Last updated: 15 October 2026
Redox reactions appear throughout A-Level Chemistry. They occur when metals react with ions, during electrolysis, inside electrochemical cells, and in reactions involving familiar oxidising agents such as acidified manganate(VII) or dichromate(VI) ions. Although these reactions may look very different, they are connected by one central idea: electrons are being transferred.
Sometimes that transfer is obvious because electrons can be written directly in half-equations. In other reactions, especially those involving covalent substances, the electrons are not shown explicitly. We therefore use oxidation numbers as a bookkeeping system to track which atoms have effectively become more oxidised or more reduced.
Try these five questions before reading the chapter. Focus on electron transfer, oxidation numbers and the roles of oxidising and reducing agents.
- 5 diagnostic questions
- Oxidation-number reasoning
- Instant explanation for every answer option
What Is a Redox Reaction?
A redox reaction is a reaction in which oxidation and reduction occur simultaneously. The word redox simply combines reduction and oxidation.
The most fundamental way to understand redox is through electrons. If one species loses electrons, those electrons must be accepted by another species. This is why oxidation cannot occur on its own and reduction cannot occur on its own.
Oxidation is loss of electrons. Reduction is gain of electrons.
Using OIL RIG
A useful mnemonic is OIL RIG:
OIL — Oxidation Is Loss: a species is oxidised when it loses electrons.
RIG — Reduction Is Gain: a species is reduced when it gains electrons.
The mnemonic is useful, but it is better to understand what the equations actually mean. Consider zinc atoms forming zinc ions:
The zinc atom releases two electrons. Zinc has therefore been oxidised.
The copper(II) ion accepts two electrons. Copper(II) ions have therefore been reduced.
Putting the two half-equations together
If zinc metal is placed in a solution containing Cu²⁺ ions, the electrons released by zinc are accepted by Cu²⁺ ions.
Zinc loses two electrons: Zn → Zn²⁺ + 2e⁻, so zinc is oxidised. Copper(II) ions gain those two electrons: Cu²⁺ + 2e⁻ → Cu, so Cu²⁺ is reduced. The number of electrons lost must equal the number gained.
Oxidation Number: Tracking Redox When Electrons Are Not Shown
Electron transfer is particularly easy to see when ions are involved. However, many redox equations do not contain explicit electrons. We therefore need another way to keep track of oxidation and reduction.
An oxidation number, also called an oxidation state, is a number assigned to an atom according to a set of rules. It allows us to track the effective redistribution of electrons during a reaction.
If the oxidation number of an element increases, that element is oxidised. If its oxidation number decreases, that element is reduced.
Rules for assigning oxidation numbers
Oxidation numbers are assigned systematically. In most A-Level questions, only a small number of rules are needed.
| Species or element | Usual oxidation number | Example |
|---|---|---|
| Element in its uncombined state | 0 | Na, O₂, Cl₂, S₈: 0 |
| Monatomic ion | Equal to ionic charge | Na⁺: +1; Mg²⁺: +2; Cl⁻: −1 |
| Group 1 metals in compounds | +1 | Na in NaCl: +1 |
| Group 2 metals in compounds | +2 | Mg in MgO: +2 |
| Fluorine in compounds | −1 | F in HF: −1 |
| Oxygen in most compounds | −2 | O in H₂O: −2 |
| Oxygen in peroxides | −1 | O in H₂O₂: −1 |
| Hydrogen in most compounds | +1 | H in H₂O: +1 |
| Hydrogen in metal hydrides | −1 | H in NaH: −1 |
The sum rule
Once the familiar oxidation numbers have been assigned, the remaining oxidation number can usually be found using the overall charge.
Neutral compound: the sum of all oxidation numbers is 0.
Polyatomic ion: the sum of all oxidation numbers equals the overall charge of the ion.
Example: oxidation number of sulfur in SO₂
Oxygen normally has oxidation number −2. Let the oxidation number of sulfur be x.
The oxidation number of sulfur in SO₂ is therefore +4.
Example: oxidation number of sulfur in SO₄²⁻
The sulfate ion has an overall charge of 2−. Again, oxygen normally has oxidation number −2.
Sulfur therefore has oxidation number +6 in SO₄²⁻.
