Chemical Bonding and Intermolecular Forces
This domain explains how bonding and intermolecular forces determine the shapes, polarity and physical properties of substances. STPM expects VSEPR reasoning and clear distinction between intramolecular and intermolecular forces.
Intramolecular bonding
Ionic bonding is electrostatic attraction between oppositely charged ions in a giant lattice; ionic solids have high melting points and conduct when molten or aqueous. Covalent bonding involves shared electron pairs, forming simple molecules or giant covalent structures. Dative covalent bonds occur when one atom supplies both electrons, as in NH4+ and H3O+. Metallic bonding is a lattice of cations in a sea of delocalised electrons.
VSEPR and molecular shape
Electron pairs repel; lone pairs repel more strongly than bond pairs. Count electron domains: four domains give tetrahedral electron geometry. Three bond pairs plus one lone pair give a trigonal pyramidal molecule; two bond pairs plus two lone pairs give a bent molecule. Bond angles: CH4 109.5°, NH3 about 107°, H2O about 104.5°.
Intermolecular forces
London dispersion forces arise from instantaneous dipoles and increase with more electrons or higher Mr. Permanent dipole–dipole forces occur in polar molecules. Hydrogen bonding occurs when H is bonded to N, O or F. It explains the unusually high boiling points of H2O, NH3 and HF.
Worked example
NH3: nitrogen has five valence electrons, giving three N–H bond pairs and one lone pair. The electron geometry is tetrahedral, but the molecular shape is trigonal pyramidal with a bond angle near 107°. The N–H bonds are polar and the lone pair gives a net dipole, so NH3 is polar and forms hydrogen bonds.
Traps that cost marks
Do not confuse electron-pair geometry with molecular shape. CO2 has polar bonds but is nonpolar overall; CCl4 is also nonpolar. Hydrogen bonding is limited to H–N, H–O and H–F. London forces increase with Mr. Ionic solids do not conduct when solid because the ions are not mobile.