CHEMISTRY • Form 5 • Chapter 1: Redox Equilibrium

1.1Oxidation and Reduction

A redox reaction is a chemical reaction in which oxidation and reduction occur simultaneously. The term "redox" is derived from the combination of Reduction and Oxidation. While early chemistry defined redox reactions solely based on the gain or loss of oxygen, modern chemistry classifies redox reactions across four distinct perspectives: gain or loss of oxygen, gain or loss of hydrogen, transfer of electrons, and changes in oxidation number. These reactions govern many chemical processes around us, from the thrust-generating combustion of rocket fuels to displacement reactions of metals and halogens.

Learning Objectives

Describe redox reactions through activities based on gain or loss of oxygen, gain or loss of hydrogen, transfer of electrons, and change in oxidation number.
Identify the oxidised substance, reduced substance, oxidising agent, and reducing agent in various chemical reactions.
Construct and balance half-equations and overall ionic equations for redox processes.
Determine the oxidation number of an element in an uncombined state, simple ion, neutral compound, or polyatomic ion.
Apply IUPAC nomenclature using Roman numerals for compounds containing transition elements or metals with multiple oxidation numbers.
Investigate displacement reactions of metals from their salt solutions based on the electrochemical series.
Investigate displacement reactions of halogens from their halide solutions based on Group 17 reactivity.

The Four Concepts of Redox and Electron Transfer at a Distance

A single comparison framework summarising oxidation vs. reduction across the four definitions, supported by a schematic of the U-tube apparatus demonstrating electron movement at a distance.

Oxidation Side

Lose e⁻

Lose H

Gain O

Increase Oxidation Number

Reduction Side

Gain e⁻

Gain H

Lose O

Decrease Oxidation Number

Electron Transfer at a Distance

FeSO₄ solution

Carbon electrode X

Negative terminal

Fe²⁺ → Fe³⁺ + e⁻

G

e⁻ flow →

Acidified KMnO₄ solution

Carbon electrode Y

Positive terminal

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

H₂SO₄ electrolyte separates both arms

Brief Explanation

Understand the four definitions of oxidation and reduction step by step.

A redox reaction occurs when both oxidation and reduction occur simultaneously. The table below summarises the four definitions.

Basis of Definition	Oxidation Reaction	Reduction Reaction
Oxygen	Reactant gains oxygen	Reactant loses oxygen
Hydrogen	Reactant loses hydrogen	Reactant gains hydrogen
Electron Transfer	Reactant loses electron(s)	Reactant gains electron(s)
Oxidation Number	Oxidation number of an element increases	Oxidation number of an element decreases

Oxidising Agent

A substance that oxidises other substances and is itself reduced in a redox reaction. It is an electron acceptor, causes loss of hydrogen, supplies oxygen, or decreases in oxidation number.

Reducing Agent

A substance that reduces other substances and is itself oxidised in a redox reaction. It is an electron donor, supplies hydrogen, removes oxygen, or increases in oxidation number.

Redox in Terms of Oxygen

Reaction between magnesium, Mg and carbon dioxide, CO₂:

\[2\text{Mg}(s) + \text{CO}_2(g) \rightarrow 2\text{MgO}(s) + \text{C}(s)\]

Mg gains oxygen to form MgO, so Mg is the oxidised substance and the reducing agent. CO₂ loses oxygen to form C, so CO₂ is the reduced substance and the oxidising agent.

Redox in Terms of Hydrogen

Reaction between hydrogen sulphide, H₂S and chlorine, Cl₂:

\[\text{H}_2\text{S}(g) + \text{Cl}_2(g) \rightarrow \text{S}(s) + 2\text{HCl}(g)\]

H₂S loses hydrogen to form S, so H₂S is the oxidised substance and the reducing agent. Cl₂ gains hydrogen to form HCl, so Cl₂ is the reduced substance and the oxidising agent.

Redox in Terms of Electron Transfer
Combustion of magnesium: \[2\text{Mg}(s) + \text{O}_2(g) \rightarrow 2\text{MgO}(s)\]

Oxidation half-equation: \(\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^-\) — Mg loses electrons; electron donor / reducing agent.
Reduction half-equation: \(\text{O}_2 + 4e^- \rightarrow 2\text{O}^{2-}\) — O₂ gains electrons; electron acceptor / oxidising agent.

Laboratory Activity 1A: Transfer of Electrons at a Distance

Apparatus: U-tube clamped vertically, carbon electrodes dipped into solutions, connecting wires, galvanometer.

Chemicals:

Arm X: Freshly prepared 0.2 mol dm⁻³ iron(II) sulphate, FeSO₄ solution.
Arm Y: 0.1 mol dm⁻³ acidified potassium manganate(VII), KMnO₄ solution.
Electrolyte separating both arms: 1.0 mol dm⁻³ sulphuric acid, H₂SO₄.

