4.1 Kössel-Lewis Approach to Chemical Bonding
- Atoms achieve stability by attaining a noble gas configuration (octet rule).
- Kössel-Lewis theory:
- Atoms share or transfer electrons to complete their valence shells.
- Lewis structures represent valence electrons as dots or lines.
- Example: H₂ → H· + H· → H:H
4.2 Ionic or Electrovalent Bond
- Formed by complete transfer of electrons from a metal to a non-metal.
- Characteristics:
- Usually between metal and non-metal
- High melting and boiling points
- Conduct electricity in molten or aqueous state
- Example: NaCl → Na⁺ + Cl⁻
4.3 Bond Parameters
- Bond Length: Distance between nuclei of two bonded atoms.
- Bond Angle: Angle between two bonds at an atom.
- Bond Energy: Energy required to break one mole of bonds.
- Bond Order: Number of chemical bonds between two atoms.
- Single → 1, Double → 2, Triple → 3
4.4 The Valence Shell Electron Pair Repulsion (VSEPR) Theory
- Electron pairs around a central atom repel each other → determine molecular shape.
- Common Shapes:
- Linear → 180°
- Trigonal Planar → 120°
- Tetrahedral → 109.5°
- Trigonal Bipyramidal → 90° & 120°
- Octahedral → 90°
4.5 Valence Bond Theory (VBT)
- Overlap of atomic orbitals forms a covalent bond.
- Types of overlap:
- σ (sigma) bond: Head-on overlap
- π (pi) bond: Sideways overlap
- Explains:
- Shape of molecules
- Bonding in diatomic and polyatomic molecules
4.6 Hybridisation
- Mixing of atomic orbitals to form equivalent hybrid orbitals.
- Types:
- sp: Linear (180°)
- sp²: Trigonal planar (120°)
- sp³: Tetrahedral (109.5°)
- sp³d: Trigonal bipyramidal (90° & 120°)
- sp³d²: Octahedral (90°)
- Example: CH₄ → sp³ hybridisation
4.7 Molecular Orbital Theory (MOT)
- Atomic orbitals combine to form molecular orbitals (MOs).
- Types of MOs:
- Bonding MO: Lower energy → stable
- Antibonding MO: Higher energy → destabilizes
- Bond Order (MOT):
- Bond Order=2(Electrons in bonding MO – Electrons in antibonding MO)
- Explains magnetism and bond strength in molecules.
4.8 Bonding in Some Homonuclear Diatomic Molecules
- H₂: σ bond (1 bond)
- O₂: σ + π bonds (double bond), paramagnetic
- N₂: σ + 2π bonds (triple bond), diamagnetic
- F₂ & Cl₂: Single σ bond, diamagnetic
4.9 Hydrogen Bonding
- Special dipole-dipole interaction between H and electronegative atoms (F, O, N).
- Types:
- Intermolecular: Between molecules (e.g., H₂O)
- Intramolecular: Within a molecule (e.g., ortho-nitrophenol)
- Effects:
- Higher boiling point
- Increased solubility
- Unique physical properties of water
✅ Quick Summary Table
| Concept | Key Points |
|---|---|
| Ionic Bond | Electron transfer, metal + non-metal, high mp/bp |
| Covalent Bond | Electron sharing, directional, σ & π bonds |
| VSEPR Theory | Predicts molecular shape, electron pair repulsion |
| Hybridisation | Mixing orbitals → hybrid orbitals, determines geometry |
| MOT | Explains bond order, magnetism, stability |
| Hydrogen Bonding | H–F/O/N, intermolecular & intramolecular, affects bp & solubility |