NEET Chemistry · NCERT Class 11

Redox Reactions notes

Oxidation & Reduction Concepts, Oxidation Number, Balancing Redox Reactions and Redox & Electrode Processes — explained topic by topic with key points, formulas, memory tricks and common mistakes.

5 topics
44 min read
19 formulas
18 PYQs

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Chapter Overview

6 min

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

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.
A species can show disproportionation when the same element is both oxidized and reduced.
Balancing redox reactions requires conservation of atoms and charge.
In acidic medium, H+ and H2O are used for balancing; in basic medium, OH- and H2O are used.
Electrode 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.

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.

Reference tables2

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

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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Formulas3

Oxidation Number Change

\text{Increase in oxidation number }=\text{ oxidation};\text{ decrease in oxidation number }=\text{ reduction}

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

O.N.: Oxidation number assigned to an atom in a species

Electron Transfer Rule

\text{Oxidation}: M → Mⁿ⁺ + ne⁻;\text{ Reduction}: X + ne⁻ → Xⁿ⁻

Represents loss and gain of electrons in redox reactions.

n: Number of electrons transferred
e⁻: Electron

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Diagrams3

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Oxidation & Reduction Concepts

8 min

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

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

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.

Reference tables3

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⁺

Function	Oxidizes another species	Reduces another species
Electron role	Accepts electrons	Donates electrons
Itself undergoes	Reduction	Oxidation
Oxidation number change	Decreases	Increases

Zn + Cu²⁺ → Zn²⁺ + Cu	Zn	Cu²⁺
2Na + Cl₂ → 2NaCl	Na	Cl₂
H₂ + Cl₂ → 2HCl	H₂	Cl₂

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Formulas3

Oxidation by Electron Loss

M → Mⁿ⁺ + ne⁻

A species loses electrons and its oxidation number increases.

M: Species undergoing oxidation
n: Number of electrons lost
e⁻: Electron

Reduction by Electron Gain

X + ne⁻ → Xⁿ⁻

A species gains electrons and its oxidation number decreases.

X: Species undergoing reduction
n: Number of electrons gained
e⁻: Electron

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Diagrams3

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Oxidation Number

10 min

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

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

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.

Reference tables3

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

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

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

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Formulas4

Neutral Compound Rule

Σ\text{ oxidation numbers }= 0

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

Σ: Sum over all atoms in the compound

Polyatomic Ion Rule

Σ\text{ oxidation numbers }=\text{ ionic charge}

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

ionic charge: Net charge on the polyatomic ion

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Diagrams3

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Balancing Redox Reactions

11 min

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

Do not balance oxygen and hydrogen before identifying the redox atoms.
Electrons must never appear in the final overall equation.
In acidic medium, H⁺ can remain in the final equation.
In basic medium, H⁺ should not remain in the final equation.
Stoichiometric coefficients from balanced equations give mole ratios for numerical problems.
Charge checking is the fastest way to catch errors in ionic redox balancing.
Both 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.

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.

Reference tables3

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

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

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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Formulas5

Oxidation Number Method Electron Balance

\text{Total increase in }O.N. =\text{ total decrease in }O.N.

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

O.N.: Oxidation number

Half-Reaction Charge Balance

Add e⁻\text{ to the more positive side to balance charge}

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

e⁻: Electron added to balance charge

Acidic Medium Balancing Rule

\text{Balance }O\text{ with }H₂O;\text{ balance }H\text{ with }H⁺

Standard rule for ion-electron balancing in acidic solution.

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

9 min

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

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

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.

Reference tables3

Process	Oxidation	Reduction
Electron role	Produces electrons	Consumes electrons
Galvanic cell sign	Negative	Positive
Electron flow	Electrons leave anode	Electrons enter cathode

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

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

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Formulas4

Anode Reaction

M → Mⁿ⁺ + ne⁻

Oxidation half-reaction occurring at the anode.

M: Metal or reducing species
n: Number of electrons lost
e⁻: Electron

Cathode Reaction

Xⁿ⁺ + ne⁻ → X

Reduction half-reaction occurring at the cathode.

Xⁿ⁺: Ion or oxidizing species
n: Number of electrons gained
e⁻: Electron

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Diagrams4

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Videos for Redox Reactions

Redox Reactions All Concepts, Tricks, PYQsVedantu NEET English · English
REDOX REACTIONS | Class 11th ONE SHOTPW NEET English · English
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NEET PYQs from Redox Reactions

The latest 3 of 18 previous-year questions from this chapter.

NEET 2026Set 50MediumQ1
Given below are two statements : Statement-I : Heating NaCl with concentrated $\mathrm{H_2SO_4}$ and $\mathrm{MnO_2}$ results in oxidation of Mn. Statement-II : Heating NaI with concentrated $\mathrm{H_2SO_4}$ and $\mathrm{MnO_2}$ results in reduction of Mn. In light of the above statements, choose the most appropriate answer from the options given below.
- ABoth Statement-I and Statement-II are correct
- BBoth Statement-I and Statement-II are incorrect
- CStatement-I is correct but Statement-II is incorrect
- DStatement-I is incorrect but Statement-II is correct
NEET 2026Set 50MediumQ2
In an acidic medium, 10 mL of 0.25 M oxalic acid is titrated with $\mathrm{KMnO_4}$ solution. If the volume of $\mathrm{KMnO_4}$ solution required to reach end point is 10 mL, the strength of the $\mathrm{KMnO_4}$ solution is
- A$0.10\ \mathrm{M}$
- B$0.20\ \mathrm{M}$
- C$0.25\ \mathrm{M}$
- D$0.15\ \mathrm{M}$
NEET 2026Set 50EasyQ3
The correct decreasing order of oxidation state of the underlined atom in each molecule is
- A$\mathrm{\underline{P}_4O_{10} > \underline{S}O_3 > H_2\underline{O}}$
- B$\mathrm{\underline{N}_2O_5 > \underline{Al}_2O_3 > H_2\underline{S}}$
- C$\mathrm{\underline{Pb}O_2 > \underline{N}_2O_3 > \underline{S}O_3}$
- D$\mathrm{\underline{P}_4O_6 > \underline{Cl}_2O_7 > \underline{Al}H_3}$

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