ChemistryNCERT Class 11

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Redox Reactions Notes

Study Notes

4 Topics19 Formulas7 PYQs33 Key Points

Chapter Overview Oxidation & Reduction Concepts Oxidation Number Balancing Redox Reactions Redox & Electrode Processes

Chapter at a Glance

Topics4

Key Points33

Formulas19

Diagrams11

NEET PYQs7

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Topics

Chapter Overview

Overview

Redox reactions are chemical reactions in which oxidation and reduction occur simultaneously. Oxidation can be understood as addition of oxygen, removal of hydrogen, loss of electrons, or increase in oxidation number. Reduction is the opposite: removal of oxygen, addition of hydrogen, gain of electrons, or decrease in oxidation number. This chapter builds from basic concepts to oxidation number rules, identification of oxidized and reduced species, balancing redox equations by oxidation number and ion-electron methods, and electrode processes. For NEET, the highest-yield areas are oxidation number calculation, disproportionation, balancing in acidic and basic media, oxidizing and reducing agents, and electron transfer in electrochemical cells.

Key Points6

1Classical oxygen-hydrogen definitions are useful but limited; electron transfer definition is broader.
2Oxidation number helps detect redox change even when electron transfer is not obvious.
3A species can show disproportionation when the same element is both oxidized and reduced.
4Balancing redox reactions requires conservation of atoms and charge.
5In acidic medium, H+ and H2O are used for balancing; in basic medium, OH- and H2O are used.
6Electrode processes convert chemical redox change into electrical effects or vice versa.

Memory Tricks3

OIL RIG

Oxidation Is Loss of electrons; Reduction Is Gain of electrons.

LEO GER

Loss of Electrons is Oxidation; Gain of Electrons is Reduction.

Agents Do Opposite

Oxidizing agent oxidizes others but itself gets reduced; reducing agent reduces others but itself gets oxidized.

Examples3

Simple Electron Transfer

Zn + Cu²⁺ → Zn²⁺ + Cu. Zinc loses electrons and is oxidized; copper ion gains electrons and is reduced.

Non-Redox Contrast

HCl + NaOH → NaCl + H₂O is neutralization, not redox, because oxidation numbers do not change.

NEET Quick Check

If any element changes oxidation number between reactants and products, the reaction is redox.

Reference Tables2

Chapter Roadmap

Main topic	What you learn	NEET focus
Oxidation & Reduction Concepts	Meaning of oxidation, reduction, oxidants and reductants	Identify agent and process
Oxidation Number	Rules, exceptions, polyatomic ions, disproportionation	Fast oxidation state calculation
Balancing Redox Reactions	Oxidation number and ion-electron methods	Acidic and basic medium balancing
Redox & Electrode Processes	Electron transfer, electrode reactions, potentials and galvanic cells	Anode-cathode and electron flow

Fast Redox Identification

Observation	Meaning	Example
O.N. increases	Species is oxidized	Fe²⁺ → Fe³⁺
O.N. decreases	Species is reduced	Cl₂ → Cl⁻
Same element increases and decreases	Disproportionation	Cl₂ in alkali
No O.N. change	Not a redox reaction	NaOH + HCl → NaCl + H₂O

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Common Mistakes3

Thinking Oxygen Is Always Required

Modern redox is based on electron transfer or oxidation number change; oxygen may be absent.

Confusing Process and Agent

The substance oxidized is the reducing agent, while the substance reduced is the oxidizing agent.

Balancing Atoms but Ignoring Charge

Ionic redox equations must balance both atoms and total charge.

Formula Cards3

Oxidation Number Change

This rule is the fastest way to identify redox changes in NEET questions.

Variables

O.N.=

Oxidation number assigned to an atom in a species

Electron Transfer Rule

Represents loss and gain of electrons in redox reactions.

