Chapter 7 · Introduction

Chapter 7 — Redox Reactions

Rusting, respiration, bleaching, batteries and the extraction of metals are all redox reactions — electrons passing from one substance to another. Oxidation numbers, the four types of redox reaction, two methods of balancing, redox titrations and the electrochemical series.

Drop a few crystals of purple potassium permanganate into a test tube of acidified iron(II) sulphate solution and the colour vanishes almost at once; keep adding and, suddenly, one drop stays pink. Dip an iron nail into blue copper sulphate and within minutes it is coated with copper. Join a zinc plate and a copper plate through a lemon or a jar of salt solution and a small bulb glows. In each case electrons have been passed from one substance to another: one has been oxidised and the other reduced. The same electron transfer rusts bridges, bleaches cloth, powers every battery and, in your own cells, turns glucose and oxygen into the energy you are using to read this sentence.

This chapter makes the idea of redox exact. It begins with the electron-transfer picture of oxidation and reduction and with oxidation numbers — a bookkeeping device that reveals a redox change even in covalent molecules where no electron is fully transferred. It then sorts redox reactions into four types (combination, decomposition, displacement and disproportionation), teaches two reliable ways of balancing even the most awkward redox equations, and uses them in redox titrations with permanganate, dichromate and iodine. Finally it separates the two halves of a redox reaction into a galvanic cell and arranges electrode potentials into the electrochemical series, which predicts which reactions can happen. Oxidation numbers, balancing and cell EMF appear in every Board, JEE and NEET paper.

Oxidation as electron loss→Competitive electron transfer→Oxidation numbers→Types of redox reaction→Balancing: ON and half-reaction methods→Redox titrations→Daniell cell & electrode potential→Electrochemical series

🎯 By the end of this chapter you can…

define oxidation and reduction in terms of oxygen and hydrogen, of electron transfer and of oxidation number, and identify oxidising and reducing agents;
explain competitive electron transfer reactions and use the activity series to predict metal displacement;
assign oxidation numbers by the rules — including peroxides, superoxides, $\mathrm{OF_{2}}$, metal hydrides and fractional averages such as $\mathrm{Fe_{3}O_{4}}$ — and use Stock notation;
classify redox reactions as combination, decomposition, displacement or disproportionation;
balance redox equations by the oxidation-number method and by the half-reaction method in acidic and basic solution;
explain permanganate, dichromate and iodometric titrations, their indicators, and calculate results from titration data;
describe the Daniell cell, the salt bridge and the standard hydrogen electrode, and calculate $E^{\circ}_{\text{cell}}$;
use the electrochemical series to predict the strengths of oxidising and reducing agents and the feasibility of redox reactions.

Pick a reaction, or press ▶ Play to tour them all. Under every formula the lab shows the oxidation number of each element (Section 7.1). An element whose oxidation number rises is oxidised (red), one whose oxidation number falls is reduced (blue), and an element that does both is purple. The verdict names the type of redox reaction (Section 7.2) — or says it is not a redox reaction at all.

Reaction:

type—

oxidised—

reduced—

agents—

Sections in this chapter

7.1

Oxidation, Reduction and Oxidation Number

Oxidation and reduction from oxygen to electrons; half-reactions, oxidising and reducing agents; competitive electron transfer and the activity series; the rules for oxidation numbers, Stock notation, and redox in terms of oxidation number.

oxidation: loss of $\mathrm{e^{-}}$, ON rises; $\mathrm{Cr_{2}O_{7}^{2-}}$: $2x-14=-2$, $x=+6$

Open section →7.2

Types of Redox Reactions and Balancing Redox Equations

Combination, decomposition, displacement and disproportionation; the oxidation-number method; the half-reaction method in acidic and basic solution; redox titrations with permanganate, dichromate and iodine; limits of the oxidation-number idea.

$\mathrm{MnO_{4}^{-}+8H^{+}+5e^{-}\rightarrow Mn^{2+}+4H_{2}O}$

Open section →7.3

Redox Reactions and Electrode Processes

The Daniell cell and the salt bridge; electrode potential and the standard hydrogen electrode; cell EMF; the electrochemical series and what it predicts about oxidising and reducing agents, displacement and corrosion.

$E^{\circ}_{\text{cell}}=E^{\circ}_{\text{cathode}}-E^{\circ}_{\text{anode}}$; Daniell: $1.10\ \mathrm{V}$

Open section →

✅ Before you start

Redox chemistry is about electrons, so you need the electronic configurations of atoms and ions (Section 2.4), and electronegativity and the periodicity of valence (Section 3.3) — oxidation numbers follow the group number. The difference between ionic and covalent bonds and the idea of a shared pair pulled towards the more electronegative atom come from Section 4.1. Balancing equations, moles and molarity — used in every titration calculation — are in Section 1.3, and Gibbs energy (for the optional link between cell EMF and spontaneity) is in Section 5.4. From Class 10 Science (Chemical Reactions and Equations; Metals and Non-metals) it helps to remember the reactivity series of metals and the reactions of metals with acids and with salt solutions.

← 6.4 Buffers, Salt Hydrolysis and Solubility Product 7.1 Oxidation, Reduction and Oxidation Number →