Oxidation number is not the same as ionic charge
An oxidation number is a bookkeeping value assigned to an individual atom. It does not necessarily mean that the atom actually exists as a free ion carrying that charge.
For example, sulfur has oxidation number +6 in SO₄²⁻, but the sulfate ion does not contain an isolated S⁶⁺ ion. The bonding within sulfate is covalent, and +6 is the oxidation number assigned to sulfur for electron bookkeeping.
Oxidation numbers are written with the sign first, such as +2 or −1. Ionic charges are conventionally written with the number first, such as Fe²⁺ or O²⁻.
Identifying Oxidation and Reduction Using Oxidation Numbers
Once oxidation numbers have been assigned, a redox reaction can be recognised by comparing the oxidation number of each relevant element before and after the reaction.
| Change | Process | Electron interpretation |
|---|---|---|
| Oxidation number increases | Oxidation | Loss of electrons |
| Oxidation number decreases | Reduction | Gain of electrons |
Example: zinc reacting with copper(II) ions
Begin by assigning oxidation numbers.
Zinc: Zn changes from 0 in Zn(s) to +2 in Zn²⁺. Its oxidation number increases, so zinc is oxidised.
Copper: Cu changes from +2 in Cu²⁺ to 0 in Cu(s). Its oxidation number decreases, so Cu²⁺ is reduced.
The oxidation-number changes tell exactly the same story as the electron half-equations. Zinc has effectively lost two electrons while Cu²⁺ has gained two electrons.
Example: iron(II) ions reacting with chlorine
Iron changes from +2 to +3. Its oxidation number increases, so Fe²⁺ is oxidised.
Chlorine changes from 0 in Cl₂ to −1 in Cl⁻. Its oxidation number decreases, so chlorine is reduced.
A useful examination habit is to write the relevant oxidation numbers directly above or beside the atoms that change. You usually do not need to assign oxidation numbers to every atom in the equation — concentrate first on the elements whose oxidation states may have changed.
Oxidising Agents and Reducing Agents
Students often understand oxidation and reduction but then reverse the names of the agents. The key is to remember that an agent is named according to what it causes another species to do.
Oxidising agent: causes another species to be oxidised by accepting electrons. The oxidising agent is itself reduced.
Reducing agent: causes another species to be reduced by donating electrons. The reducing agent is itself oxidised.
Oxidising agent → gains electrons → is reduced.
Reducing agent → loses electrons → is oxidised.
Identifying the agents from an equation
Return to the reaction between zinc and copper(II) ions:
Zinc is oxidised from oxidation number 0 to +2. It supplies electrons to Cu²⁺ ions and therefore acts as the reducing agent.
Cu²⁺ is reduced from +2 to 0. It accepts electrons from zinc and therefore acts as the oxidising agent.
Another example
Fe²⁺ changes from +2 to +3 and is oxidised. Fe²⁺ therefore acts as the reducing agent.
Chlorine changes from 0 to −1 and is reduced. Cl₂ therefore acts as the oxidising agent.
| Species | What happens to it? | Role |
|---|---|---|
| Loses electrons | Oxidised | Reducing agent |
| Gains electrons | Reduced | Oxidising agent |
Connecting Electron Transfer with Oxidation-Number Changes
The electron-transfer definition and the oxidation-number definition are not two unrelated ways of describing redox. They are different ways of tracking the same underlying process.
Electron-transfer view: oxidation means electrons are lost, while reduction means electrons are gained.
Oxidation-number view: oxidation means the oxidation number increases, while reduction means the oxidation number decreases.
Why does loss of electrons increase oxidation number?
Electrons carry negative charge. When an atom is considered to have lost electron density during oxidation, its oxidation number becomes more positive. For example:
The oxidation number increases from +2 to +3 as one electron is lost.
Why does gain of electrons decrease oxidation number?
Gaining negatively charged electrons makes the oxidation number less positive or more negative.
Each chlorine atom changes from oxidation number 0 to −1. The oxidation number decreases because chlorine gains electrons.
Increase in oxidation number = oxidation = electron loss.
Decrease in oxidation number = reduction = electron gain.
A Reliable Method for Analysing Redox Equations
When an examination question asks you to identify what has been oxidised, reduced, or which species acts as the oxidising or reducing agent, use a consistent sequence.
