IMAT Chemistry · Redox Reactions

Recognising Redox: Oxidation, Reduction and the Agents

In a redox reaction one species loses electrons (its oxidation number rises) and another gains them; the species that takes electrons is the oxidising agent, the one that gives them is the reducing agent.

All Redox Reactions notes

Why this matters

The ministry papers asked for the reducing agent in a metal extraction (2023) and when zinc reacts with nitric acid (2024). The older papers asked which reaction in a list is or is not redox, and which species act as oxidising or reducing agents.

Concept 1 of 4: Oxidation and reduction as electron loss and gain

Electrons cannot vanish: if one atom loses them, another must take them. So oxidation and reduction always happen together, which is why the reaction is called redox. Losing negative electrons makes an atom more positive, so oxidation raises the oxidation number and reduction lowers it.

Definition

Oxidation is loss of electrons, an increase in oxidation number. Reduction is gain of electrons, a decrease in oxidation number. Memory aid: OIL RIG (Oxidation Is Loss, Reduction Is Gain).
Older definitions still appear: oxidation as gain of oxygen or loss of hydrogen, reduction as loss of oxygen or gain of hydrogen.
Electrons lost = electrons gained in every complete redox equation.
A half-equation shows one side of the transfer with the electrons written in.

Two half-equations, one transfer

\underbrace{\mathrm{Zn \rightarrow Zn^{2+} + 2e^-}}_{\text{oxidation}} \qquad \underbrace{\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}}_{\text{reduction}}

$\mathrm{e^-}$ an electron; on the right in oxidation, on the left in reduction

Worked example

In the thermite reaction,

\mathrm{2Al + Fe_2O_3 \rightarrow Al_2O_3 + 2Fe}

, what is oxidised, what is reduced, and how many electrons move per equation as written?

Practice this conceptself-check · 3 quick reps

Oxidation raises the oxidation number

Losing electrons, which are negative, makes an atom more positive. So oxidation is an INCREASE in oxidation number and reduction is a DECREASE. Mixing up the direction reverses every later answer about agents.

Concept 2 of 4: Telling redox reactions from non-redox reactions

The only sure test is to look for a change in oxidation number. Many reactions just swap partners or pass protons around without moving any electrons: precipitations, neutralisations and acid-base reactions. Those are not redox, however dramatic they look.

Definition

Assign oxidation numbers to every element on both sides. Any change means redox; no change anywhere means not redox.
A reaction with an uncombined element on one side and the same element combined on the other is redox (the change from 0 is unavoidable). Converting one form of an element to another, such as $\mathrm{O_2}$ to $\mathrm{O_3}$ , is not.
Usually not redox: neutralisation, acid-base proton transfer, precipitation, double exchange of ions (a salt with an acid), formation of complexes, combining two oxides.
Decomposition can go either way: $\mathrm{CaCO_3 \rightarrow CaO + CO_2}$ is not redox; $\mathrm{2KClO_3 \rightarrow 2KCl + 3O_2}$ is.

Worked example

Decide which of these are redox: (1)

\mathrm{CaCO_3 \rightarrow CaO + CO_2}

; (2)

\mathrm{2H_2O_2 \rightarrow 2H_2O + O_2}

; (3)

\mathrm{AgNO_3 + NaCl \rightarrow AgCl + NaNO_3}

Practice this conceptself-check · 3 quick reps

No free element does not mean no redox

A reaction can be redox with compounds on both sides: in

\mathrm{2FeCl_3 + SnCl_2 \rightarrow 2FeCl_2 + SnCl_4}

, iron falls from +3 to +2 and tin rises from +2 to +4. Always check the oxidation numbers rather than looking for an element.

A decomposition is not automatically redox

Breaking a carbonate into an oxide and carbon dioxide changes no oxidation numbers. Breaking a chlorate into a chloride and oxygen does. Decide by the numbers, not by the reaction type.

Concept 3 of 4: Identifying the oxidising agent and the reducing agent

An oxidising agent makes something else oxidised by taking its electrons, so the oxidising agent itself gains electrons and is reduced. The reducing agent does the opposite. The names describe what the substance does to its partner, not what happens to itself.

Definition

The oxidising agent (oxidant) accepts electrons: it is reduced, and contains the atom whose oxidation number falls.
The reducing agent (reductant) gives electrons: it is oxidised, and contains the atom whose oxidation number rises.
Name the whole species as it appears among the reactants: $\mathrm{Zn(s)}$ , not $\mathrm{Zn^{2+}}$ ; $\mathrm{HNO_3}$ , not $\mathrm{NO}$ .
A reagent can play two roles at once: some of it is reduced while the rest only supplies ions or acts as an acid.

Worked example

Copper dissolves in hot concentrated sulfuric acid:

\mathrm{Cu + 2H_2SO_4 \rightarrow CuSO_4 + SO_2 + 2H_2O}

. Name the oxidising agent and the reducing agent.

