Chemical Bonding and Intermolecular Forces
This domain tests the link between electron arrangement and the three-dimensional structure, polarity and physical properties of the substances atoms form.
Bond types and evidence
Ionic bonding is electron transfer giving a giant lattice: high melting point, brittle, conducts when molten or aqueous but not solid. Covalent bonding is shared pairs; draw Lewis structures and check formal charge. Metallic bonding is a lattice of cations in a sea of delocalised electrons: malleable, ductile, conducting. Bond order controls length and strength: triple > double > single, and shorter bonds are stronger.
VSEPR shapes
Count electron domains around the central atom (bonding regions plus lone pairs). Two domains = linear; three = trigonal planar; four = tetrahedral; three plus one lone pair = trigonal pyramidal; two plus two = bent; five = trigonal bipyramidal, with lone pairs occupying equatorial positions (SF4 is a seesaw, ClF3 is T-shaped); six = octahedral, with two lone pairs trans (XeF4 is square planar).
Polarity and intermolecular forces
A molecule is polar only if its polar bonds do not cancel. CO2 is linear, so the dipoles cancel and it is non-polar; SO2 is bent and polar. Intermolecular forces in increasing strength are London dispersion forces, permanent dipole–dipole attraction and hydrogen bonding. Dispersion forces increase with electron count and molecular surface area. Hydrogen bonding requires H covalently bonded to N, O or F, with a lone pair available on the acceptor atom.
Worked example
Explain why water boils at 100 °C but hydrogen sulfide at −60 °C. Both are bent, polar molecules of similar size. Oxygen is far more electronegative than sulfur and carries lone pairs, so each water molecule donates two hydrogen bonds and accepts two. H2S has only weak dipole–dipole and dispersion forces. The extra energy needed to break water's hydrogen-bond network raises the boiling point.
Traps that cost marks
- Claiming hydrogen bonding whenever hydrogen is present; HCl has only dipole–dipole forces.
- Forgetting that C–H bonds are essentially non-polar.
- Describing boiling as breaking covalent bonds; it overcomes intermolecular forces only.
- Miscounting lone pairs on the central atom, which changes the shape completely.
- Ranking hydrogen bonds below dipole–dipole forces; they are weak covalent interactions but strong intermolecular ones.