Chemical bonding is one of the most important topics in JEE Main Chemistry. It explains how atoms combine to form molecules and compounds and why different substances have different structures, properties and stability.

The major concepts include Kossel-Lewis approach, ionic bonding, covalent bonding, Lewis structures, VSEPR theory, hybridisation, molecular orbital theory, hydrogen bonding and molecular properties.

1. Why Do Atoms Form Chemical Bonds?

Atoms combine with one another to attain a more stable electronic configuration.

Generally, atoms tend to achieve a configuration similar to that of a noble gas.

The stability is commonly associated with:

• Complete valence shell

• Lower potential energy

• Greater stability of the resulting molecule or compound

The electrons involved in chemical bonding are mainly valence electrons.

2. Valence Electrons

Electrons present in the outermost shell of an atom are called valence electrons.

For example:

• Na: 2, 8, 1 → 1 valence electron

• Mg: 2, 8, 2 → 2 valence electrons

• Cl: 2, 8, 7 → 7 valence electrons

• Ne: 2, 8 → complete valence shell

Valence electrons play the major role in chemical bonding.

3. Octet Rule

The octet rule states that atoms tend to achieve eight electrons in their valence shell during chemical bonding.

Examples:

• Na loses one electron → Na⁺

• Cl gains one electron → Cl⁻

Both achieve a stable noble-gas-like configuration.

Exceptions to Octet Rule

The octet rule does not apply universally.

Important exceptions include:

Incomplete Octet

Examples:

• BeCl₂

• BF₃

• AlCl₃

Odd-Electron Molecules

Examples:

• NO

• NO₂

Expanded Octet

Elements from the third period and beyond can have more than eight electrons around the central atom.

Examples:

• PCl₅

• SF₆

• XeF₄

4. Lewis Structures

Lewis structures represent valence electrons using dots and chemical bonds using lines.

Basic Steps

1. Count total valence electrons.

2. Select the central atom.

3. Connect atoms using single bonds.

4. Complete octets of terminal atoms.

5. Place remaining electrons on the central atom.

6. Create multiple bonds if necessary.

Example: CO₂

Carbon has four valence electrons.

Each oxygen has six.

Total:

4 + 6 + 6 = 16 electrons

The Lewis structure is:

O=C=O

Carbon and each oxygen achieve an octet.

5. Formal Charge

Formal charge helps determine the most appropriate Lewis structure.

Formula:

Formal Charge = V − L − B/2

Where:

• V = valence electrons of free atom

• L = lone-pair electrons

• B = bonding electrons

Important Points

A good Lewis structure generally has:

• Minimum formal charges

• Minimum charge separation

• Negative charge preferably on the more electronegative atom

6. Ionic Bond

An ionic bond is formed by the electrostatic attraction between oppositely charged ions.

It generally involves transfer of electrons from one atom to another.

Example:

NaCl

Na loses one electron:

Na → Na⁺ + e⁻

Cl gains one electron:

Cl + e⁻ → Cl⁻

Then:

Na⁺ + Cl⁻ → NaCl

7. Characteristics of Ionic Compounds

Ionic compounds generally:

• Have high melting and boiling points

• Are crystalline solids

• Conduct electricity in molten or aqueous state

• Are often soluble in polar solvents

• Have strong electrostatic forces

8. Lattice Enthalpy

Lattice enthalpy is associated with the formation or separation of an ionic crystal.

The magnitude of lattice enthalpy generally increases with:

• Higher ionic charge

• Smaller ionic radius

A simplified relationship is:

Lattice energy ∝ |z⁺z⁻| / r

where:

• z⁺ = cation charge

• z⁻ = anion charge

• r = distance between ions

Therefore, highly charged and smaller ions generally form stronger ionic lattices.

9. Factors Affecting Ionic Bond Formation

Important factors include:

Ionisation Enthalpy

Lower ionisation enthalpy generally makes electron loss easier.

Electron Gain Enthalpy

Greater tendency to accept electrons favours formation of anions.

Lattice Enthalpy

Higher lattice enthalpy generally increases the stability of an ionic solid.

10. Covalent Bond

A covalent bond is formed by sharing of electron pairs between atoms.

Examples:

• H₂

• Cl₂

• O₂

• N₂

• CH₄

Types of Covalent Bonds

Single Bond

One shared electron pair.

Example:

H—H

Double Bond

Two shared electron pairs.

Example:

O=O

Triple Bond

Three shared electron pairs.

Example:

N≡N

11. Sigma and Pi Bonds

Covalent bonds can involve sigma (σ) and pi (π) bonds.

Sigma Bond

A sigma bond is formed by head-on overlap of orbitals.

Characteristics:

• Stronger than a π bond

• Electron density lies along the internuclear axis

• Allows rotation in many single-bond systems

Pi Bond

A pi bond is formed by sideways overlap of orbitals.

