What you'll learn
Oxidation and reduction in terms of electrons and oxygen is a topic that runs through the whole of AQA GCSE Chemistry, connecting metal extraction, reactions of acids, electrolysis, corrosion and the reactivity series. What makes it demanding is that two definitions are in use at once: an older one in terms of oxygen, and a more general one in terms of electron transfer. Both are correct, both are examined, and a question may require either. By the end of this guide you should be able to apply both definitions confidently, use the mnemonic OIL RIG accurately, identify which species is oxidised and which reduced in any reaction, write and balance half equations, recognise a redox reaction from an equation, explain the role of oxidising and reducing agents, and apply all of this to displacement, extraction and electrolysis.
Key terms and definitions
Oxidation — gain of oxygen, or loss of electrons
Reduction — loss of oxygen, or gain of electrons
Redox reaction — a reaction in which oxidation and reduction occur together
Half equation — an equation showing either the loss or the gain of electrons for one species
Oxidising agent — the substance that causes another to be oxidised, and is itself reduced
Reducing agent — the substance that causes another to be reduced, and is itself oxidised
Spectator ion — an ion present in solution that takes no part in the reaction and appears unchanged on both sides
Ionic equation — an equation showing only the species that actually change, with spectator ions omitted
Displacement reaction — a reaction in which a more reactive element takes the place of a less reactive one in a compound
Cathode — the negative electrode in electrolysis, where reduction occurs
Anode — the positive electrode in electrolysis, where oxidation occurs
Core concepts
The two definitions
The original definition concerns oxygen. A substance that gains oxygen is oxidised; a substance that loses oxygen is reduced.
When magnesium burns in air, magnesium gains oxygen to form magnesium oxide, so the magnesium is oxidised.
When copper oxide is heated with carbon, the copper oxide loses its oxygen to leave copper, so the copper oxide is reduced. The carbon gains that oxygen, so the carbon is simultaneously oxidised. Notice that both processes occur in the same reaction, which is why such reactions are called redox.
The more general definition concerns electrons, and it applies to reactions in which no oxygen is present at all. A substance that loses electrons is oxidised; a substance that gains electrons is reduced.
The standard memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain — of electrons. It applies only to the electron definition, so check which definition a question is using before applying it.
The two definitions agree because gaining oxygen involves losing electrons to the oxygen. They are two descriptions of the same underlying process.
Identifying oxidation and reduction from an equation
The reliable method is to track charges.
Consider magnesium reacting with dilute hydrochloric acid. Magnesium atoms start uncharged and end as magnesium ions with a charge of plus 2, so each atom has lost two electrons and magnesium is oxidised. Hydrogen ions start with a charge of plus 1 and end as uncharged hydrogen molecules, so each ion has gained an electron and hydrogen is reduced.
The general rule is that an increase in positive charge means electrons have been lost, so that species is oxidised; a decrease in positive charge, or an increase in negative charge, means electrons have been gained, so that species is reduced.
Because electrons cannot simply disappear, oxidation and reduction always occur together. If you identify one, the other must be present.
Half equations
A half equation shows what happens to one species, with the electrons written explicitly.
For oxidation, the electrons appear on the right, because they are being lost. Magnesium losing two electrons to become a magnesium ion is written with the metal on the left and the ion plus two electrons on the right.
For reduction, the electrons appear on the left, because they are being gained. Two hydrogen ions gaining two electrons to form a hydrogen molecule is written with the ions and the electrons on the left and the molecule on the right.
Two rules govern a correct half equation. The atoms must balance on both sides, and the total charge must balance on both sides. Checking charge is the step candidates most often omit, and it is where the marks are lost.
Adding the two half equations together, with the electrons cancelling, gives the overall ionic equation. The number of electrons lost must equal the number gained, which sometimes requires multiplying one half equation through before adding.
Oxidising and reducing agents
These terms confuse students because each describes a substance by what it does to something else rather than by what happens to it.
An oxidising agent oxidises another substance, which means it takes electrons from it. Taking electrons means the oxidising agent is itself reduced.
A reducing agent reduces another substance, which means it gives electrons to it. Giving electrons means the reducing agent is itself oxidised.
The relationship is always inverted, and stating it clearly is worth practising: the oxidising agent is reduced, and the reducing agent is oxidised.
Common oxidising agents include oxygen, chlorine and other halogens. Common reducing agents include carbon, carbon monoxide, hydrogen and reactive metals.
In the extraction of iron, carbon monoxide acts as the reducing agent: it removes oxygen from iron oxide, reducing the iron oxide, and is itself oxidised to carbon dioxide.
Redox in the reactivity series and displacement
Metals react by losing electrons to form positive ions, which means every reaction of a metal in which it forms a compound is an oxidation of that metal. The reactivity series is therefore an order of how readily each metal is oxidised.
In a displacement reaction, a more reactive metal gives its electrons to the ions of a less reactive metal. When iron is placed in copper sulfate solution, iron atoms lose electrons and are oxidised to iron ions, while copper ions gain those electrons and are reduced to copper atoms. The blue colour fades and brown copper is deposited.
The sulfate ions take no part at all and appear unchanged on both sides, which makes them spectator ions. Removing them gives the ionic equation, which shows the redox process clearly and is what examination questions usually request.
Halogen displacement works the same way with the electron transfer reversed, since halogens gain electrons. A more reactive halogen such as chlorine takes electrons from the ions of a less reactive halogen such as bromide, so the chlorine is reduced and the bromide ions are oxidised.
