Class 11 Chemistry Notes

Redox Reactions Class 11 Notes

Complete, exam-ready notes on redox reactions: oxidation and reduction, oxidation number and its rules, balancing redox equations by the oxidation-number and ion-electron methods, types of redox reactions, and electrochemical cells with electrode potentials and the Nernst equation — written for CBSE, JEE and NEET revision.

Class11SubjectChemistryCoversCBSE · JEE · NEET

Written byDeep Narayan· Science & Mathematics EducatorReviewed byPushpanjali

What is a redox reaction in one line?

A redox reaction is one in which electrons are transferred: one species loses electrons (oxidation) while another gains them (reduction) simultaneously.

Oxidation and Reduction

Redox reaction

Redox couples oxidation and reduction in one reaction. Oxidation is the loss of electrons (oxidation number increases: OIL — Oxidation Is Loss); reduction is the gain of electrons (oxidation number decreases: RIG — Reduction Is Gain). The oxidising agent is reduced, and the reducing agent is oxidised.

Classical view: oxidation meant adding oxygen or removing hydrogen; reduction meant removing oxygen or adding hydrogen.
Modern view (by electron transfer): oxidation loses electrons, reduction gains electrons.
The species that loses electrons is the reducing agent; the species that gains electrons is the oxidising agent.

Oxidation Number and Its Rules

Oxidation number

The charge an atom would have if all bonds were ionic — the hypothetical charge after assigning shared electrons to the more electronegative atom. It tracks electron transfer even in covalent compounds.

Free elements (O₂, N₂, Fe) have oxidation number 0.
Monatomic ions carry their charge: Na⁺ is +1, Cl⁻ is −1.
Hydrogen is +1 except in metal hydrides (NaH, CaH₂) where it is −1.
Oxygen is −2, except in peroxides (H₂O₂, −1), superoxides, OF₂ (+2) and O₂F₂ (+1).
Fluorine is always −1; alkali metals +1; alkaline earth metals +2.
The sum of oxidation numbers in a neutral molecule is 0; in a polyatomic ion it equals the ion's charge.

\text{sum of oxidation numbers} = 0\;(\text{molecule}),\quad = \text{charge}\;(\text{ion})

Charge-balance rule

Balancing Redox Equations — Oxidation-Number Method

Assign oxidation numbers and identify the atoms that change state.
Balance the total change in oxidation number: the increase on the oxidised atoms must equal the decrease on the reduced atoms.
Balance atoms other than O and H, then add H₂O, H⁺ (acidic) or OH⁻ (basic) to balance O and H.
Check that atoms and charge balance on both sides.

Acidic vs basic medium

In acidic solution add H⁺ and H₂O to balance; in basic solution add OH⁻ and H₂O. The same redox equation can look different in the two media — a common exam trap.

Ion-Electron (Half-Reaction) Method

Half-reaction method

Split the redox reaction into separate oxidation and reduction half-reactions, balance each for atoms (using H⁺/OH⁻ and H₂O) and charge (using electrons), then multiply and add so that electrons cancel.

Oxidation half: MnO₄⁻ → Mn²⁺ gains 5 e⁻ per manganese (Mn +7 → +2).
Reduction half: Fe²⁺ → Fe³⁺ loses 1 e⁻ per iron (+2 → +3).
Multiply the iron half by 5 so electrons cancel: 5Fe²⁺ + MnO₄⁻ + ... → 5Fe³⁺ + Mn²⁺ + ...
Balancing in acid: add H₂O to the side short of oxygen, H⁺ to the side short of hydrogen.

Types of Redox Reactions

Combination: A + B → product, e.g. Na + ½Cl₂ → NaCl.
Decomposition: a single compound splits, e.g. 2H₂O₂ → 2H₂O + O₂.
Displacement: a more reactive element replaces a less reactive one, e.g. Zn + CuSO₄ → ZnSO₄ + Cu.
Disproportionation: the same species is both oxidised and reduced — e.g. 2H₂O₂ → 2H₂O + O₂, and Cl₂ in alkali gives Cl⁻ and ClO⁻.
Comproportionation: two compounds of the same element at different oxidation states meet at an intermediate one.

