# Redox

> CIE IGCSE Chemistry · CIE IGCSE 0620 (2026-2028)
> Source: https://www.owlsprep.com/study/cie-0620-u6-redox/

This guide covers all Core and Extended redox content for CIE IGCSE Chemistry 0620, including oxygen transfer definitions, electron transfer rules, oxidation state calculations and half equation writing.

**Prerequisites:** [Ability to write balanced chemical equations](https://www.owlsprep.com/study/cie-0620-u3-chemical-equations/); [Knowledge of ion formation and ion charges](https://www.owlsprep.com/study/cie-0620-u2-ionic-compounds/)

## Learning objectives

- Define redox reactions using oxygen gain/loss (Core)
- Identify oxidising and reducing agents using oxygen transfer (Core)
- Define redox using electron transfer (Extended)
- Calculate oxidation states to identify redox processes (Extended)
- Write balanced half equations for redox reactions (Extended)

## Core: Redox as Oxygen Transfer

**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: $CuO + H_2 \rightarrow Cu + H_2O$ (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: $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$

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.

## Extended: Redox as Electron Transfer

> **mnemonic**
>
> OIL RIG: **O**xidation **I**s **L**oss of electrons, **R**eduction **I**s **G**ain of electrons.

**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: $2Na + Cl_2 \rightarrow 2NaCl$

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

## Extended: Oxidation State Calculations

1. Oxidation state of an uncombined element = 0
2. Oxidation state of a monatomic ion = its charge
3. Oxygen = -2 (exception: peroxides = -1, or bonded to fluorine)
4. Hydrogen = +1 (exception: metal hydrides = -1)
5. Sum of oxidation states in a neutral compound = 0
6. 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 $H_2SO_4$, and state if S is oxidised or reduced when it forms $SO_2$ (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.

## Extended: Writing Redox Half Equations

1. First balance the atoms of the element being oxidised or reduced on both sides
2. For oxidation (electron loss), add electrons to the right hand side
3. For reduction (electron gain), add electrons to the left hand side
4. 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 $Br_2$ to $Br^-$.

1. Oxidation half equation: balance I atoms first: $2I^- \rightarrow I_2$
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: $Br_2 \rightarrow 2Br^-$
5. Charge on left = 0, charge on right = -2: add 2 electrons to left to balance charge
6. $$Br_2 + 2e^- \rightarrow 2Br^-$$

*Calculator:* forbidden

## Common pitfalls

- **Wrong:** Claiming a reaction with only oxidation or only reduction is redox
  - Why it fails: Redox is a paired reaction: oxidation and reduction always occur together, you cannot have one without the other.
  - Correct: Always confirm both oxidation and reduction are present before classifying a reaction as redox.
- **Wrong:** Mixing up electron transfer definitions for oxidation and reduction
  - Why it fails: Students often reverse loss and gain of electrons when defining redox.
  - Correct: Use the OIL RIG mnemonic to recall: Oxidation Is Loss of electrons, Reduction Is Gain of electrons.
- **Wrong:** Assigning oxygen an oxidation state of -2 in all compounds
  - Why it fails: Exceptions exist for peroxides and compounds where oxygen is bonded to fluorine.
  - Correct: Only use -2 for oxygen unless the question explicitly states it is a peroxide or bonded to fluorine.
- **Wrong:** Writing electrons on the wrong side of a half equation
  - Why it fails: Students confuse oxidation (electrons lost) and reduction (electrons gained) side placement.
  - Correct: Check charge balance after writing the half equation to confirm electrons are on the correct side.
- **Wrong:** Mixing up oxidising and reducing agent roles
  - Why it fails: Students link the agent name to the process it undergoes, rather than the process it causes.
  - Correct: Remember: oxidising agents are reduced, reducing agents are oxidised.

## 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 |

## 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.

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