Variables

n=

Number of electrons transferred

e⁻=

Electron

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Diagrams3

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Oxidation & Reduction Concepts

Overview

Oxidation and reduction were first defined using oxygen and hydrogen. Oxidation meant addition of oxygen or removal of hydrogen, while reduction meant removal of oxygen or addition of hydrogen. These classical definitions explain many reactions but fail for reactions without oxygen or hydrogen. The modern concept defines oxidation as loss of electrons and reduction as gain of electrons. Since electrons lost by one species must be gained by another, oxidation and reduction are always simultaneous. The species that causes oxidation is the oxidizing agent and itself gets reduced. The species that causes reduction is the reducing agent and itself gets oxidized. This topic is essential for identifying redox roles quickly.

Key Points6

1Oxygen-hydrogen concept is useful for combustion, metallurgy and organic examples.
2Electron transfer concept is broader and applies to ionic reactions.
3In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc is oxidized and Cu²⁺ is reduced.
4Oxidizing agents usually contain elements in high oxidation states or electronegative atoms.
5Reducing agents usually contain electropositive metals or species that can donate electrons.
6The total number of electrons lost equals the total number of electrons gained.

Memory Tricks3

OIL RIG

Oxidation Is Loss; Reduction Is Gain. This is the fastest memory trick for electron transfer questions.

Agent Opposite Rule

The oxidizing agent gets reduced; the reducing agent gets oxidized. Agent name tells what it does to others, not itself.

Oxygen-Hydrogen Opposites

Oxidation adds oxygen or removes hydrogen; reduction removes oxygen or adds hydrogen.

Examples4

Addition of Oxygen

2Mg + O₂ → 2MgO. Magnesium is oxidized because oxygen is added to it.

Removal of Hydrogen

H₂S + Cl₂ → 2HCl + S. Hydrogen is removed from H₂S, so H₂S is oxidized.

Electron Transfer

2Na + Cl₂ → 2NaCl. Sodium loses electrons to become Na⁺, while chlorine gains electrons to become Cl⁻.

Practice Question

In Fe²⁺ + Ce⁴⁺ → Fe³⁺ + Ce³⁺, Fe²⁺ is oxidized and acts as reducing agent; Ce⁴⁺ is reduced and acts as oxidizing agent.

Reference Tables3

Classical Concept of Oxidation and Reduction

Basis	Oxidation	Reduction	Example
Oxygen	Addition of oxygen	Removal of oxygen	2Mg + O₂ → 2MgO
Hydrogen	Removal of hydrogen	Addition of hydrogen	H₂S + Cl₂ → 2HCl + S
Electron	Loss of electrons	Gain of electrons	Zn + Cu²⁺ → Zn²⁺ + Cu
Oxidation number	Increase in O.N.	Decrease in O.N.	Fe²⁺ → Fe³⁺ and MnO₄⁻ → Mn²⁺

Oxidizing Agent vs Reducing Agent

Feature	Oxidizing agent	Reducing agent
Function	Oxidizes another species	Reduces another species
Electron role	Accepts electrons	Donates electrons
Itself undergoes	Reduction	Oxidation
Oxidation number change	Decreases	Increases
Common examples	KMnO₄, K₂Cr₂O₇, Cl₂, O₂	H₂, CO, Zn, Fe²⁺, I⁻

Simultaneous Oxidation and Reduction

Reaction	Oxidized species	Reduced species
Zn + Cu²⁺ → Zn²⁺ + Cu	Zn	Cu²⁺
2Na + Cl₂ → 2NaCl	Na	Cl₂
H₂ + Cl₂ → 2HCl	H₂	Cl₂

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Common Mistakes3

Calling Electron Acceptor a Reducing Agent

An electron acceptor causes oxidation of another species, so it is an oxidizing agent.

Separating Oxidation from Reduction

Free electrons do not accumulate in ordinary chemical reactions; electron loss and gain happen together.

Using Only Classical Definition

Many redox reactions have no oxygen or hydrogen, so use electron transfer or oxidation number change.

Formula Cards3

Oxidation by Electron Loss

A species loses electrons and its oxidation number increases.

Variables

M=

Species undergoing oxidation

n=

Number of electrons lost

e⁻=

Electron

Reduction by Electron Gain

A species gains electrons and its oxidation number decreases.