Practice this conceptself-check · 3 quick reps

The oxidising agent is the one that is reduced

The name says what it does to the OTHER substance. An oxidising agent takes electrons, so its own oxidation number goes down. Picking the species that is oxidised as the oxidising agent is the classic slip.

Name the reactant, not the product

When zinc reacts and becomes

\mathrm{Zn^{2+}}

, the reducing agent is zinc metal, not the zinc ion it turns into. Options often list both forms of the same element.

Concept 4 of 4: Common oxidising and reducing agents

A substance whose key atom is already at its highest oxidation number can only go down, so it can only oxidise others. One at its lowest can only go up, so it can only reduce. Species in between can do either, depending on the partner.

Definition

Highest oxidation state: oxidising agent only ( $\mathrm{MnO_4^-}$ with Mn +7, $\mathrm{Cr_2O_7^{2-}}$ with Cr +6, $\mathrm{HNO_3}$ with N +5).
Lowest oxidation state: reducing agent only (metals at 0, $\mathrm{I^-}$ , $\mathrm{S^{2-}}$ , $\mathrm{H^-}$ ).
In between: can do both ( $\mathrm{H_2O_2}$ , $\mathrm{SO_2}$ , $\mathrm{Fe^{2+}}$ , $\mathrm{NO_2^-}$ ).
Electronegative non-metals ( $\mathrm{F_2}$ , $\mathrm{O_2}$ , $\mathrm{Cl_2}$ ) are oxidising agents; reactive metals, carbon, carbon monoxide and hydrogen are reducing agents.

Agent	Acts as	Becomes (in acid)	Electrons per particle
Permanganate $\mathrm{MnO_4^-}$ (purple)	Oxidising agent	$\mathrm{Mn^{2+}}$ (almost colourless)	5 gained
Dichromate $\mathrm{Cr_2O_7^{2-}}$ (orange)	Oxidising agent	$\mathrm{Cr^{3+}}$ (green)	6 gained (3 per Cr)
Chlorine $\mathrm{Cl_2}$	Oxidising agent	$\mathrm{Cl^-}$	2 gained
Concentrated $\mathrm{HNO_3}$	Oxidising agent	$\mathrm{NO_2}$ (dilute acid gives NO)	1 gained (3 for NO)
Hydrogen peroxide $\mathrm{H_2O_2}$	Either	$\mathrm{H_2O}$ as oxidant; $\mathrm{O_2}$ as reductant	2
Reactive metals (Mg, Zn, Fe)	Reducing agent	Metal ions	Equal to the ion's charge, lost
Carbon, CO, $\mathrm{H_2}$	Reducing agent	$\mathrm{CO_2}$ , $\mathrm{H_2O}$	Used to extract metals from oxides

Practice this conceptself-check · 3 quick reps

Hydrogen peroxide can be either agent

With oxygen at −1, between 0 and −2,

\mathrm{H_2O_2}

oxidises iodide ions (becoming water) but reduces permanganate (becoming oxygen gas). Options that call it only an oxidising agent are incomplete.

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

Formulas (1)

Oxidation and reduction as electron loss and gain
Two half-equations, one transfer
$\underbrace{\mathrm{Zn \rightarrow Zn^{2+} + 2e^-}}_{\text{oxidation}} \qquad \underbrace{\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}}_{\text{reduction}}$

Reference tables (1)

Common oxidising and reducing agents7 rows

Agent	Acts as	Becomes (in acid)	Electrons per particle
Permanganate $\mathrm{MnO_4^-}$ (purple)	Oxidising agent	$\mathrm{Mn^{2+}}$ (almost colourless)	5 gained
Dichromate $\mathrm{Cr_2O_7^{2-}}$ (orange)	Oxidising agent	$\mathrm{Cr^{3+}}$ (green)	6 gained (3 per Cr)
Chlorine $\mathrm{Cl_2}$	Oxidising agent	$\mathrm{Cl^-}$	2 gained
Concentrated $\mathrm{HNO_3}$	Oxidising agent	$\mathrm{NO_2}$ (dilute acid gives NO)	1 gained (3 for NO)
Hydrogen peroxide $\mathrm{H_2O_2}$	Either	$\mathrm{H_2O}$ as oxidant; $\mathrm{O_2}$ as reductant	2
Reactive metals (Mg, Zn, Fe)	Reducing agent	Metal ions	Equal to the ion's charge, lost
Carbon, CO, $\mathrm{H_2}$	Reducing agent	$\mathrm{CO_2}$ , $\mathrm{H_2O}$	Used to extract metals from oxides

Watch out for (6)

Oxidation raises the oxidation number
→ Oxidation and reduction as electron loss and gain
No free element does not mean no redox
→ Telling redox reactions from non-redox reactions
A decomposition is not automatically redox
→ Telling redox reactions from non-redox reactions
The oxidising agent is the one that is reduced
→ Identifying the oxidising agent and the reducing agent
Name the reactant, not the product
→ Identifying the oxidising agent and the reducing agent
Hydrogen peroxide can be either agent
→ Common oxidising and reducing agents

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