Step 1: Assign the relevant oxidation numbers
Focus on elements that could reasonably change oxidation state. Free elements begin at 0, while ions and familiar atoms such as oxygen and hydrogen can usually be assigned immediately.
Step 2: Compare reactants and products
Look for increases and decreases in oxidation number.
Step 3: Identify the agents
The species that is reduced is the oxidising agent. The species that is oxidised is the reducing agent.
For 2Fe²⁺ + Cl₂ → 2Fe³⁺ + 2Cl⁻: Fe changes from +2 to +3, so Fe²⁺ is oxidised and is the reducing agent. Cl changes from 0 to −1, so Cl₂ is reduced and is the oxidising agent.
Common Mistakes
| Common mistake | Why it is wrong | Better approach |
|---|---|---|
| Thinking that a decrease in oxidation number means oxidation | Oxidation involves loss of electrons, which corresponds to an increase in oxidation number. | Remember: increase → oxidation; decrease → reduction. |
| Calling the species that is oxidised the oxidising agent | An agent is named according to what it does to another species. | The species oxidised is the reducing agent; the species reduced is the oxidising agent. |
| Assuming oxygen always has oxidation number −2 | Oxygen has oxidation number −1 in peroxides such as H₂O₂. | Use −2 normally, but check whether the compound is a peroxide. |
| Treating oxidation number as an actual ionic charge | Oxidation number is a formal bookkeeping value and does not mean that the atom exists as a free ion of that charge. | Use oxidation numbers to track redox changes, while keeping ionic charge and chemical bonding conceptually separate. |
What you should be able to do
Recognise electron transfer
Identify oxidation as electron loss and reduction as electron gain.
Assign oxidation numbers
Apply the standard oxidation-number rules correctly to elements, compounds and ions.
Use the sum rule
Calculate an unknown oxidation number using the overall charge of a compound or ion.
Identify oxidation and reduction
Recognise oxidation from an increase and reduction from a decrease in oxidation number.
Identify oxidising agents
Recognise that an oxidising agent accepts electrons and is itself reduced.
Identify reducing agents
Recognise that a reducing agent donates electrons and is itself oxidised.
These ten questions test whether you can apply electron-transfer and oxidation-number reasoning to familiar and unfamiliar redox examples.
- 10 mastery questions
- Oxidation states and redox agents
- Instant explanation for every answer option
Frequently asked questions
Questions Students Often Ask About Redox Processes
Redox becomes much easier once electron transfer, oxidation numbers and the names of the agents are treated as different descriptions of the same process. These questions clear up the distinctions that most often cause confusion.
Is oxidation always a reaction with oxygen?
No. Reactions with oxygen are often oxidation reactions, but the modern definition is broader. Oxidation is loss of electrons or an increase in oxidation number. For example, Fe²⁺ → Fe³⁺ + e⁻ is oxidation even though oxygen does not appear in the equation.
Why must oxidation and reduction always occur together?
Electrons cannot simply disappear. If one species loses electrons, another species must accept them. The electron loss is oxidation and the corresponding electron gain is reduction, so the two processes occur together in a redox reaction.
Is an oxidation number the same as the charge on an atom?
Not necessarily. For a monatomic ion such as Fe³⁺, the oxidation number does equal the ionic charge. In a covalent compound or polyatomic ion, however, oxidation number is a formal bookkeeping value. Sulfur having oxidation number +6 in SO₄²⁻ does not mean that sulfate contains a free S⁶⁺ ion.
How do I stop confusing oxidising agents and reducing agents?
Concentrate first on what happens to the agent itself. An oxidising agent gains electrons and is reduced. A reducing agent loses electrons and is oxidised. You can then check the name: the oxidising agent causes the other species to be oxidised, while the reducing agent causes the other species to be reduced.
Is oxygen always −2 and hydrogen always +1?
Those are their usual oxidation numbers, but there are important exceptions. Oxygen is −1 in peroxides such as H₂O₂, while hydrogen is −1 in metal hydrides such as NaH. Recognising these exceptions prevents incorrect oxidation-number calculations.
Redox questions become far more systematic once you connect the three viewpoints: electron transfer, changes in oxidation number, and the roles of oxidising and reducing agents. Track what happens to the electrons or oxidation numbers first, and the identity of each agent then follows naturally.