Variables

X=

Species undergoing reduction

n=

Number of electrons gained

e⁻=

Electron

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Diagrams3

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Oxidation Number

10m

Overview

Oxidation number is the apparent charge assigned to an atom in a molecule or ion by assuming complete transfer of bonding electrons to the more electronegative atom. It is not always the same as valency; valency shows combining capacity, while oxidation number can be positive, negative, zero, fractional or even average. Oxidation number rules help calculate oxidation states in elements, compounds and polyatomic ions. In redox reactions, the element whose oxidation number increases is oxidized and the element whose oxidation number decreases is reduced. Disproportionation occurs when the same element in one oxidation state undergoes both oxidation and reduction. NEET frequently tests rules, exceptions and quick calculations.

Key Points7

1Oxidation number is a formal assigned value, not always real ionic charge.
2Valency has no sign, but oxidation number has sign and magnitude.
3In H₂O₂, oxygen has oxidation number -1 due to peroxide linkage.
4In OF₂, oxygen has +2 because fluorine is more electronegative.
5In superoxides like KO₂, oxygen has average oxidation number -1/2.
6Fractional oxidation numbers are average values, not actual charge on every atom.
7Disproportionation needs an element in an intermediate oxidation state.

Memory Tricks3

F is Fixed

Fluorine is always -1 in compounds. Start with fluorine before assigning other elements.

Oxygen Usually Owns -2

Use oxygen as -2 first, but remember the exceptions: peroxide, superoxide and OF₂.

Sum Rule Saves Time

For neutral compounds sum is zero; for ions sum is charge. This solves most oxidation number questions.

Examples4

KMnO₄

K is +1 and O is -2. Let Mn be x: +1 + x - 8 = 0, so x = +7.

Cr₂O₇²⁻

Let Cr be x: 2x + 7(-2) = -2, so 2x = +12 and Cr = +6.

Disproportionation Example

Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O. Chlorine goes from 0 to -1 and +1, so it is disproportionation.

PYQ-Style Check

In MnO₄⁻ → Mn²⁺, Mn changes from +7 to +2, so Mn is reduced and MnO₄⁻ acts as oxidizing agent.

Reference Tables3

Rules for Assigning Oxidation Number

Species or element	Usual oxidation number	Important exception or note
Free element	0	Na, Fe, H₂, O₂, P₄ all have 0
Monoatomic ion	Equal to ionic charge	Na⁺ = +1, S²⁻ = -2
Group 1 metals	+1	Almost always +1 in compounds
Group 2 metals	+2	Almost always +2 in compounds
Fluorine	-1	Always -1 in compounds
Oxygen	-2	-1 in peroxides, -1/2 in superoxides, +2 in OF₂
Hydrogen	+1	-1 in metal hydrides like NaH
Halogens except F	-1	Can be positive with oxygen or fluorine

Oxidation Number vs Valency

Feature	Oxidation number	Valency
Meaning	Apparent charge assigned by rules	Combining capacity
Sign	Can be positive, negative or zero	Usually no sign
Fractional value	Possible as average	Generally not used
Use in redox	Tracks electron shift	Does not directly identify redox

Solved Oxidation Number Examples

Species	Calculation	Oxidation number
KMnO₄	+1 + x + 4(-2) = 0	Mn = +7
Cr₂O₇²⁻	2x + 7(-2) = -2	Cr = +6
SO₄²⁻	x + 4(-2) = -2	S = +6
H₂O₂	2(+1) + 2x = 0	O = -1
OF₂	x + 2(-1) = 0	O = +2

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Common Mistakes4

Confusing Oxidation Number with Valency

Valency is combining capacity, while oxidation number is an assigned signed value used for redox tracking.

Forgetting Peroxide Exception

In H₂O₂ and Na₂O₂, oxygen is -1, not -2.

Ignoring Ion Charge

For MnO₄⁻, the sum of oxidation numbers must equal -1, not zero.

Treating Fractional O.N. as Real Charge

Fractional oxidation number is often an average value, as in superoxides or mixed-valence compounds.

Formula Cards4

Neutral Compound Rule

The sum of oxidation numbers of all atoms in a neutral molecule is zero.

