Redox

Q: How do I tell oxidising and reducing agents apart?

Remember the rule: Oxidising agents are reduced, reducing agents are oxidised. For Extended tier, you can also check oxidation state changes: oxidising agents have decreasing oxidation states, reducing agents have increasing oxidation states.

CIE IGCSE Chemistry· 6.4 Redox Reactions· 50 min read

1. Core: Redox as Oxygen Transfer★★☆☆☆⏱ 10 min

📘 Definition

Redox Reaction (Core)

A reaction where oxidation (gain of oxygen) and reduction (loss of oxygen) occur at the same time. The substance that gains oxygen is oxidised, the substance that loses oxygen is reduced.

Example:

Reaction of copper(II) oxide with hydrogen: $C u O + H_{2} \to C u + H_{2} O$ (CuO is reduced, H2 is oxidised)

For Core tier, you only need to use oxygen transfer to identify redox reactions, oxidising agents and reducing agents. The oxidising agent gives away oxygen (it is reduced), while the reducing agent takes oxygen (it is oxidised).

📐 Worked Example

Identify the oxidised substance, reduced substance, oxidising agent and reducing agent in the reaction: $F e_{2} O_{3} + 3 C O \to 2 F e + 3 C O_{2}$

1
Check which substance gains oxygen: CO gains oxygen to form CO2, so CO is oxidised, making it the reducing agent.
2
Check which substance loses oxygen: Fe2O3 loses oxygen to form Fe, so Fe2O3 is reduced, making it the oxidising agent.
3
Confirm both oxidation and reduction occur, so this is a redox reaction.

Exam tip:

Core tier questions will never ask for electron transfer or oxidation state answers unless explicitly marked for Extended candidates.

2. Extended: Redox as Electron Transfer★★★☆☆Extended only⏱ 12 min

📘 Definition

Redox Reaction (Extended)

A reaction where electrons are transferred between reactants. Oxidation = loss of electrons, reduction = gain of electrons, even when no oxygen is present in the reaction.

You can use electron transfer to identify redox reactions for any reaction, including those between metals and non-metals to form ionic compounds. The species that loses electrons is the reducing agent, the species that gains electrons is the oxidising agent.

📐 Worked Example

Identify the oxidising and reducing agent in the reaction between sodium and chlorine: $2 N a + C l_{2} \to 2 N a C l$

1
Identify ion charges in the product: Na forms $N a^{+}$ , Cl forms $C l^{-}$ .
2
Na loses 1 electron to form $N a^{+}$ : this is oxidation, so Na is the reducing agent.
3
Each Cl atom in Cl2 gains 1 electron to form $C l^{-}$ : this is reduction, so Cl2 is the oxidising agent.

Exam tip:

Always confirm charge balance when using electron transfer definitions: total electrons lost = total electrons gained in any redox reaction.

3. Extended: Oxidation State Calculations★★★★☆Extended only⏱ 15 min

Oxidation state of an uncombined element = 0
Oxidation state of a monatomic ion = its charge
Oxygen = -2 (exception: peroxides = -1, or bonded to fluorine)
Hydrogen = +1 (exception: metal hydrides = -1)
Sum of oxidation states in a neutral compound = 0
Sum of oxidation states in a polyatomic ion = its overall charge

An increase in oxidation state of an element between reactant and product means it is oxidised. A decrease in oxidation state means it is reduced.

📐 Worked Example

Calculate the oxidation state of S in $H_{2} S O_{4}$ , and state if S is oxidised or reduced when it forms $S O_{2}$ (oxidation state of S = +4).

1
Assign known oxidation states: H = +1 each, O = -2 each, let oxidation state of S = x.
2
$(2 \times +1) + x + (4 \times -2) = 0$
3
Solve for x: $2 + x - 8 = 0$ → $x = + 6$
4
Oxidation state of S changes from +6 to +4: it decreases, so S is reduced.

Exam tip:

Always show your working for oxidation state calculations: you can get marks for correct method even if your final answer is wrong.

