RedoxCambridge IGCSE Chemistry: Revision notes

Section 1

What are oxidation and reduction in terms of oxygen transfer?

Redox reactions involve simultaneous oxidation and reduction of substances. The traditional definitions focus on oxygen transfer:

Oxidation = gain of oxygen by a substance
Reduction = loss of oxygen by a substance

These definitions are useful for identifying simple reactions involving combustion and thermal decomposition. For example, when copper is heated in air, it gains oxygen to form copper oxide — this is oxidation of the copper.

However, not all redox reactions involve oxygen directly, so these definitions have limitations and must be extended using electron transfer concepts.

Key termsoxidationreductionredox reactionsoxygen transfer

Example

When magnesium burns in oxygen: 2Mg + O₂ → 2MgO. The magnesium gains oxygen (oxidation), and the oxygen is reduced (gains electrons from Mg). Both processes occur together.

Section 2

How do we define redox reactions using electron transfer?

Electron transfer provides a more complete definition of redox reactions:

Oxidation = loss of electrons (or an increase in oxidation number)
Reduction = gain of electrons (or a decrease in oxidation number)

Every redox reaction involves the transfer of electrons from one substance to another. The substance losing electrons is oxidised; the substance gaining electrons is reduced.

This electron-based definition works for all redox reactions, including those that do not involve oxygen. It is the preferred method for identifying complex redox reactions in ionic solutions and transition metal chemistry.

Key principle: Electrons lost by one substance must equal electrons gained by another substance.

Key termselectron transferoxidation numberoxidisedreduced

Exam tip

Examiners expect you to explain redox using both oxygen transfer AND electron transfer definitions. Use electron transfer when oxygen is not involved; use oxygen transfer for combustion reactions.

Think of it like this

Think of oxidation and reduction like a game of pass-the-parcel: electrons are the parcel. The person who gives it away is oxidised; the person who receives it is reduced.

Section 3

What are oxidation numbers and how do we use them to identify redox reactions?

Oxidation numbers are assigned to elements in compounds according to strict rules. They allow us to track electron transfer and identify redox reactions:

Rules for assigning oxidation numbers:

The oxidation number of an element in its uncombined (elemental) state is zero
The oxidation number of a monatomic ion equals the charge on that ion
The sum of oxidation numbers in a neutral compound is zero
The sum of oxidation numbers in an ion equals the charge on that ion
Oxygen usually has oxidation number −2 (except in peroxides where it is −1)
Hydrogen usually has oxidation number +1 (except in metal hydrides where it is −1)

Identifying redox reactions using oxidation numbers:

A redox reaction occurs when the oxidation numbers of elements change during the reaction. If oxidation numbers stay the same, it is not a redox reaction.

Roman numerals are used to indicate oxidation numbers: Fe²⁺ is written Fe(II), Fe³⁺ is written Fe(III), and so on.

Key termsoxidation numbersmonatomic ionuncombined state

Example

In the reaction 2Fe + 3Cl₂ → 2FeCl₃: Fe goes from 0 (uncombined) to +3 (in FeCl₃), so Fe is oxidised. Cl goes from 0 to −1, so Cl is reduced. This is a redox reaction.

Exam tip

Always assign oxidation numbers to every element in the reactants and products, then check which ones change. Only elements with changed oxidation numbers are involved in redox.

Common mistake

Students often forget that in FeCl₃, each Cl has oxidation number −1, so the sum is: Fe(+3) + 3×(−1) = 0. Do not assign all three chlorines' charges to a single number.

Section 4

How do we identify redox reactions using colour changes?

Certain colour changes in solution provide visual evidence of redox reactions:

Potassium manganate(VII) [KMnO₄] — acidified aqueous solution:

Colour: Deep purple
In a redox reaction: The purple colour fades or disappears (the MnO₄⁻ ion is reduced to Mn²⁺, which is colourless)
When it fades: A redox reaction is occurring; the substance being tested is a reducing agent (it causes the manganate to be reduced)
If colour persists: No redox reaction with that substance

Potassium iodide [KI] — aqueous solution:

Colour in neutral solution: Colourless (contains I⁻ ions)
In a redox reaction: If an oxidising agent is present, I⁻ is oxidised to I₂ (iodine)
Colour change: Solution turns brown or blue-black (when starch is added, as iodine forms a complex with starch)
When colour develops: A redox reaction is occurring; the substance being tested is an oxidising agent

These colour tests are qualitative methods for detecting redox reactions and identifying the type of substance (oxidising or reducing agent).

Key termspotassium manganate(VII)potassium iodidecolour changequalitative test

Example

When hydrogen peroxide is added to acidified potassium manganate(VII), the purple colour fades immediately. This shows hydrogen peroxide is a reducing agent and is being oxidised in a redox reaction.

Exam tip

In exam questions, you will be asked to identify redox based on colour changes. Remember: purple fades = KMnO₄ test (reducing agent present); brown develops = KI test (oxidising agent present).

Section 5

What are oxidising and reducing agents?

In any redox reaction, two key substances play different roles:

Oxidising agent:

A substance that oxidises another substance (causes another substance to lose electrons)
The oxidising agent is itself reduced (it gains electrons)
It contains an element whose oxidation number decreases
Examples: Oxygen, chlorine, potassium manganate(VII), acidified potassium dichromate

Reducing agent:

A substance that reduces another substance (causes another substance to gain electrons)
The reducing agent is itself oxidised (it loses electrons)
It contains an element whose oxidation number increases
Examples: Hydrogen, carbon, metals, hydrogen peroxide (as a reducing agent)

Key relationship:

	Oxidising Agent	Reducing Agent
Effect on other substance	Oxidises it	Reduces it
What happens to agent	Gets reduced	Gets oxidised
Oxidation number change	Decreases	Increases
Electrons	Gains electrons	Loses electrons

Note that a substance can be both an oxidising and reducing agent in different reactions (e.g., hydrogen peroxide). In any single reaction, identify which role each substance plays.

Key termsoxidising agentreducing agentoxidation number decreaseoxidation number increase

Example

In the reaction: Zn + CuSO₄ → ZnSO₄ + Cu. Zn goes from 0 to +2 (oxidation), so Zn is the reducing agent. Cu²⁺ goes from +2 to 0 (reduction), so CuSO₄ (or Cu²⁺) is the oxidising agent.

Common mistake

Students confuse oxidising agents with oxidised substances. Remember: the oxidising agent causes oxidation in another substance, but the oxidising agent itself is reduced.

Must Know

Redox reactions involve simultaneous oxidation and reduction; the substance oxidised loses electrons (or gains oxygen), and the substance reduced gains electrons (or loses oxygen)
Oxidation numbers allow identification of redox reactions: assign numbers using the rules, then check if any element's number changes; if it does, it is a redox reaction
Electron transfer is the core definition: oxidation = loss of electrons (oxidation number increases), reduction = gain of electrons (oxidation number decreases)
Colour changes identify redox: acidified potassium manganate(VII) turns colourless (reducing agent present), and potassium iodide turns brown (oxidising agent present)
Oxidising agents are reduced (gain electrons, oxidation number decreases); reducing agents are oxidised (lose electrons, oxidation number increases)
Use Roman numerals to denote oxidation numbers (e.g., Fe(II), Fe(III)) in written answers

That's the notes covered.

Carry on to the next subtopic.