Variables

Σ=

Sum over all atoms in the compound

Polyatomic Ion Rule

The sum of oxidation numbers in a polyatomic ion equals its net charge.

Variables

ionic charge=

Net charge on the polyatomic ion

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Diagrams3

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Balancing Redox Reactions

11m

Overview

Balancing redox reactions requires both mass conservation and charge conservation. The oxidation number method balances total increase and decrease in oxidation numbers, then balances remaining atoms. The ion-electron method splits the reaction into oxidation and reduction half-reactions, balances atoms and charge separately, equalizes electrons, and adds the half-reactions. In acidic medium, oxygen is balanced using H₂O and hydrogen using H⁺. In basic medium, OH⁻ and H₂O are used, often by first balancing as acidic and then neutralizing H⁺ with OH⁻. NEET questions commonly test acidic/basic balancing, identification of oxidized and reduced species, and stoichiometric coefficients.

Key Points7

1Do not balance oxygen and hydrogen before identifying the redox atoms.
2Electrons must never appear in the final overall equation.
3In acidic medium, H⁺ can remain in the final equation.
4In basic medium, H⁺ should not remain in the final equation.
5Stoichiometric coefficients from balanced equations give mole ratios for numerical problems.
6Charge checking is the fastest way to catch errors in ionic redox balancing.
7Both oxidation number and ion-electron methods must give the same final balanced equation.

Memory Tricks3

Acidic: O-H-H

In acidic medium: balance Oxygen with H₂O, then Hydrogen with H⁺.

Basic: Add OH to H

In basic medium, if H⁺ appears, add equal OH⁻ to both sides and convert H⁺ + OH⁻ into H₂O.

Final Check: A-C-E

Check Atoms, Charge and Electrons. Electrons must be absent in the final equation.

Examples5

Acidic Medium Example

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. Here Mn changes from +7 to +2, so 5 electrons are gained.

Oxidation Half-Reaction Example

Fe²⁺ → Fe³⁺ + e⁻. Iron loses one electron and is oxidized.

Combined Reaction Example

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O. This is a standard NEET redox balancing result.

Basic Medium Practice

For reactions in alkaline medium, first balance as if acidic, then add OH⁻ to both sides equal to H⁺ present and cancel extra water.

Stoichiometric Calculation

In the balanced reaction MnO₄⁻ reacts with 5Fe²⁺, so 1 mole of permanganate oxidizes 5 moles of Fe²⁺ in acidic medium.

Reference Tables3

Oxidation Number Method Steps

Step	Action	Purpose
1	Write skeletal equation	Know reactants and products
2	Assign oxidation numbers	Detect redox atoms
3	Calculate increase and decrease	Find electron loss and gain
4	Equalize increase and decrease	Balance redox change
5	Balance remaining atoms	Conserve mass
6	Balance H and O according to medium	Complete equation
7	Check atoms and charge	Avoid final errors

Ion-Electron Method in Acidic vs Basic Medium

Feature	Acidic medium	Basic medium
Oxygen balancing	Add H₂O	Add H₂O as needed
Hydrogen balancing	Add H⁺	Use H₂O and OH⁻ or remove H⁺ using OH⁻
Charge balancing	Add electrons	Add electrons
Final allowed ions	H⁺ may remain	OH⁻ may remain, H⁺ should not remain

Common Oxidized and Reduced Species

Change	Process	Agent role
Fe²⁺ → Fe³⁺	Oxidation	Fe²⁺ is reducing agent
MnO₄⁻ → Mn²⁺	Reduction	MnO₄⁻ is oxidizing agent
Cr₂O₇²⁻ → Cr³⁺	Reduction	Cr₂O₇²⁻ is oxidizing agent
I⁻ → I₂	Oxidation	I⁻ is reducing agent

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Common Mistakes4

Leaving Electrons in Final Answer

Electrons must cancel completely when oxidation and reduction half-reactions are added.

Using H⁺ in Basic Final Equation

Basic medium final equations should not contain free H⁺; neutralize it with OH⁻.