Key Observations and Inferences:

Galvanometer needle: Deflects, showing electron flow through the external wire from reducing agent to oxidising agent.
Electrode X (Negative terminal): Iron(II) ion, Fe²⁺ releases electrons to become iron(III) ion, Fe³⁺. Solution turns from green to brown. Confirmed by adding NaOH(aq), which produces a brown precipitate of iron(III) hydroxide.
Oxidation half-equation: \(\text{Fe}^{2+}(aq) \rightarrow \text{Fe}^{3+}(aq) + e^-\)
Electrode Y (Positive terminal): Acidified manganate(VII) ion, MnO₄⁻ receives electrons to become manganese(II) ion, Mn²⁺. Solution turns from purple to colourless.
Reduction half-equation: \(\text{MnO}_4^-(aq) + 8\text{H}^+(aq) + 5e^- \rightarrow \text{Mn}^{2+}(aq) + 4\text{H}_2\text{O}(l)\)
Overall Ionic Equation: \[5\text{Fe}^{2+}(aq) + \text{MnO}_4^-(aq) + 8\text{H}^+(aq) \rightarrow 5\text{Fe}^{3+}(aq) + \text{Mn}^{2+}(aq) + 4\text{H}_2\text{O}(l)\]

Note on Set-up: The carbon electrode must not touch the layer of sulphuric acid to prevent premature mixing or bypassing the specific half-cell solutions. The acid acts as an electrolyte allowing movement of ions to complete the internal circuit.

Common Oxidising Agents	Half-equation
Acidified KMnO₄	\(\text{MnO}_4^- + 8\text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}\)
Acidified K₂Cr₂O₇	\(\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^+ + 6e^- \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O}\)
Chlorine water, Cl₂	\(\text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^-\)
Bromine water, Br₂	\(\text{Br}_2 + 2e^- \rightarrow 2\text{Br}^-\)
Iron(III) chloride, FeCl₃	\(\text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+}\)
Acidified hydrogen peroxide, H₂O₂	\(\text{H}_2\text{O}_2 + 2\text{H}^+ + 2e^- \rightarrow 2\text{H}_2\text{O}\)

Common Reducing Agents	Half-equation
Potassium iodide, KI	\(2\text{I}^- \rightarrow \text{I}_2 + 2e^-\)
Potassium bromide, KBr	\(2\text{Br}^- \rightarrow \text{Br}_2 + 2e^-\)
Iron(II) sulphate, FeSO₄	\(\text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^-\)
Reactive metals (e.g. Zn)	\(\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-\)
Sulphur dioxide, SO₂	\(\text{SO}_2 + 2\text{H}_2\text{O} \rightarrow \text{SO}_4^{2-} + 4\text{H}^+ + 2e^-\)

Redox in Terms of Change in Oxidation Number

Oxidation number is the charge of the element in a compound if electron transfer occurs completely.

Rules to Determine Oxidation Number:

All uncombined elements have an oxidation number of 0 (e.g. Na, C, He, O₂, Cl₂, S₈).
For monoatomic ions, oxidation number equals ionic charge (e.g. Cu²⁺ = +2, Cl⁻ = −1).
Group 1, Group 2, and Group 13 elements in compounds are always +1, +2, and +3 respectively.
Sum of oxidation numbers in a neutral compound equals 0.
Sum of oxidation numbers in a polyatomic ion equals the charge of the ion (e.g. CO₃²⁻ = −2).
Hydrogen is usually +1, except in metal hydrides (e.g. NaH, AlH₃) where it is −1.
Oxygen is usually −2, except in peroxides (e.g. H₂O₂) where it is −1.
Group 17 elements are normally −1 (fluorine is always −1; Cl, Br, I have positive values when combined with more electronegative elements like oxygen, e.g. in KClO, Cl = +1; in BrO₃⁻, Br = +5).

Naming via IUPAC Nomenclature

Roman numerals denote oxidation numbers of elements with variable oxidation states (mostly transition metals).

Fe(NO₃)₂ is Iron(II) nitrate; Fe(NO₃)₃ is Iron(III) nitrate.
MnO₂ = Manganese(IV) oxide; KMnO₄ = Potassium manganate(VII).
Metals with a single oxidation state do not need Roman numerals (e.g. Mg(NO₃)₂ is simply magnesium nitrate; Al₂O₃ is aluminium oxide).

Conversion of Fe²⁺ to Fe³⁺ (Oxidation)

Reactant: Freshly prepared FeSO₄ solution (green).

Reagent added: Bromine water, Br₂ (orange/brown oxidising agent, reduced to colourless Br⁻).

