Redox Reactions

Introduction to Redox Reactions

Chemical reactions involving transfer of electrons from one substance to another are termed oxidation-reduction or redox reactions.

Molecular and Ionic Equations

Molecular Equations

When reactants and products are written in molecular forms, it is a molecular equation.

MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂

Ionic Equations

When reactants and products are written as ions, it is an ionic equation.

MnO₂ + 4H⁺ + 4Cl⁻ → Mn²⁺ + 2Cl⁻ + 2H₂O + Cl₂

Spectator Ions

Ions that do not change and are equal in number on both sides are spectator ions, omitted from final equation.

Ionic Equation:

Zn + 2H⁺ + 2Cl⁻ → Zn²⁺ + 2Cl⁻ + H₂

Final Equation:

Zn + 2H⁺ → Zn²⁺ + H₂

Oxidation

Process involving: addition of oxygen, removal of hydrogen, addition of non-metal, removal of metal, increase in +ve valency, loss of electrons, increase in oxidation number.

Addition of Oxygen

2Mg + O₂ → 2MgO

Removal of Hydrogen

H₂S + Cl₂ → 2HCl + S

Addition of Non-metal

Fe + S → FeS

Removal of Metal

2KI + H₂O₂ → 2KOH + I₂

Increase in +ve Valency

Fe²⁺ → Fe³⁺ + e⁻

Loss of Electrons (De-electronation)

H⁰ → H⁺ + e⁻

MnO₄²⁻ → MnO₄⁻ + e⁻

Increase in Oxidation Number

Mg⁰ → Mg²⁺

[Fe(CN)₆]⁴⁻ → [Fe(CN)₆]³⁻

Reduction

Reverse of oxidation: removal of oxygen, addition of hydrogen, removal of non-metal, addition of metal, decrease in +ve valency, gain of electrons, decrease in oxidation number.

Removal of Oxygen

CuO + C → Cu + CO

Addition of Hydrogen

Cl₂ + H₂ → 2HCl

Removal of Non-metal

2HgCl₂ + SnCl₂ → Hg₂Cl₂ + SnCl₄

Addition of Metal

HgCl₂ + Hg → Hg₂Cl₂

Decrease in +ve Valency

Fe³⁺ → Fe²⁺

Gain of Electrons (Electronation)

Zn²⁺ + 2e⁻ → Zn

Pb²⁺ + 2e⁻ → Pb

Decrease in Oxidation Number

Mg²⁺ → Mg⁰

[Fe(CN)₆]³⁻ → [Fe(CN)₆]⁴⁻

Redox Reactions

Overall reaction where oxidation and reduction occur simultaneously, involving electron transfer.

Types of Redox Reactions

Direct Redox

Oxidation and reduction in same vessel

Indirect Redox

Oxidation and reduction in different vessels (electrochemical cells)

Intermolecular

One substance oxidized, another reduced

2Al + Fe₂O₃ → Al₂O₃ + 2Fe

Intramolecular

One element oxidized, another reduced in same compound

2KClO₃ → 2KCl + 3O₂

Oxidising and Reducing Agents

Oxidising Agent

Gains electrons, gets reduced; oxidation number decreases.

Reducing Agent

Loses electrons, gets oxidized; oxidation number increases.

Important Oxidising Agents

Electronegative Elements

O₂, O₃, X₂ (halogens)

Highest Oxidation State

KMnO₄, K₂Cr₂O₇, HNO₃, H₂SO₄, FeCl₃

Oxides

MgO, CuO, CrO₃, CO₂, P₄O₁₀

Strongest

Fluorine

Important Reducing Agents

Metals

Na, Zn, Fe, Al

Non-metals

C, H₂, S

Hydracids

HCl, HBr, HI, H₂S

Lower Oxidation State

FeSO₄, SnCl₂, HgCl₂, Cu₂O

Metallic Hydrides

NaH, LiH

Organic

HCOOH, (COOH)₂, aldehydes, alkanes

Strongest

Lithium (in solution), Cesium (no water)

Dual Agents

Act as both: H₂O₂, SO₂, H₂SO₃, HNO₂, NaNO₂, Na₂SO₃, O₃

Identification Tips

Highest Oxidation State

Oxidising agent (e.g., KMnO₄, HNO₃)

Lowest Oxidation State

Reducing agent (e.g., H₂S, SnCl₂)

Intermediate State

Both (e.g., H₂O₂, SO₂)

Electronegative High State

Powerful oxidising (e.g., KClO₃, KIO₃)

Electronegative Low State

Powerful reducing (e.g., I⁻, Br⁻)

Equivalent Weight of Oxidising/Reducing Agents

Eq. wt. of O.A. = Molecular weight / No. of electrons gained per molecule
= Molecular weight / Change in O.N. per mole

Eq. wt. of R.A. = Molecular weight / No. of electrons lost per molecule
= Molecular weight / Change in O.N. per mole

Agents	O.N.	Product	O.N.	Change in O.N. per atom	Total Change in O.N. per mole	Eq. wt.
Cr₂O₇²⁻	+6	Cr³⁺	+3	3	6	Mol. wt./6
C₂O₄²⁻	+3	CO₂	+4	1	2	Mol. wt./2
S₂O₃²⁻	+2	S₄O₆²⁻	+2.5	0.5	1	Mol. wt./1
H₂O₂	-1	H₂O	-2	1	2	Mol. wt./2
H₂O₂	-1	O₂	0	1	2	Mol. wt./2
MnO₄⁻ (Acidic)	+7	Mn²⁺	+2	5	5	Mol. wt./5
MnO₄⁻ (Neutral)	+7	MnO₂	+4	3	3	Mol. wt./3
MnO₄⁻ (Alkaline)	+7	MnO₄²⁻	+6	1	1	Mol. wt./1

Oxidation Number or Oxidation State

Charge on an atom produced by donating or accepting electrons.

Valency vs Oxidation Number

Oxidation Number	Valency
Charge (real/imaginary) on atom in combination; + or - sign	Combining capacity; no sign
May vary depending on compound	Usually fixed
Whole number or fractional	Always whole number
May be zero	Never zero (except noble gases)

Nomenclature

-ous (lower O.N.), -ic (higher O.N.) e.g., cuprous (Cu⁺), cupric (Cu²⁺)

Stock system: Roman numerals e.g., Copper(I) oxide (Cu₂O), Iron(II) chloride (FeCl₂)

Important Points to Remember

Key JEE Points

Redox: Oxidation + Reduction simultaneously
Oxidation: Loss of e⁻, ↑ O.N.; Reduction: Gain of e⁻, ↓ O.N.
Spectator ions omitted from net ionic equation
Oxidising agent gets reduced; Reducing agent gets oxidized
Eq. wt. = Mol. wt. / Change in O.N. per mole
O.N. can be fractional; Valency is whole number
Fluorine: Strongest oxidising; Lithium: Strongest reducing in solution
Direct redox: Same vessel; Indirect: Different vessels