Balancing Charge Before Atoms

In ion-electron method, balance atoms first, then charge using electrons.

Forgetting Coefficients in Electron Change

Total oxidation number increase or decrease must include the number of atoms involved.

Formula Cards5

Oxidation Number Method Electron Balance

Coefficients are chosen so that oxidation and reduction changes are equal.

Variables

O.N.=

Oxidation number

Half-Reaction Charge Balance

Used in the ion-electron method after atoms are balanced.

Variables

e⁻=

Electron added to balance charge

Acidic Medium Balancing Rule

Standard rule for ion-electron balancing in acidic solution.

Variables

H₂O=

Water used to balance oxygen atoms

H⁺=

Hydrogen ion used to balance hydrogen atoms

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Diagrams3

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Redox & Electrode Processes

Overview

Redox reactions can be separated into oxidation and reduction half-reactions at electrodes. An electrode is a conducting surface where electron transfer occurs. Oxidation takes place at the anode and reduction takes place at the cathode. In a galvanic cell, a spontaneous redox reaction produces electrical energy; electrons flow through the external circuit from anode to cathode. Oxidation potential measures tendency to lose electrons, while reduction potential measures tendency to gain electrons. A species with higher reduction potential is more easily reduced and acts as a stronger oxidizing agent. Redox processes explain batteries, corrosion, extraction of metals, electrolysis, respiration and many analytical reactions.

Key Points7

1Electrode reactions are half-reactions written separately for oxidation and reduction.
2Anode is the source of electrons in a galvanic cell.
3Cathode receives electrons in a galvanic cell.
4Conventional current direction is opposite to electron flow.
5Reduction potential values help predict feasibility of redox reactions.
6The oxidizing agent is reduced at the cathode.
7The reducing agent is oxidized at the anode.

Memory Tricks3

AN OX and RED CAT

ANode has OXidation; REDuction happens at CAThode.

Electrons Leave Anode

In a galvanic cell, electrons are produced at the anode and travel to the cathode.

High E° Reduces

A higher reduction potential means the species more strongly wants electrons and is a better oxidizing agent.

Examples5

Daniell Cell

Zn | Zn²⁺ || Cu²⁺ | Cu. Zinc is oxidized at the anode and copper ion is reduced at the cathode.

Rusting

Iron undergoes oxidation in the presence of oxygen and moisture, forming hydrated iron oxide.

Metallurgical Reduction

Fe₂O₃ + 3CO → 2Fe + 3CO₂. Iron oxide is reduced to iron, while CO is oxidized to CO₂.

Redox Titration

Acidified KMnO₄ oxidizes Fe²⁺ to Fe³⁺ and itself gets reduced from Mn(+7) to Mn(+2).

PYQ-Style Concept

If a metal electrode dissolves into solution as Mⁿ⁺, oxidation is occurring at that electrode.

Reference Tables3

Electrode Process Comparison

Feature	Anode	Cathode
Process	Oxidation	Reduction
Electron role	Produces electrons	Consumes electrons
Galvanic cell sign	Negative	Positive
Electron flow	Electrons leave anode	Electrons enter cathode
Memory	AN OX	RED CAT

Oxidation Potential vs Reduction Potential

Property	Oxidation potential	Reduction potential
Meaning	Tendency to lose electrons	Tendency to gain electrons
Favoured species	Reducing agents	Oxidizing agents
Sign relation	Negative of reduction potential	Negative of oxidation potential
Higher value means	Easier oxidation	Easier reduction

Applications of Redox Reactions

Application	Redox idea involved	Example
Batteries	Spontaneous redox produces electricity	Dry cell, lead storage battery
Corrosion	Metal oxidation by environment	Rusting of iron
Metallurgy	Reduction of metal oxides	Fe₂O₃ reduced by CO
Respiration	Oxidation of glucose	C₆H₁₂O₆ converted to CO₂ and H₂O
Redox titration	Electron-transfer stoichiometry	KMnO₄ vs Fe²⁺

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Common Mistakes4

Reversing Electron Flow

In a galvanic cell, electrons flow externally from anode to cathode, not cathode to anode.