Half-equations:

\(\text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^-\) (Oxidation number of Fe increases from +2 to +3).
\(\text{Br}_2 + 2e^- \rightarrow 2\text{Br}^-\)

Confirmatory test: Add NaOH(aq) → brown precipitate of Fe(OH)₃ forms.

Conversion of Fe³⁺ to Fe²⁺ (Reduction)

Reactant: FeCl₃ solution (brown).

Reagent added: Zinc powder, Zn (reducing agent, oxidised to Zn²⁺), followed by gentle heating and filtration.

Half-equations:

\(\text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+}\) (Oxidation number of Fe decreases from +3 to +2).
\(\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-\)

Confirmatory test: Add NaOH(aq) → green precipitate of Fe(OH)₂ forms.

Displacement Reactions of Metals from Salt Solutions

A more electropositive metal (higher in the Electrochemical Series / more negative E⁰) displaces a less electropositive metal from its aqueous salt solution.

Example: Zinc plate in copper(II) nitrate, Cu(NO₃)₂ solution:

Zinc plate dissolves and becomes thinner: \(\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-\) (Oxidation).
Blue colour of copper(II) nitrate fades as Cu²⁺ ions are consumed.
Brown solid deposited: \(\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)\) (Reduction).
Overall ionic equation: \(\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)\).

Displacement Reactions of Halogens from Halide Solutions

A more reactive halogen (higher in Group 17, stronger oxidising agent) displaces a less reactive halogen from its halide solution.

Reactivity decreases down Group 17 (F₂ > Cl₂ > Br₂ > I₂ > At₂).

Identifying Halogens using 1,1,1-trichloroethane (CH₃CCl₃):

Chlorine (Cl₂): Pale yellow in water; colourless/pale yellow in CH₃CCl₃.
Bromine (Br₂): Brown / yellowish brown in water; brown / reddish brown in CH₃CCl₃.
Iodine (I₂): Brown in water; purple in CH₃CCl₃.

Example: Chlorine water added to potassium bromide, KBr:

Reaction: \(\text{Cl}_2(aq) + 2\text{KBr}(aq) \rightarrow 2\text{KCl}(aq) + \text{Br}_2(aq)\).
Ionic equation: \(\text{Cl}_2(aq) + 2\text{Br}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{Br}_2(aq)\).
Br⁻ loses electrons (oxidised; reducing agent).
Cl₂ gains electrons (reduced; oxidising agent).

Try to Answer First

Answer mentally, then check the explanation.

1

In the reaction \(2\text{NH}_3(g) + 3\text{CuO}(s) \rightarrow \text{N}_2(g) + 3\text{Cu}(s) + 3\text{H}_2\text{O}(l)\), identify the oxidised substance and the reducing agent in terms of hydrogen transfer.

Answer: The oxidised substance is NH₃ because it loses hydrogen to form N₂. Since it undergoes oxidation, NH₃ is also the reducing agent.

2

What is the oxidation number of manganese in the permanganate ion, MnO₄⁻?

Answer: +7. Workings: Let the oxidation number of Mn be \(x\). Each oxygen is −2. The overall charge of the ion is −1. Thus, \(x + 4(-2) = -1 \Rightarrow x - 8 = -1 \Rightarrow x = +7\).

3

Explain why no reaction is observed when an iodine solution is added to a test tube containing potassium bromide solution.

Answer: Iodine is located below bromine in Group 17, meaning iodine is less reactive and a weaker oxidising agent than bromine. Therefore, iodine cannot accept electrons from bromide ions, and no displacement reaction occurs.

Common Mistakes

Avoid these errors when identifying redox substances and agents.

!

Confusing the terms: students often identify the product as the oxidised or reduced substance instead of the reactant. Always identify the substance oxidised, reduced, oxidising agent, or reducing agent from the reactants.

!

Mixing up terminal designations in electron transfer at a distance: students frequently assume the positive terminal releases electrons. In reality, the negative terminal is the electrode where oxidation occurs and electrons are released into the external circuit.

!

Forgetting to write the sign before the digit in oxidation numbers: writing 2+ or 2− represents an ionic charge, whereas oxidation numbers must be written with the sign in front, such as +2 or −2.

!

Applying Roman numerals to non-variable metals: students write names like "magnesium(II) nitrate" or "aluminium(III) oxide". Metals from Group 1, 2, and 13 with single oxidation states do not take Roman numerals.

Concept Misunderstandings

Compare each misconception with the correct fact.

Misconception

Oxidation and reduction can occur independently in separate containers.

Correct Fact

Oxidation and reduction must occur simultaneously. Even in the U-tube experiment, electron loss at one electrode cannot proceed without simultaneous electron acceptance at the other electrode, linked via the electrolyte and external conductor.

Misconception

Oxidising agents are oxidised during a redox reaction.

