Chemical Bonding and Molecular Structure — Class 11 Chemistry

This unit explains why atoms combine and how the shapes of molecules arise. The Kössel-Lewis approach relates bonding to the attainment of a stable noble-gas configuration: electron transfer gives ionic (electrovalent) bonds while electron sharing gives covalent bonds, both governed by the octet rule. Lewis dot structures, formal charge and the limitations of the octet rule (incomplete octet, odd-electron molecules and the expanded octet) are discussed. The stability of an ionic bond is measured by lattice enthalpy rather than by octet attainment alone. Bond parameters — bond length, bond angle, bond enthalpy and bond order — are defined alongside resonance (canonical structures and the resonance hybrid) and bond polarity expressed through dipole moment. The valence shell electron pair repulsion (VSEPR) theory predicts molecular geometry from electron-pair repulsions, where lone pair-lone pair repulsion exceeds lone pair-bond pair, which exceeds bond pair-bond pair. Valence bond theory explains bonding through orbital overlap, distinguishing sigma and pi bonds and introducing hybridisation (sp, sp, sp, spd, spd) to account for geometries such as those of methane, ammonia and water. Molecular orbital theory builds bonding and antibonding orbitals by the linear combination of atomic orbitals, giving bond order, stability and magnetic behaviour (explaining why dioxygen is paramagnetic). The unit closes with hydrogen bonding.