Introduction
Chemical bonding and molecular structure is one of the most important chapters in Class 11 Chemistry. It explains how atoms combine to form molecules and compounds. Understanding this chapter helps in predicting molecular shape, bond strength, polarity, and chemical reactivity.
Atoms combine to achieve stability, usually by attaining noble gas configuration.
Stability condition:
Stable configuration = ns^2 np^6
Kossel-Lewis Approach
According to this theory, atoms combine by transfer or sharing of electrons.
Octet rule:
Valence electrons = 8 (stable configuration)
Exception cases exist for incomplete octet and expanded octet.
Types of Chemical Bonds
Ionic Bond
Formed by transfer of electrons.
Example process:
Na → Na+ + e−
Cl + e− → Cl−
Electrostatic force:
F = (1 / (4π ε0)) × (q1 q2 / r^2)
Lattice energy relation:
U ∝ (q1 q2) / r
Properties:
- High melting point
- Conduct electricity in molten state
- Soluble in polar solvents
Covalent Bond
Formed by sharing of electrons.
Single bond:
H + H → H2
Bond order:
Bond order = (Nb − Na) / 2
Where:
Nb = bonding electrons
Na = antibonding electrons
Coordinate Bond
Both electrons donated by one atom.
Example:
NH3 + H+ → NH4+
Lewis Structures
Steps:
- Count valence electrons
- Arrange atoms
- Complete octet
- Assign lone pairs
Formal charge formula:
Formal charge = Valence electrons − Lone pair electrons − 1/2 × Bonding electrons
VSEPR Theory
Valence Shell Electron Pair Repulsion theory explains shape of molecules.
Repulsion order:
Lone pair − Lone pair > Lone pair − Bond pair > Bond pair − Bond pair
Shapes of Molecules
Linear:
Bond angle = 180°
Trigonal planar:
Bond angle = 120°
Tetrahedral:
Bond angle = 109.5°
Hybridization
Mixing of atomic orbitals.
sp hybridization:
1s + 1p → 2 sp
sp2 hybridization:
1s + 2p → 3 sp2
sp3 hybridization:
1s + 3p → 4 sp3
Molecular Orbital Theory
Combination of atomic orbitals forms molecular orbitals.
Constructive overlap:
ψ = ψA + ψB
Destructive overlap:
ψ = ψA − ψB
Bond order:
Bond order = (Nb − Na) / 2
Hydrogen Bonding
Weak attractive force between H and electronegative atom.
Condition:
H bonded with F, O, or N
Dipole Moment
Measure of polarity.
μ = q × r
Where:
μ = dipole moment
q = charge
r = distance
Bond Parameters
Bond length:
Distance between nuclei
Bond energy:
Energy required to break bond
Bond angle:
Angle between bonds
Fajan’s Rule
Covalent character increases when:
Polarization ∝ charge / size
Important Concepts for Exams
- Octet rule exceptions
- Hybridization vs shape
- Bond order calculation
- Dipole moment questions
Periodic Trends in Bonding
Electronegativity difference:
ΔEN = EN1 − EN2
If ΔEN large → ionic bond
If ΔEN small → covalent bond
Advanced Concepts
Resonance:
Actual structure = average of contributing structures
Resonance energy:
Stability ∝ delocalization
Conclusion
Chemical bonding explains how atoms form stable molecules. It is essential for understanding reactions, molecular geometry, and properties.
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FAQs
Q1. What is bond order?
Bond order = (Nb − Na) / 2
Q2. What is hybridization?
Mixing of orbitals to form new orbitals.
Q3. What is dipole moment?
μ = q × r