Characteristics:

• Weaker than σ bond

• Electron density is above and below the internuclear axis

• Restricts rotation

Important Rule

A single bond contains:

1 σ bond

A double bond contains:

1 σ + 1 π

A triple bond contains:

1 σ + 2 π

12. Coordinate Covalent Bond

In a coordinate covalent bond, both shared electrons are contributed by the same atom.

Examples:

• NH₄⁺

• H₃O⁺

• BF₄⁻

The donor atom provides the electron pair, while the acceptor provides an empty orbital or suitable electron-deficient site.

13. Electronegativity

Electronegativity is the tendency of an atom to attract shared electrons towards itself.

On the Pauling scale, fluorine is the most electronegative element.

General trend:

• Across a period → electronegativity generally increases

• Down a group → electronegativity generally decreases

14. Bond Polarity

When two atoms have different electronegativities, the bonding electrons are shared unequally.

This creates a polar covalent bond.

Example:

H—Cl

Chlorine attracts the shared electron pair more strongly.

Therefore:

Hδ+—Clδ−

15. Dipole Moment

Dipole moment measures the polarity of a bond or molecule.

For a simple dipole:

μ = q × r

where:

• q = magnitude of charge

• r = distance between charges

The SI unit is C m.

A commonly used chemistry unit is Debye (D).

Molecular Dipole Moment

The molecular dipole moment is the vector sum of individual bond moments.

Therefore, molecular geometry is extremely important.

16. VSEPR Theory

VSEPR stands for:

Valence Shell Electron Pair Repulsion Theory

According to VSEPR theory, electron pairs around the central atom arrange themselves to minimize repulsion.

Repulsion Order

Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair

Therefore, lone pairs occupy more space than bonding pairs.

17. Common Molecular Geometries

18. Important Bond Angles

BeCl₂

180°

BF₃

120°

CH₄

109.5°

NH₃

Approximately 107°

H₂O

Approximately 104.5°

The bond angle decreases in:

CH₄ > NH₃ > H₂O

because the number of lone pairs increases.

19. Hybridisation

Hybridisation is the mixing of atomic orbitals of comparable energy to form new equivalent hybrid orbitals.

Important hybridisations:

20. sp Hybridisation

One s orbital combines with one p orbital.

Number of hybrid orbitals:

2

Geometry:

Linear

Bond angle:

180°

Examples:

• BeCl₂

• CO₂

• C₂H₂

21. sp² Hybridisation

One s orbital combines with two p orbitals.

Number of hybrid orbitals:

3

Geometry:

Trigonal planar

Bond angle:

120°

Examples:

• BF₃

• C₂H₄

22. sp³ Hybridisation

One s orbital combines with three p orbitals.

Number of hybrid orbitals:

4

Ideal geometry:

Tetrahedral

Bond angle:

109.5°

Examples:

• CH₄

• NH₃

• H₂O

The actual shape changes when lone pairs are present.

23. Hybridisation and Molecular Shape

CH₄

Hybridisation: sp³

Shape: Tetrahedral

Bond angle: 109.5°

NH₃

Hybridisation: sp³

Shape: Trigonal pyramidal

Bond angle: ~107°

H₂O

Hybridisation: sp³

Shape: Bent

Bond angle: ~104.5°

24. Valence Bond Theory

According to Valence Bond Theory, a covalent bond is formed by overlap of half-filled atomic orbitals.

Greater overlap generally results in a stronger bond.

Types of orbital overlap:

• s-s

• s-p

• p-p

• hybrid-hybrid

• hybrid-s

• hybrid-p

25. Molecular Orbital Theory

Molecular Orbital Theory explains molecular structure by considering molecular orbitals formed from atomic orbitals.

When atomic orbitals combine, they form:

• Bonding molecular orbitals

• Antibonding molecular orbitals

Bonding Molecular Orbital

Has lower energy than the original atomic orbitals.

It increases molecular stability.

Antibonding Molecular Orbital

Has higher energy.

It decreases molecular stability.

Antibonding orbitals are represented using an asterisk:

σ*

π*

26. Bond Order

Bond order is calculated using:

Bond Order = 1/2 (Nb − Na)

where:

• Nb = number of electrons in bonding molecular orbitals

• Na = number of electrons in antibonding molecular orbitals

Interpretation

Higher bond order generally means:

• Stronger bond

• Shorter bond length

• Greater stability

If bond order is zero, the molecule is generally unstable with respect to formation.

27. Magnetic Nature

A molecule is:

Paramagnetic

If it contains one or more unpaired electrons.

Diamagnetic

If all electrons are paired.

Important JEE Main Example

O₂ is paramagnetic because it contains unpaired electrons in its molecular orbitals.

