Metal Oxides & Oxidation
Most metals react with oxygen to form a metal oxide. How quickly and how vigorously they do this is our first clue to how reactive a metal is — magnesium burns with a brilliant white flame, iron slowly rusts, and gold never tarnishes at all.
metal + oxygen → metal oxide
For example, magnesium burns in air to make magnesium oxide:
2Mg + O2 → 2MgO
Oxidation is the gain of oxygen. Reduction is the loss of oxygen.
When a metal reacts with oxygen it gains oxygen, so the metal is oxidised. Forming a metal oxide is always an oxidation reaction.
A handy way to remember which way round it goes:
- A substance that gains oxygen has been OXIDISED.
- A substance that loses oxygen has been REDUCED.
So when copper is heated in air it is oxidised to black copper oxide:
2Cu + O2 → 2CuO
This oxygen definition is all that Foundation Tier needs. Higher Tier students also learn to describe oxidation and reduction in terms of electrons — that comes up in Displacement Reactions and again in electrolysis.
Discovery activity
How did chemists decide which metal is most reactive?
Below are eight key metals used to build the experimental order, displayed in a random order. This is the challenge that faced early chemists — by looking at them you cannot tell which is most reactive. How would you find out?
You need to do experiments and compare how vigorously each metal reacts. The more vigorous the reaction, the more reactive the metal.
Two experiments give us the evidence we need. The length of each bar shows how vigorously the metal reacts — the longer the bar, the more reactive.
Experiment 1 — reactions with cold water
Experiment 2 — reactions with dilute acid
K, Na, Li and Ca are too dangerously reactive to test with acid — water already tells us their order.
Water separates the most reactive metals; dilute acid separates the middle ones. Together, the two experiments give us enough information to rank all of them.
Combining the evidence from both experiments — and adding the non-metals carbon and hydrogen as reference points — lets us build the reactivity series:
A few more metals — aluminium, silver and gold — slot into the full series too. See the complete diagram below.
🧪 Exam-style questions
Complete the sentence. Oxidation is the… Tick (✓) one box.
Magnesium burns in air to form magnesium oxide. Explain, in terms of oxygen, why this is an oxidation reaction.
Show answer
- The magnesium gains oxygen (it combines with oxygen from the air to form magnesium oxide). 1 mark
- Gaining oxygen is oxidation, so the magnesium has been oxidised. 1 mark
The balanced equation is 2Mg + O2 → 2MgO.
In which change has a substance been reduced? Tick (✓) one box.
The Reactivity Series
The reactivity series is a league table of metals, listed in order from the most reactive at the top to the least reactive at the bottom. We build it by comparing how metals react with water and with dilute acid.
When a metal reacts, its atoms lose electrons to form positive ions. The reactivity of a metal is its tendency to form positive ions — the more easily a metal loses its outer electrons, the more reactive it is.
Reactions with water
The most reactive metals react with cold water to give a metal hydroxide (an alkaline solution) and hydrogen gas:
metal + water → metal hydroxide + hydrogen
For example, sodium fizzes across the surface of water:
2Na + 2H2O → 2NaOH + H2
Potassium, sodium and lithium (the alkali metals) and calcium all react with cold water. How vigorously they fizz tells us their order: potassium reacts most violently, then sodium, then lithium, then calcium more gently. Less reactive metals react only very slowly — or not at all — which makes water a poor test for telling them apart.
At GCSE these reactions are limited to room temperature: you are not expected to know the reactions of metals with steam.
Pick a metal to drop onto the water and watch how vigorously it reacts.
Reactions with dilute acid
To compare the metals that barely react with water, we use dilute acid instead. Metals more reactive than hydrogen react with acid to give a salt and hydrogen, and the rate of fizzing (effervescence of hydrogen) shows how reactive each one is:
metal + acid → salt + hydrogen
Magnesium fizzes rapidly, zinc steadily, iron slowly, and copper not at all — placing them in the order Mg > Zn > Fe > Cu.
Each reactive metal gives off hydrogen, which fizzes off as bubbles — the faster the fizzing, the more reactive the metal (Mg > Zn > Fe > Cu). Copper lies below hydrogen in the reactivity series, so it does not react with dilute acid at all.
