Chemical Bonding and Molecular Structure - Quick Revision
A chemical bond is the attractive force that holds atoms/ions together. Bonding lowers the energy of the system to attain stability.
Kossel-Lewis approach
- Lewis symbols show valence electrons as dots; group valence = dots or (8 - dots).
- Octet rule: atoms combine by transfer or sharing of electrons to attain a noble-gas octet.
- Ionic (electrovalent) bond: electrostatic attraction between ions formed by electron transfer (e.g. NaCl, CaF2).
- Covalent bond: sharing of electron pairs (single/double/triple). Lewis dot structures, formal charge.
- Octet exceptions: incomplete octet (LiCl, BeH2, BCl3), odd-electron molecules (NO, NO2), expanded octet (PF5, SF6, H2SO4 - need d orbitals).
Ionic bond and lattice enthalpy
- Favoured by low ionization enthalpy + high negative electron gain enthalpy.
- Lattice enthalpy = energy to separate 1 mole of ionic solid into gaseous ions (NaCl = 788 kJ/mol). Stability depends on lattice enthalpy, not just octet.
Bond parameters
- Bond length: distance between bonded nuclei; covalent radius = half of it.
- Bond angle: angle between bonding orbitals at the central atom.
- Bond enthalpy: energy to break 1 mole of bonds (gaseous); H-H = 435.8 kJ/mol.
- Bond order: number of bonds; higher order -> higher enthalpy, shorter length.
- Resonance: when one Lewis structure is inadequate (O3, CO3 2-, CO2); hybrid is more stable.
- Polarity: nonpolar (similar atoms) vs polar (unequal electronegativity); dipole moment mu = Q x r (Debye). Fajans rules govern covalent character of ionic bonds.
Theories of bonding
- VSEPR: electron pairs repel; shapes from minimising repulsion (lp-lp > lp-bp > bp-bp).
- Valence bond theory: bond by orbital overlap; sigma (head-on) and pi (sidewise).
- Hybridisation: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), sp3d (trigonal bipyramidal), sp3d2 (octahedral).
- Molecular orbital theory: LCAO gives bonding/antibonding MOs; bond order = 1/2 (Nb - Na); explains paramagnetism of O2.
- Hydrogen bond: weak attraction between H (bonded to F, O, N) and an electronegative atom; intermolecular or intramolecular.