Chemical Bonding, Structure and Properties
Chemical bonding, structure and properties tests whether you can link particle-level models to observable behaviour: melting point, conductivity, solubility and molecular shape.
Core rules
- Ionic compounds form giant lattices of alternating ions; strong electrostatic attractions give high melting points. They conduct only when ions are mobile (molten or aqueous), not when solid.
- Simple molecular substances have weak intermolecular forces between molecules; boiling breaks these forces, not covalent bonds. Giant covalent substances (diamond, graphite, silicon dioxide) have high melting points because many covalent bonds must be broken.
- Graphite conducts because each carbon is bonded to three others, leaving delocalised electrons; diamond does not.
- Metals conduct through delocalised electrons and are malleable because layers of ions slide without breaking the metallic bond.
- VSEPR: electron pairs repel; lone pairs repel more strongly than bonding pairs. CH4 109.5°, NH3 107°, H2O 104.5°, CO2 180°, BF3 120°.
Worked example
Predict the shape and bond angle of NH3. Nitrogen has five valence electrons; three form single bonds to H and one remains as a lone pair. Four electron pairs adopt a tetrahedral arrangement, but the shape is trigonal pyramidal. The lone pair repels the N–H bonding pairs more than they repel each other, reducing the H–N–H angle from 109.5° to about 107°.
Traps
Saying covalent bonds break when a simple molecular substance boils; assuming all covalent substances are non-conductors; ignoring lone pairs when naming a shape; assuming a molecule with polar bonds must be polar overall (CO2 is non-polar because dipoles cancel); confusing the number of electron pairs with the number of atoms bonded.