Forgetting Potential Sign Relation

For the same half-reaction, oxidation potential is the negative of reduction potential.

Mixing Cell Types

In galvanic cells anode is negative and cathode is positive; in electrolytic cells signs are different, but oxidation is still at anode.

Calling Salt Bridge an Electron Path

The salt bridge allows ion migration to maintain neutrality; electrons flow through the external wire.

Formula Cards4

Anode Reaction

Oxidation half-reaction occurring at the anode.

Variables

M=

Metal or reducing species

n=

Number of electrons lost

e⁻=

Electron

Cathode Reaction

Reduction half-reaction occurring at the cathode.

Variables

Xⁿ⁺=

Ion or oxidizing species

n=

Number of electrons gained

e⁻=

Electron

Cell Potential

Standard cell potential for a galvanic cell using reduction potentials.

Variables

E°cell=

Standard cell potential

E°cathode=

Standard reduction potential of cathode

E°anode=

Standard reduction potential of anode

Oxidation and Reduction Potential Relation

For the same half-cell, oxidation potential is the negative of reduction potential.

Variables

E°oxidation=

Standard oxidation potential

E°reduction=

Standard reduction potential

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Formula Sheet

Oxidation Number Change

This rule is the fastest way to identify redox changes in NEET questions.

Variables

O.N.=

Oxidation number assigned to an atom in a species

Electron Transfer Rule

Represents loss and gain of electrons in redox reactions.

Variables

n=

Number of electrons transferred

e⁻=

Electron

Charge Conservation

A balanced ionic redox equation must conserve total charge as well as atoms.

Variables

charge=

Algebraic sum of ionic charges

Oxidation by Electron Loss

A species loses electrons and its oxidation number increases.

Variables

M=

Species undergoing oxidation

n=

Number of electrons lost

e⁻=

Electron

Reduction by Electron Gain

A species gains electrons and its oxidation number decreases.

Variables

X=

Species undergoing reduction

n=

Number of electrons gained

e⁻=

Electron

Electron Conservation

Every redox reaction follows conservation of charge through equal electron loss and gain.

Variables

electrons lost=

Electrons released during oxidation

electrons gained=

Electrons accepted during reduction

Neutral Compound Rule

The sum of oxidation numbers of all atoms in a neutral molecule is zero.

Variables

Σ=

Sum over all atoms in the compound

Polyatomic Ion Rule

The sum of oxidation numbers in a polyatomic ion equals its net charge.

Variables

ionic charge=

Net charge on the polyatomic ion

Oxidation Number Change

Positive change indicates oxidation; negative change indicates reduction.

Variables

O.N.=

Oxidation number

Disproportionation Condition

A reaction is disproportionation when one species acts as both oxidant and reductant.

Variables

O.N.=

Oxidation number of the same element

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Oxidation and reduction always occur together.

Oxidation means loss of electrons and increase in oxidation number.

Reduction means gain of electrons and decrease in oxidation number.

Oxidizing agent gets reduced; reducing agent gets oxidized.

Oxidation number is a bookkeeping charge assigned using fixed rules.

Redox equations can be balanced by oxidation number method or ion-electron method.

At anode oxidation occurs; at cathode reduction occurs.

Electron flow in a galvanic cell is from anode to cathode through the external circuit.

Classical oxygen-hydrogen definitions are useful but limited; electron transfer definition is broader.

Oxidation number helps detect redox change even when electron transfer is not obvious.

Classical oxidation: addition of oxygen or removal of hydrogen.

Classical reduction: removal of oxygen or addition of hydrogen.

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NEET PYQs — Redox Reactions

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NEET 2024Set T1EasyQ1

Which reaction is NOT a redox reaction?

NEET 2023Set E1MediumQ2

On balancing the given redox reaction, a Cr₂O₇²⁻ + b SO₃²⁻(aq) + c H⁺(aq) → 2a Cr³⁺(aq) + b SO₄²⁻(aq) + c/2 H₂O(l) the coefficients a, b and c are found to be, respectively -