Electrochemical Cells and Electrode Potentials

Electrochemical cell

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

A galvanic cell converts spontaneous redox energy to electricity. Oxidation happens at the anode (negative), reduction at the cathode (positive); a salt bridge balances charge. The standard emf E°cell equals the cathode potential minus the anode potential.

Standard hydrogen electrode (SHE) is the reference: E° = 0 V for 2H⁺ + 2e⁻ → H₂ under standard conditions.
Positive E° means the species is more easily reduced than H⁺; negative means it is oxidised more easily.
A reaction is spontaneous when E°cell > 0.

Nernst Equation and Gibbs Energy

E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.059}{n}\log Q\;(25^\circ\text{C})

Nernst equation

\Delta G^\circ = -nFE^\circ_{\text{cell}},\qquad \Delta G^\circ = -RT\ln K

Gibbs energy and emf

Solved Examples

Example: Find the oxidation number of S in H₂SO₄.

Solution: H is +1 (×2), O is −2 (×4), and the molecule is neutral: 2(+1) + S + 4(−2) = 0, so S = +6.

Example: In MnO₄⁻ → Mn²⁺, is manganese oxidised or reduced, and by how many electrons?

Solution: Mn goes from +7 to +2 — oxidation number decreases by 5, so it is reduced and each MnO₄⁻ gains 5 electrons.

Revision

Key formulas at a glance

Memorise these before attempting numericals — most exam questions hinge on one of them.

Standard cell potential

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

Nernst equation (25 °C)

E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.059}{n}\log Q

Gibbs energy vs emf

\Delta G^\circ = -nFE^\circ_{\text{cell}}

Gibbs energy vs K

\Delta G^\circ = -RT\ln K

Charge balance

\sum \text{ox. no.} = 0\;(\text{mol}),\; = \text{charge}\;(\text{ion})

Oxidation-state difference

\text{change} = |\text{ox. no.}(\text{final}) - \text{ox. no.}(\text{initial})|

Spontaneity

\text{spontaneous when } E^\circ_{\text{cell}} > 0

Exam tips

How this chapter is asked

Where this topic appears in CBSE, JEE Main and NEET papers.

OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain.
Hydrogen is −1 in metal hydrides; oxygen is −1 in peroxides.
Oxidation number of an element in its free state is 0.
In acidic medium balance with H⁺ and H₂O; in basic medium with OH⁻ and H₂O.
MnO₄⁻ → Mn²⁺ involves 5 electrons — a favourite titration and balancing question.
E°cell > 0 means spontaneous; in a galvanic cell the anode is the negative terminal.
SHE reference: E° = 0 V for 2H⁺ + 2e⁻ → H₂.

FAQ

Common questions

What is the difference between oxidation and reduction?

Oxidation is the loss of electrons, which raises an atom's oxidation number; reduction is the gain of electrons, which lowers it. The two always occur together, so the reaction is called a redox reaction.

How do you calculate oxidation number?

Apply the fixed rules (free elements 0, H +1, O −2, alkali metals +1, etc.) and set the sum equal to the molecule's charge — 0 for neutral molecules or the ion's charge. Solve for the unknown atom's value.

What are the two methods to balance a redox equation?

The oxidation-number method balances the total increase and decrease in oxidation state; the ion-electron method balances separate oxidation and reduction half-reactions and cancels electrons. Both work in acidic or basic media with the right balancing ions.

What does E°cell tell you about a redox reaction?

A positive E°cell means the reaction is spontaneous and can drive a galvanic cell; a negative value means it is non-spontaneous. E°cell also links to Gibbs energy through ΔG° = −nFE°cell.

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