The reactivity series. The non-metals carbon and hydrogen are included as reference points — their positions decide how a metal can be extracted and whether it reacts with acid.
Carbon and hydrogen are non-metals, but they are placed in the series as reference points. A metal’s position relative to carbon decides how it is extracted (see §4), and its position relative to hydrogen decides whether it reacts with acids (see §5).
🧪 Exam-style questions
The reactivity of a metal is best described as its tendency to…
A student adds four metals to dilute hydrochloric acid. Metal W fizzes very fast, X fizzes slowly, Y fizzes steadily and Z does not react. What is the order of reactivity, most reactive first?
Displacement Reactions
The core idea here is for everyone; the ionic and half equations at the end are Higher Tier only.
A displacement reaction is one in which a more reactive metal displaces a less reactive metal from a compound (usually a solution of its salt).
If you put a more reactive metal into a solution of a less reactive metal’s salt, the more reactive metal “pushes out” the less reactive one and takes its place. For example, magnesium is more reactive than copper, so it displaces copper from copper sulfate solution:
magnesium + copper sulfate → magnesium sulfate + copper
Mg + CuSO4 → MgSO4 + Cu
You can see it happen: the blue colour of the copper sulfate solution fades, and a coating of brown copper forms on the magnesium. If the added metal is less reactive than the one in the compound, nothing happens — silver cannot displace copper, for instance.
Use the reactivity series: the reaction only happens if the metal you add is higher than the metal in the compound. Displacement is the basis of the whole reactivity series — ranking metals by which can displace which gives exactly the same order as ranking them by their reactions with water and acid.
🧪 Test a displacement reaction
Pick a salt solution, then a metal to drop in. Both metals light up in the reactivity series so you can compare them — then see whether a reaction happens.
Ionic equations and redox Higher
In a displacement reaction, electrons are transferred from one metal to the other. One metal is oxidised (loss of electrons) while the other is reduced (gain of electrons), so it is a redox reaction (reduction and oxidation happening in one reaction). Higher Tier defines these in terms of electrons:
Oxidation Is Loss of electrons · Reduction Is Gain of electrons.
Let’s build the ionic equation for magnesium displacing copper from copper(II) sulfate solution, one step at a time.
Step 1 — Write the normal balanced equation. Magnesium is more reactive than copper, so it takes copper’s place in the compound:
Mg(s) + CuSO4(aq) → MgSO4(aq) + Cu(s)
Step 2 — Split the dissolved ionic compounds into their ions. Any ionic compound that is dissolved in water (aq) has had its lattice broken up, so its ions are pulled apart and float about separately in the solution. We rewrite each aqueous ionic compound as its free ions. (The magnesium and copper metals are solid elements, not in solution, so they stay exactly as they are.)
Mg(s) + Cu2+(aq) + SO42−(aq) → Mg2+(aq) + SO42−(aq) + Cu(s)
Step 3 — Cancel the spectator ions. The sulfate ion (SO42−) is identical on both sides — it starts and ends as the same free aqueous ion and takes no part in the reaction. An ion like this is called a spectator ion, and we cancel it from both sides:
Mg(s) + Cu2+(aq) + SO42−(aq) → Mg2+(aq) + SO42−(aq) + Cu(s)
What is left is the ionic equation — it shows only the species that actually change:
Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s)
Half equations — oxidation and reduction separately Higher
A half equation goes one step further: it shows the oxidation or reduction of just one element on its own, including the electrons that are transferred. We can split the ionic equation above into its two halves.
Start with “what becomes what”. Magnesium atoms turn into magnesium ions:
Mg → Mg2+
Now finish the half equation with a two-point checklist:
- Balance the atoms. One Mg on each side — already balanced.
- Balance the charge using electrons. The left side has a total charge of 0; the right side is 2+. Each electron is 1−, so add electrons to the more positive side until both sides match — here, 2 electrons on the right:
Mg → Mg2+ + 2e− (oxidation — magnesium loses electrons)
The copper does the opposite — copper ions gain those electrons to become copper atoms:
Cu2+ + 2e− → Cu (reduction — copper ions gain electrons)
The more reactive metal is always oxidised (it loses electrons to become an ion); the less reactive metal ion is always reduced (it gains electrons to become an atom). The electrons lost by one are exactly the electrons gained by the other.
