IMAT Chemistry · Free notes

Redox Reactions: Oxidation Numbers, Agents, Balancing and Cells

Redox Reactions has 18 past questions since 2011, and the ministry papers from 2023 on have asked 7 of them. Almost every one comes down to a single skill: assigning oxidation numbers quickly and without error. Once you can do that, deciding whether a reaction is redox, naming the reducing agent and balancing an equation are all bookkeeping with electrons. The 2026 paper added balancing questions whose options are lists of coefficients, so practise counting electrons lost and gained until it is automatic.

The handout has every subtopic, worked example, formula and trap with answers shown. The formula sheet is the chapter’s formulas and traps alone, on one page.

IMAT home All formulas on one page

Subtopic notes

Formula & revision sheet

8 formulas · 5 reference tables · 16 gotchas across all subtopics — the exam-eve cheat-sheet

Oxidation Numbers: Rules and Calculation

Formulas (2)

Working out an oxidation number from a formula · Sum of oxidation numbers $\sum (\text{number of atoms} \times \text{oxidation number}) = \text{overall charge}$
The range of oxidation numbers, and fractional averages · Limits for a main-group element $\text{lowest} = \text{outer electrons} - 8 \qquad \text{highest} = \text{outer electrons}$

Reference tables (1)

The rules for assigning oxidation numbers7 rows

Atom or situation	Oxidation number	Exception or example
An uncombined element	0	$\mathrm{Na}$ , $\mathrm{O_2}$ , $\mathrm{Cl_2}$ , $\mathrm{P_4}$ , $\mathrm{S_8}$ are all 0
A simple (monatomic) ion	Equal to its charge	$\mathrm{Mg^{2+}}$ is +2, $\mathrm{Cl^-}$ is −1
Fluorine in a compound	−1 always	The most electronegative element, no exceptions
Group 1, Group 2, aluminium	+1, +2, +3	Almost no exceptions in IMAT
Hydrogen	+1	−1 in metal hydrides: $\mathrm{NaH}$ , $\mathrm{LiH}$ , $\mathrm{CaH_2}$
Oxygen	−2	−1 in peroxides ( $\mathrm{H_2O_2}$ , $\mathrm{Na_2O_2}$ ); +2 in $\mathrm{OF_2}$ Peroxides are the exception IMAT likes most: in hydrogen peroxide each O is −1.
Cl, Br, I	−1 in halides	Positive when bonded to O: +1 in $\mathrm{HClO}$ , +7 in $\mathrm{HClO_4}$

Watch out for (3)

Oxygen is not always −2, and hydrogen is not always +1
→ The rules for assigning oxidation numbers
Divide by the number of atoms of the unknown element
→ Working out an oxidation number from a formula
Oxidation numbers can be fractions, but not irrational
→ The range of oxidation numbers, and fractional averages

Recognising Redox: Oxidation, Reduction and the Agents

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

Balancing Redox Equations

Formulas (3)

Balancing by matching electrons lost and gained · The balancing condition $(\text{atoms oxidised}) \times (\text{rise per atom}) = (\text{atoms reduced}) \times (\text{fall per atom})$
Half-equations in acidic solution · Two common reduction half-equations in acid $\mathrm{MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O} \qquad \mathrm{Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O}$
Disproportionation and comproportionation · Splitting one species between two products $(\text{atoms going up}) \times (\text{rise}) = (\text{atoms going down}) \times (\text{fall})$

Watch out for (3)

Atom balance alone can hide a wrong electron count
→ Balancing by matching electrons lost and gained
Electrons go on the side that is more positive
→ Half-equations in acidic solution
In disproportionation one species is both agents
→ Disproportionation and comproportionation

The Reactivity Series and Displacement Reactions

Formulas (1)

Predicting displacement reactions of metals and halogens · Metal displacement, as an ionic equation $\mathrm{M_{(more\ reactive)} + N^{n+}_{(aq)} \rightarrow M^{n+}_{(aq)} + N_{(less\ reactive)}}$

Reference tables (1)

