Because the metal has gained oxygen, the reaction is described as oxidation. This is the first, simplest way to define oxidation, and it works for every metal-plus-oxygen reaction you will meet.

For example, magnesium burns in air with a bright white flame to form magnesium oxide:

$2Mg+O_2\to 2MgO$

The magnesium has gained oxygen, so the magnesium has been oxidised. Copper heated in air behaves the same way, slowly forming black copper oxide:

$2Cu+O_2\to 2CuO$

Oxidation is the gain of oxygen. Reduction is the loss of oxygen. A memory hook: the metal that ends up joined to oxygen has been oxidised.

The reverse process, in which a metal oxide loses its oxygen, is called reduction. Reduction becomes important when metals are extracted from their ores, which is the topic of the next lesson.

Metals react by losing electrons to form positive ions. When magnesium reacts, each atom loses two electrons to become an $M{g}^{2+}$ ion. The easier it is for a metal to lose electrons and form its positive ion, the more reactive that metal is.

This single idea explains the whole reactivity series. Potassium loses its outer electron very easily, so it reacts violently. Copper holds onto its electrons far more tightly, so it barely reacts at all.

$Mg\to M{g}^{2+}+2e^{-}$

$K\to K^{+}+e^{-}$

The more readily a metal forms a positive ion, the more reactive it is. Reactivity is really a measure of how strongly a metal wants to give away electrons.

This is why the metals near the top of the reactivity series are the Group 1 and Group 2 metals: their outer electrons are lost most easily. Writing the ionic half equation, as shown above, is a clear way to show which ion a metal forms.

The reactivity series lists metals from most to least reactive. You must be able to recall this order for the eight metals AQA specifies. The non-metals carbon and hydrogen are included because they are used as reference points, especially when deciding how a metal is extracted.

Element	Symbol	Reactivity	Notes
Potassium	K	Most reactive	Group 1
Sodium	Na		Group 1
Lithium	Li		Group 1
Calcium	Ca		Group 2
Magnesium	Mg		Group 2
Carbon	C	(non-metal reference)	Used to extract metals below it
Zinc	Zn
Iron	Fe
Hydrogen	H	(non-metal reference)	Metals above it react with acids
Copper	Cu	Least reactive	Found almost native

A common way to remember the metals is a phrase such as "Please Stop Little Cats Messing Zebras In Cages" for K, Na, Li, Ca, Mg, Zn, Fe, Cu. Carbon sits above zinc (and above iron) and hydrogen sits just above copper.

Learn the exact position of carbon and hydrogen. Carbon above a metal means that metal can be extracted by heating with carbon; a metal above hydrogen means it reacts with dilute acids.

Metals near the top of the series react with cold water to produce a metal hydroxide and hydrogen gas. The reaction becomes gentler as you move down the series, and the metals from zinc downwards do not react noticeably with cold water at all.

All these reactions are carried out at room temperature. Reactions with steam are not part of this specification, so describe only the room-temperature reaction with liquid water.

Metal	Reaction with cold water
Potassium	Very vigorous; ignites, fizzes and moves fast
Sodium	Vigorous; melts into a ball, fizzes rapidly
Lithium	Steady fizzing, floats and moves
Calcium	Steady stream of bubbles
Magnesium	Very slow with cold water
Zinc, iron, copper	No noticeable reaction with cold water

For example, sodium reacts with water to give sodium hydroxide and hydrogen:

$2Na+2H_{2}O\to 2NaOH+H_2$

The general pattern is: metal + water → metal hydroxide + hydrogen. The fizzing you see is the hydrogen gas being released.

Any metal placed above hydrogen in the reactivity series reacts with a dilute acid to give a salt and hydrogen gas. The more reactive the metal, the faster the fizzing. Copper sits below hydrogen and does not react with dilute acids.

The general pattern is:

$\text{metal}+\text{acid}\to \text{salt}+\text{hydrogen}$

You can rank metals by how quickly they bubble. Placing magnesium, zinc and iron each into dilute hydrochloric acid gives:

Metal	Rate of fizzing with dilute acid	Position
Magnesium	Fast, vigorous bubbling	Most reactive
Zinc	Moderate bubbling	Middle
Iron	Slow bubbling	Least reactive
Copper	No reaction	Below hydrogen

For example, magnesium reacts with dilute hydrochloric acid:

$Mg+2HCl\to MgC{l}_2+H_2$

Comparing the speed of fizzing lets you deduce a reactivity order from experimental results — a skill the exam tests directly. The faster the reaction, the more reactive the metal.

A more reactive metal displaces a less reactive metal from its compound. The more reactive metal takes the place of the less reactive one, which is pushed out as a pure element. This happens because the more reactive metal forms its positive ion more readily.

Worked example. An iron nail is placed into blue copper(II) sulfate solution. Iron is more reactive than copper, so iron displaces the copper:

$Fe+CuS{O}_4\to FeS{O}_4+Cu$

What you observe: the blue solution fades as iron sulfate (pale green) forms, and a red-brown coating of copper builds up on the nail. The iron has taken the place of copper in the compound.

The reverse does not happen: putting copper into iron sulfate solution gives no reaction, because copper is less reactive than iron and cannot displace it. Displacement reactions can therefore be used to work out a reactivity order — if metal A displaces metal B, then A is more reactive than B.

If a metal is added to the solution of a salt and a reaction happens, the added metal is more reactive than the metal in the salt. No reaction means it is less reactive.

1. Getting oxidation and reduction the wrong way round

Oxidation is the gain of oxygen; reduction is the loss of oxygen. When a metal reacts with oxygen to form an oxide, the metal is oxidised because it gains oxygen. Do not call this reduction.

2. Forgetting where carbon and hydrogen sit

Carbon and hydrogen are non-metals but appear in the reactivity series as references. Carbon sits above zinc and iron; hydrogen sits just above copper. Their positions decide extraction method and whether a metal reacts with acids.

3. Describing reactions with steam

The specification limits metal reactions with water to cold water at room temperature. Do not bring in reactions with steam; describe the cold-water reaction only.

4. Saying reactivity is about "how hard the metal is" or its melting point

Reactivity is about how readily a metal loses electrons to form a positive ion. It is not about hardness, density or melting point. Explain reactivity in terms of ion formation.

5. Writing the wrong products for metal plus acid

Metal + acid gives a salt + hydrogen, not "metal oxide" or water. The gas released is hydrogen, which you can test with a lit splint that gives a squeaky pop.

6. Claiming copper reacts with dilute acid

Copper is below hydrogen in the reactivity series, so it does not react with dilute hydrochloric or sulfuric acid. Only metals above hydrogen give a salt and hydrogen with dilute acids.