Chemical Bonding and Intermolecular Forces
What this domain tests
Predicting shapes, polarity and physical properties from bonding, and explaining why substances with similar Mr can have very different boiling points.
Core rules
- VSEPR: electron domains repel; lone pairs repel more than bonding pairs. CH4 109.5°, NH3 107°, H2O 104.5°; PCl5 is trigonal bipyramidal, SF6 octahedral, XeF4 square planar.
- Polar bonds do not guarantee a polar molecule. CO2 is linear and non-polar; H2O is bent and polar.
- Intermolecular forces, weakest to strongest: London dispersion, permanent dipole–dipole, hydrogen bonding. Hydrogen bonds form when H is bonded to N, O or F and a lone pair on a neighbouring N, O or F is available.
- Giant covalent (diamond, graphite, SiO2), ionic and metallic lattices behave very differently on heating from simple molecular solids.
- Lattice enthalpy increases with ionic charge and decreases with ionic radius: MgO is far higher than NaCl.
Worked example
Explain why HF (Mr 20) boils at 19.5 °C but HCl (Mr 36.5) boils at −85 °C. Dispersion forces increase with Mr, which would predict HCl higher. However the H–F bond is very polar and F carries lone pairs, so HF molecules form strong intermolecular hydrogen bonds; HCl has only dipole–dipole and dispersion forces. The hydrogen bonding outweighs the Mr effect.
Traps
Assuming polar bonds always give a polar molecule; saying lone pairs occupy 'less space' than bonds; using Mr alone to predict hydrogen bonding; writing that covalent bonds break when a molecular substance boils - only intermolecular forces break.