Correct Fact

An oxidising agent causes another substance to be oxidised while it itself undergoes reduction (it receives electrons).

Misconception

A displacement reaction occurs whenever any metal is placed into any salt solution.

Correct Fact

Displacement occurs only if the solid metal is more electropositive (higher in the electrochemical series) than the metal cation in the salt solution.

Summary

Key facts to remember.

A redox reaction is a chemical reaction where oxidation and reduction take place simultaneously.
Oxidation involves: gain of oxygen, loss of hydrogen, loss of electrons, or an increase in oxidation number.
Reduction involves: loss of oxygen, gain of hydrogen, gain of electrons, or a decrease in oxidation number.
Agents: An oxidising agent gains electrons and is reduced; a reducing agent loses electrons and is oxidised.
Transfer of electrons at a distance uses an electrolyte barrier (e.g. H₂SO₄) in a U-tube where electrons flow externally from the reducing agent (negative terminal) to the oxidising agent (positive terminal).
Oxidation numbers follow specific IUPAC assignment rules; Roman numerals are used for elements showing variable oxidation states.
Displacement reactions depend directly on relative positions: higher metals in the electrochemical series displace lower metal ions; higher halogens in Group 17 displace lower halide ions.

Short Activity

Complete all parts before checking your score.

A. Objective Quiz

1

In the reaction between magnesium and carbon dioxide: \[2\text{Mg}(s) + \text{CO}_2(g) \rightarrow 2\text{MgO}(s) + \text{C}(s)\] Which statement correctly identifies the substance and its role?

CO₂ is the reducing agent because it loses oxygen. Mg is the oxidised substance because it gains oxygen. Mg is the oxidising agent because it releases electrons. CO₂ is oxidised to carbon.

2

During an experiment investigating the transfer of electrons at a distance using a U-tube with FeSO₄ solution and acidified KMnO₄ solution:

Electrons flow from the acidified KMnO₄ solution to the FeSO₄ solution. The electrode in acidified KMnO₄ solution acts as the negative terminal. Fe²⁺ ions release electrons to form Fe³⁺ ions at the negative terminal. Manganate(VII) ions are oxidised to manganese(II) ions.

B. Fill in the Blanks

3

The oxidation number of chromium in the dichromate ion, Cr₂O₇²⁻, is ________.

4

In a displacement reaction of halogens, chlorine can displace bromine from potassium bromide solution because chlorine is a stronger ________ agent than bromine.

C. Match the Redox Concept with the Description

Drag each concept on desktop. On mobile, tap a concept and then tap the matching box.

Answer Choices

Negative Terminal in U-tube Reducing Agent 1,1,1-trichloroethane layer turns purple Brown precipitate with NaOH(aq)

Correct Description

1Electrode where electrons are released by the reducing agent into the external wire

2Substance that undergoes oxidation by losing electrons or increasing in oxidation number

3Confirms the formation of iodine molecules, I₂

4Confirms the presence of iron(III) ions, Fe³⁺

Complete every answer before checking.

Keywords

Tap or hover over a term to read its meaning.

Redox reaction Oxidation Reduction Oxidising agent Reducing agent Oxidation number Electrochemical series Displacement reaction

Quick Notes

1.1Oxidation and Reduction

Learning Objectives

The Four Concepts of Redox and Electron Transfer at a Distance

Brief Explanation

Oxidising Agent

Reducing Agent

Redox in Terms of Oxygen

Redox in Terms of Hydrogen

Laboratory Activity 1A: Transfer of Electrons at a Distance

Redox in Terms of Change in Oxidation Number

Naming via IUPAC Nomenclature

Conversion of Fe2+ to Fe3+ (Oxidation)

Conversion of Fe3+ to Fe2+ (Reduction)

Displacement Reactions of Metals from Salt Solutions

Displacement Reactions of Halogens from Halide Solutions

Try to Answer First

In the reaction \(2\text{NH}_3(g) + 3\text{CuO}(s) \rightarrow \text{N}_2(g) + 3\text{Cu}(s) + 3\text{H}_2\text{O}(l)\), identify the oxidised substance and the reducing agent in terms of hydrogen transfer.

What is the oxidation number of manganese in the permanganate ion, MnO4−?

Explain why no reaction is observed when an iodine solution is added to a test tube containing potassium bromide solution.

Common Mistakes

Concept Misunderstandings

Summary

Short Activity

A. Objective Quiz

B. Fill in the Blanks

C. Match the Redox Concept with the Description

Answer Choices

Correct Description

Keywords

Conversion of Fe²⁺ to Fe³⁺ (Oxidation)

Conversion of Fe³⁺ to Fe²⁺ (Reduction)

What is the oxidation number of manganese in the permanganate ion, MnO₄⁻?