The reactivity series of metals7 rows

Metal	With cold water	With dilute HCl	How it is extracted
Potassium, sodium	Violent: hydroxide and $\mathrm{H_2}$	Dangerously violent	Electrolysis of the molten chloride
Calcium	Steady: hydroxide and $\mathrm{H_2}$	Vigorous	Electrolysis of the molten chloride
Magnesium	Very slow (fast with steam)	Vigorous, $\mathrm{H_2}$	Electrolysis
Aluminium	None (oxide layer)	Slow at first, then steady	Electrolysis of molten $\mathrm{Al_2O_3}$
Zinc, iron	None (react with steam when hot)	Moderate, $\mathrm{H_2}$	Heating the oxide with carbon or CO
Lead	None	Very slow	Heating the oxide with carbon
Copper, silver, gold	None	None	Reduction of ores; silver and gold also found as the metal

Higher in the table means a stronger reducing agent; the ions of metals lower down are stronger oxidising agents.

Watch out for (1)

For halogens, the more reactive one is the oxidising agent
→ Predicting displacement reactions of metals and halogens

Electrochemical Cells and Electrolysis

Formulas (1)

Standard reduction potentials and the cell voltage · Standard cell potential $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$

Reference tables (2)

Galvanic and electrolytic cells compared5 rows

Feature	Galvanic (voltaic) cell	Electrolytic cell
Energy change	Chemical to electrical; the reaction is spontaneous	Electrical to chemical; the reaction is forced
Anode	Oxidation; negative electrode	Oxidation; positive electrode
Cathode	Reduction; positive electrode	Reduction; negative electrode
Electrons in the wire	From anode to cathode	Pushed by the supply into the cathode and drawn out of the anode
Examples	Daniell cell (zinc and copper), batteries, fuel cells	Electrolysis of brine, aluminium extraction, electroplating

Electrolysis: what forms at each electrode6 rows

Electrolyte (inert electrodes unless stated)	At the cathode (negative)	At the anode (positive)
Molten sodium chloride	Sodium metal	Chlorine
Concentrated aqueous sodium chloride (brine)	Hydrogen; sodium hydroxide is left in solution	Chlorine
Dilute sulfuric acid	Hydrogen	Oxygen, half the volume of the hydrogen
Aqueous copper(II) sulfate	Copper	Oxygen; the solution becomes acidic
Aqueous copper(II) sulfate, copper electrodes	Copper is deposited	The copper anode dissolves as $\mathrm{Cu^{2+}}$ (used to purify copper)
Molten aluminium oxide in cryolite	Aluminium	Oxygen, which burns the carbon anode to $\mathrm{CO_2}$

Watch out for (3)

The anode's sign changes, its job does not
→ Galvanic and electrolytic cells compared
Never multiply E° by the equation's coefficients
→ Standard reduction potentials and the cell voltage
A reactive metal is not deposited from its aqueous solution
→ Electrolysis: what forms at each electrode

PYQ weightage by concept

15 concepts · 0 PYQs — where the marks actually sit, so you know what to drill first

Oxidation Numbers: Rules and Calculation0 PYQs · 0%

Concept	PYQs	Share
The rules for assigning oxidation numbersfoundation	—	—
Working out an oxidation number from a formulafoundation	—	—
The range of oxidation numbers, and fractional averagesfoundation	—	—

Recognising Redox: Oxidation, Reduction and the Agents0 PYQs · 0%

Concept	PYQs	Share
Oxidation and reduction as electron loss and gainfoundation	—	—
Telling redox reactions from non-redox reactionsfoundation	—	—
Identifying the oxidising agent and the reducing agentfoundation	—	—
Common oxidising and reducing agentsfoundation	—	—

Balancing Redox Equations0 PYQs · 0%

Concept	PYQs	Share
Balancing by matching electrons lost and gainedfoundation	—	—
Half-equations in acidic solutionfoundation	—	—
Disproportionation and comproportionationfoundation	—	—

The Reactivity Series and Displacement Reactions0 PYQs · 0%

Concept	PYQs	Share
The reactivity series of metalsfoundation	—	—
Predicting displacement reactions of metals and halogensfoundation	—	—

Electrochemical Cells and Electrolysis0 PYQs · 0%

Concept	PYQs	Share
Galvanic and electrolytic cells comparedfoundation	—	—
Standard reduction potentials and the cell voltagefoundation	—	—
Electrolysis: what forms at each electrodefoundation	—	—