Redox in electrolysis
Electrolysis is redox driven by an electric current, and the electrode names map directly onto the two processes.
At the cathode, the negative electrode, positive ions gain electrons. Gaining electrons is reduction, so reduction always occurs at the cathode.
At the anode, the positive electrode, negative ions lose electrons. Losing electrons is oxidation, so oxidation always occurs at the anode.
A useful memory aid links the two: reduction and cathode both relate to the negative electrode, and an oxidation occurs at the anode. Candidates who learn this pairing rarely lose the marks.
In the electrolysis of molten lead bromide, lead ions are reduced at the cathode to lead metal, and bromide ions are oxidised at the anode to bromine.
Redox in the reactions of acids
When a metal reacts with an acid to produce a salt and hydrogen, the metal is oxidised as it loses electrons to form its ion, and the hydrogen ions from the acid are reduced as they gain electrons to form hydrogen gas.
By contrast, a neutralisation reaction between an acid and an alkali is not a redox reaction. Hydrogen ions combine with hydroxide ions to form water, and no species changes its charge, so no electrons are transferred. Recognising which acid reactions are redox and which are not is a point examiners test.
Worked examples
Example 1: Applying both definitions (4 marks)
In the reaction between zinc oxide and carbon, zinc metal and carbon monoxide are produced. Identify what is oxidised and what is reduced, using both definitions.
In terms of oxygen, the zinc oxide loses oxygen to become zinc, so the zinc oxide is reduced. The carbon gains oxygen to become carbon monoxide, so the carbon is oxidised.
In terms of electrons, the zinc ions in zinc oxide carry a charge of plus 2 and end as uncharged zinc atoms, so each has gained two electrons and the zinc ions are reduced. The carbon atoms begin uncharged and end combined with oxygen, having lost electrons, so the carbon is oxidised. Carbon is therefore acting as the reducing agent.
Example 2: Writing half equations (4 marks)
Write the half equations for the reaction in which magnesium displaces copper from copper sulfate solution, and identify the oxidising agent.
For the oxidation, each magnesium atom loses two electrons to form a magnesium ion with a charge of plus 2, so the electrons appear on the right-hand side of the half equation.
For the reduction, each copper ion with a charge of plus 2 gains two electrons to form a copper atom, so the electrons appear on the left-hand side.
Both half equations balance for atoms and for charge, and both involve two electrons, so they combine directly. The copper ions are the oxidising agent, because they take electrons from the magnesium and are themselves reduced in the process.
Example 3: Identifying a non-redox reaction (3 marks)
Explain why the reaction between hydrochloric acid and sodium hydroxide is not a redox reaction, whereas the reaction between hydrochloric acid and magnesium is.
In the reaction with sodium hydroxide, hydrogen ions combine with hydroxide ions to form water. The sodium ions and chloride ions are spectator ions. No species changes its charge and therefore no electrons are transferred, so neither oxidation nor reduction occurs.
In the reaction with magnesium, magnesium atoms change from uncharged to a charge of plus 2, losing electrons and being oxidised, while hydrogen ions change from plus 1 to uncharged hydrogen molecules, gaining electrons and being reduced. Since electrons are transferred, this is a redox reaction.
Common mistakes and how to avoid them
The most frequent error is applying OIL RIG to the oxygen definition, which reverses the answer. OIL RIG describes electrons only; gaining oxygen is oxidation.
Students routinely confuse oxidising agents with the substance oxidised. The oxidising agent is the one reduced. Writing the relationship out before answering prevents the error.
In half equations, many candidates balance the atoms but not the charge. Always check that the total charge on the left equals the total charge on the right.
Another common slip is placing the electrons on the wrong side. For oxidation the electrons are produced, so they go on the right; for reduction they are consumed, so they go on the left.
Finally, candidates frequently assume that every reaction involving an acid is redox. Neutralisation is not, because no electrons are transferred.
Exam technique for "Oxidation and reduction in terms of electrons and oxygen"
Decide which definition the question requires before you begin. If the equation contains oxygen and no charges are shown, use the oxygen definition; if ions or charges appear, use the electron definition.
Track charges systematically. Write the charge on each species before and after the reaction, and the direction of electron transfer becomes obvious.
For half equations, write the atoms first, then add electrons to balance the charge, then check both sides again. Working in that fixed order is faster than trying to produce the equation whole.
When asked for an ionic equation, identify and remove the spectator ions explicitly. Marks are often awarded for recognising which ions take no part.
For electrolysis questions, state the electrode, the process and the direction of electron movement together: reduction occurs at the cathode as positive ions gain electrons.
Quick revision summary
Oxidation is gain of oxygen or loss of electrons; reduction is loss of oxygen or gain of electrons, summarised for electrons only as OIL RIG. The two definitions describe the same process and both are examinable. Identify each by tracking charge: an increase in positive charge means electrons lost and oxidation, a decrease means electrons gained and reduction. Oxidation and reduction always occur together, making the reaction redox. Half equations place electrons on the right for oxidation and the left for reduction, and must balance for both atoms and charge; adding them with electrons cancelling gives the ionic equation. The oxidising agent is itself reduced and the reducing agent is itself oxidised, with carbon, carbon monoxide, hydrogen and reactive metals acting as common reducing agents. Metals are oxidised when they react, so the reactivity series ranks ease of oxidation, and in displacement the more reactive metal is oxidised while the less reactive metal's ions are reduced, with spectator ions omitted from the ionic equation. In electrolysis, reduction occurs at the cathode and oxidation at the anode. Neutralisation is not redox, because no electrons are transferred.