This was one of the important successes of Molecular Orbital Theory.

28. Hydrogen Bonding

Hydrogen bonding occurs when hydrogen is attached to a highly electronegative atom and interacts with another electronegative atom having a lone pair.

Common atoms involved:

F, O and N

Examples:

• H₂O

• HF

• NH₃

29. Types of Hydrogen Bonding

Intermolecular Hydrogen Bonding

Occurs between different molecules.

Example:

H₂O molecules

Intramolecular Hydrogen Bonding

Occurs within the same molecule.

This can occur in suitable compounds where donor and acceptor groups are positioned appropriately.

30. Effects of Hydrogen Bonding

Hydrogen bonding can affect:

• Boiling point

• Melting point

• Solubility

• Viscosity

• Molecular association

For example, water has an unusually high boiling point for its molecular mass because of extensive hydrogen bonding.

31. Resonance

Some molecules cannot be represented accurately by a single Lewis structure.

Such molecules are represented by two or more contributing structures called resonance structures.

The actual molecule is a resonance hybrid.

Example:

CO₃²⁻

The three C—O bonds are equivalent in the actual carbonate ion.

32. Bond Length

Bond length is the average distance between the nuclei of two bonded atoms.

General trend:

Triple bond < Double>

For the same atoms:

• Triple bond is shortest

• Double bond is intermediate

• Single bond is longest

33. Bond Strength

Bond strength generally increases with bond order.

Therefore:

Triple bond > Double bond > Single bond

For example:

N≡N is a very strong bond.

34. Fajan's Rule

Fajan's rules help predict the covalent character of ionic compounds.

Covalent character generally increases when:

• Cation is small

• Cation has high positive charge

• Anion is large

• Anion has high negative charge

Example

AlCl₃ has considerably more covalent character than NaCl.

35. Important Factors Affecting Covalent Character

Covalent character increases with:

Small, highly charged cation + large, highly polarizable anion

Polarising power of cation increases with:

• Smaller size

• Greater positive charge

Polarizability of anion increases with:

• Larger size

• Greater negative charge

36. Dipole Moment and Molecular Shape

Dipole moment is strongly affected by molecular geometry.

CO₂

Individual C=O bond moments are present, but they cancel because the molecule is linear.

Therefore:

Net dipole moment = 0

H₂O

The molecule is bent, so bond moments do not cancel.

Therefore:

H₂O has a non-zero dipole moment.

37. Important JEE Main Comparisons

Bond Strength

Triple > Double > Single

Bond Length

Single > Double > Triple

Electron Pair Repulsion

LP–LP > LP–BP > BP–BP

Hybridisation

sp → linear

sp² → trigonal planar

sp³ → tetrahedral

sp³d → trigonal bipyramidal

sp³d² → octahedral

38. Important Formula & Quick Revision Sheet

Formal Charge

FC = V − L − B/2

Dipole Moment

μ = q × r

Bond Order

BO = 1/2(Nb − Na)

Maximum Number of Electrons

For a shell with principal quantum number n:

Maximum electrons = 2n²

Number of Orbitals in a Shell

n²

Number of Orbitals in a Subshell

2l + 1

Maximum Electrons in a Subshell

2(2l + 1)

39. Most Important Chemical Bonding Topics for JEE Main

Students should give special attention to:

1. Lewis structures

2. Octet rule and exceptions

3. Formal charge

4. Ionic bonding

5. Lattice enthalpy

6. Fajan's rule

7. Covalent bonding

8. Sigma and pi bonds

9. Electronegativity

10. Dipole moment

11. VSEPR theory

12. Molecular geometry

13. Hybridisation

14. Valence Bond Theory

15. Molecular Orbital Theory

16. Bond order

17. Magnetic properties

18. Resonance

19. Hydrogen bonding

20. Bond length and bond strength

40. Quick JEE Main Revision

Before solving questions, remember these key relationships:

LP–LP > LP–BP > BP–BP

Triple bond > Double bond > Single bond in bond strength.

Single bond > Double bond > Triple bond in bond length.

Higher bond order → stronger and shorter bond

Paramagnetic → unpaired electrons

Diamagnetic → all electrons paired

sp → 180°

sp² → 120°

sp³ → 109.5°

PCl₅ → sp³d

SF₆ → sp³d²

CH₄ → tetrahedral

NH₃ → trigonal pyramidal

H₂O → bent

CO₂ → linear

Conclusion

Chemical Bonding forms the foundation for understanding molecular structure, geometry, polarity, hybridisation and many concepts in inorganic and organic chemistry. For JEE Main, students should combine conceptual study with hybridisation questions, molecular geometry, formal-charge problems, bond-order calculations, MOT questions and previous-year questions.

After completing these notes, practice topic-wise MCQs and numerical questions to strengthen your preparation.