Marks are lost on ionic and half equations even when the chemistry is understood. Watch for: impossible species such as Zn2− (a metal forms a positive ion, Zn2+); writing full formulae instead of separate ions (use Cu2+, not CuSO4); a charge that does not balance on the two sides; and missing or incorrect state symbols. For zinc displacing copper the full mark answer is Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s). Check in order: species → balancing → charge → state symbols.
🧪 Build a half equation
Pick a displacement reaction and one of its metals, then build that metal’s half equation: choose which side the atom and the ion go on, and add electrons until the atoms and charges balance.
🧪 Exam-style questions
A piece of zinc is left in blue copper(II) sulfate solution. Describe what you would see.
Show answer
- A red-brown (copper) coating / deposit forms on the zinc. 1 mark
- The blue colour of the solution fades (goes paler / colourless). 1 mark
Zinc is more reactive than copper, so it displaces copper: Zn + CuSO4 → ZnSO4 + Cu.
Balance the equation for zinc displacing silver from silver nitrate solution.Type a balancing number in each box (leave it as 1 if no number is needed), then press Check. Any correct set of numbers is accepted.
Show answer
Zn + 2AgNO3 → Zn(NO3)2 + 2Ag
Zinc forms a 2+ ion, so it needs two nitrate groups — which means two AgNO3 and two Ag. 1 mark for the fully balanced equation.
Zinc displaces copper from copper(II) sulfate solution. Write the two half equations for this reaction, and state which species is oxidised.
Show answer
- Oxidation: Zn → Zn2+ + 2e− — the zinc is oxidised (it loses electrons). 1 mark
- Reduction: Cu2+ + 2e− → Cu. 1 mark
Accept the electrons written as e or e−. Both half equations must have the charges balanced by the two electrons.
Extracting Metals: Reduction with Carbon
A few very unreactive metals — gold, silver and platinum — are found in the Earth as the metal itself (“native” metals), because they are so unreactive they have stayed uncombined for billions of years. Most metals, though, are found locked inside compounds in rocks called ores, and have to be chemically “pulled apart” to get the metal out.
An ore is a rock that contains enough of a metal (or its compound) to make it economical to extract.
Most metals are found as compounds inside ores. Crushing and concentrating the ore is physical — the metal is only freed by a chemical reduction step, using carbon or electrolysis.
Many metals occur as oxides. To get the metal, the oxygen has to be removed — in other words the oxide must be reduced. How we do this depends on the metal’s position relative to carbon in the reactivity series.
A metal less reactive than carbon can be extracted from its oxide by reduction with carbon. The carbon displaces the metal because it is more reactive, taking the oxygen for itself.
A metal more reactive than carbon (aluminium, magnesium, calcium, sodium, potassium) cannot — it has to be extracted by electrolysis instead.
The same reactivity series, regrouped by extraction method. A metal’s position relative to carbon decides how it is extracted: more reactive → electrolysis; less reactive → reduction with carbon; unreactive → found native as the element.
For example, iron oxide is heated with carbon in a blast furnace. The carbon gains the oxygen (it is oxidised) and the iron loses it (it is reduced):
2Fe2O3 + 3C → 4Fe + 3CO2
The same happens with copper oxide and lead oxide:
2CuO + C → 2Cu + CO2
2PbO + C → 2Pb + CO2
| In a reduction-with-carbon reaction… | Gains / loses oxygen | Oxidised or reduced? |
|---|---|---|
| The metal oxide | loses oxygen | reduced |
| The carbon | gains oxygen | oxidised |
🧪 Exam-style questions
Which of these metals cannot be extracted by heating its oxide with carbon?
In the reaction 2CuO + C → 2Cu + CO2, name the substance that is reduced and explain your choice.
Show answer
- Copper oxide is reduced. 1 mark
- It has lost oxygen (to the carbon), and loss of oxygen is reduction. 1 mark