Chapter 4 Chemical Bonding and Molecular Structure

Chemical Bonding and Molecular Structure

By Aarish Sir

1. Introduction

• Chemistry is all about atoms combining to form molecules and compounds.

• Atoms combine because:

  1. They want to achieve a stable electronic configuration (like noble gases).

  2. Bonding lowers the energy of the system, making it more stable.

• Chemical Bond: The attractive force that holds atoms together in a molecule or compound.

2. Types of Chemical Bonds

(a) Ionic Bond (Electrovalent Bond)

• Formed by transfer of electrons from one atom to another.

• Metals lose electrons (form cations), non-metals gain electrons (form anions).

• Oppositely charged ions attract each other.

Example: NaCl

• Na → Na⁺ + e⁻

• Cl + e⁻ → Cl⁻

• Na⁺ + Cl⁻ → NaCl

Properties:

• Hard, brittle solids.

• High melting and boiling points.

• Conduct electricity in molten or aqueous state.

(b) Covalent Bond

• Formed by sharing of electrons between atoms.

• Both atoms contribute electrons to achieve octet.

Example: H₂ (each H shares 1 electron → H–H bond).

Properties:

• Exist as gases, liquids or soft solids.

• Low melting and boiling points (except giant covalent like diamond, SiO₂).

• Poor conductors of electricity.

(c) Coordinate (Dative) Bond

• A type of covalent bond where one atom donates both electrons for sharing.

Example: NH₄⁺ ion

• NH₃ + H⁺ → NH₄⁺

• N donates a lone pair to H⁺.

(d) Metallic Bond

• Found in metals.

• Positive metal ions surrounded by a sea of delocalized electrons.

Properties:

• Good conductors of heat and electricity.

• Malleable and ductile.

• Metallic lustre.

3. Theories of Chemical Bonding

(a) Lewis Theory (Octet Rule)

• Atoms tend to achieve 8 electrons in their outermost shell (octet).

• Explains simple molecules.

Limitations:

• Does not explain shape of molecules.

• Cannot explain molecules with odd number of electrons (NO, NO₂).

• Cannot explain stability of some compounds (BF₃).

(b) Valence Bond Theory (VBT)

• A covalent bond is formed by the overlap of atomic orbitals.

• Types of overlap:

  1. s–s overlap

  2. s–p overlap

  3. p–p overlap

• Bonds:

  • Sigma (σ) bond: Formed by head-on overlap (stronger).

  • Pi (π) bond: Formed by sidewise overlap (weaker).

(c) VSEPR Theory (Valence Shell Electron Pair Repulsion)

• Shape of molecules is determined by repulsion between electron pairs around the central atom.

• Electron pairs (bonding + lone pairs) arrange to minimize repulsion.

• Examples:

• BeCl₂ → Linear (180°)

• BF₃ → Trigonal planar (120°)

• CH₄ → Tetrahedral (109.5°)

• NH₃ → Pyramidal (107°)

• H₂O → Bent/V-shaped (104.5°)

(d) Hybridisation Theory

• Mixing of atomic orbitals to form new hybrid orbitals of equal energy.

• Explains bond angles and geometry.

Types:

• sp → Linear (180°) → BeCl₂, CO₂

• sp² → Trigonal planar (120°) → BF₃, C₂H₄

• sp³ → Tetrahedral (109.5°) → CH₄, NH₃, H₂O

• sp³d → Trigonal bipyramidal (90°, 120°) → PCl₅

• sp³d² → Octahedral (90°) → SF₆

(e) Molecular Orbital Theory (MOT)

• Atomic orbitals combine to form molecular orbitals (MOs).

• Explains bond order, magnetic nature, and stability of molecules.

Key Rules:

• Bond Order = (Number of bonding e⁻ – Number of antibonding e⁻)/2

  • If BO > 0 → molecule is stable.

  • If BO = 0 → molecule unstable.

Examples:

• H₂ → BO = 1 (stable)

• He₂ → BO = 0 (unstable)

• O₂ → BO = 2 (paramagnetic, has unpaired electrons)

4. Polarity of Bonds and Molecules

(a) Polar Covalent Bond

• Shared electrons are unequally distributed due to difference in electronegativity.

• Results in partial charges (δ⁺, δ⁻).

Example: H–Cl (Cl is more electronegative).

(b) Non-Polar Covalent Bond

• Equal sharing of electrons.

• Example: H₂, O₂, Cl₂.

(c) Dipole Moment (μ)

• Measure of polarity of a molecule.

• μ = q × d (charge × distance).

• Linear molecules like CO₂ → μ = 0 (non-polar).

• Bent molecules like H₂O → μ ≠ 0 (polar).

5. Hydrogen Bonding

• Special type of intermolecular attraction between H and highly electronegative atoms (N, O, F).

• Stronger than van der Waals forces but weaker than covalent bond.

Types:

1. Intermolecular → between molecules (e.g., H₂O molecules).

2. Intramolecular → within a molecule (e.g., o-nitrophenol).

Effects:

• High boiling points (H₂O, NH₃, HF).

• Unusual properties of water (ice is less dense than water).

6. Molecular Shapes (Summary Table – VSEPR)

Molecule	Hybridisation	Shape	Bond Angle
BeCl₂	sp	Linear	180°
BF₃	sp²	Trigonal planar	120°
CH₄	sp³	Tetrahedral	109.5°
NH₃	sp³	Trigonal pyramidal	107°
H₂O	sp³	Bent	104.5°
PCl₅	sp³d	Trigonal bipyramidal	90°,120°
SF₆	sp³d²	Octahedral	90°

7. Importance of Chemical Bonding

• Explains why and how atoms combine.

• Helps in understanding the structure and shape of molecules.

• Predicts properties like polarity, magnetism, boiling/melting points.

• Essential for studying biological molecules (DNA, proteins), materials (metals, alloys), and reactions.