4. Extended: Writing Redox Half Equations★★★★☆Extended only⏱ 15 min

🚫 No Calculator

First balance the atoms of the element being oxidised or reduced on both sides
For oxidation (electron loss), add electrons to the right hand side
For reduction (electron gain), add electrons to the left hand side
Check total charge on left side equals total charge on right side

📐 Worked Example

Write the oxidation half equation for conversion of $I^{-}$ to $I_{2}$ , and the reduction half equation for conversion of $B r_{2}$ to $B r^{-}$ .

1
Oxidation half equation: balance I atoms first: $2 I^{-} \to I_{2}$
2
Charge on left = -2, charge on right = 0: add 2 electrons to right to balance charge
3
$2I^- \rightarrow I_2 + 2e^-$
4
Reduction half equation: balance Br atoms first: $B r_{2} \to 2 B r^{-}$
5
Charge on left = 0, charge on right = -2: add 2 electrons to left to balance charge
6
$Br_2 + 2e^- \rightarrow 2Br^-$

5. Common Pitfalls

Wrong move:

Claiming a reaction with only oxidation or only reduction is redox

Why:

Redox is a paired reaction: oxidation and reduction always occur together, you cannot have one without the other.

Correct move:

Always confirm both oxidation and reduction are present before classifying a reaction as redox.

Wrong move:

Mixing up electron transfer definitions for oxidation and reduction

Why:

Students often reverse loss and gain of electrons when defining redox.

Correct move:

Use the OIL RIG mnemonic to recall: Oxidation Is Loss of electrons, Reduction Is Gain of electrons.

Wrong move:

Assigning oxygen an oxidation state of -2 in all compounds

Why:

Exceptions exist for peroxides and compounds where oxygen is bonded to fluorine.

Correct move:

Only use -2 for oxygen unless the question explicitly states it is a peroxide or bonded to fluorine.

Wrong move:

Writing electrons on the wrong side of a half equation

Why:

Students confuse oxidation (electrons lost) and reduction (electrons gained) side placement.

Correct move:

Check charge balance after writing the half equation to confirm electrons are on the correct side.

Wrong move:

Mixing up oxidising and reducing agent roles

Why:

Students link the agent name to the process it undergoes, rather than the process it causes.

Correct move:

Remember: oxidising agents are reduced, reducing agents are oxidised.

6. Quick Reference Cheatsheet

Concept	Core Rule	Extended Rule
Redox Definition	Oxidation = O gain, Reduction = O loss	Oxidation = e⁻ loss / oxidation state increase; Reduction = e⁻ gain / oxidation state decrease
Oxidising Agent	Loses oxygen, causes oxidation	Gains electrons, is reduced, oxidation state decreases
Reducing Agent	Gains oxygen, causes reduction	Loses electrons, is oxidised, oxidation state increases
Half Equations	Not required	Balance atoms first, then charge with e⁻: e⁻ on right for oxidation, left for reduction

7. Frequently Asked

Do I need to know oxidation states for Core tier?

No, oxidation states, electron transfer definitions and half equations are only required for Extended tier candidates. Core tier only uses oxygen gain/loss to define redox.

How do I tell oxidising and reducing agents apart?

Remember the rule: Oxidising agents are reduced, reducing agents are oxidised. For Extended tier, you can also check oxidation state changes: oxidising agents have decreasing oxidation states, reducing agents have increasing oxidation states.

Going deeper

syllabusOfficial CIE IGCSE Chemistry 0620 SyllabusRefer to section 6.4 for redox requirements

What's Next

Redox reactions are a foundational concept across the rest of the CIE IGCSE Chemistry 0620 syllabus. You will apply this knowledge to predict reactivity series displacement reactions, understand electrolysis processes, and analyse oxidation of alcohols in organic chemistry. Extended tier candidates should practice writing half equations and calculating oxidation states regularly, as these are common high-mark questions in Paper 2 and Paper 4. Make sure you can distinguish Core and Extended content to avoid wasting time studying topics you do not need for your tier.