Chemical Bonding and Molecular Structure
Core concepts
Bonding is driven by achieving a stable electron arrangement. Ionic bonding transfers electrons between metals and non-metals, producing lattices held by electrostatic attraction. Covalent bonding shares electrons; it may be polar or non-polar depending on electronegativity difference. Metallic bonding involves delocalised electrons.
Candidates must predict shapes and polarity using VSEPR theory: count electron domains (bonding pairs plus lone pairs), then apply the sequence linear (2), trigonal planar (3), tetrahedral (4), trigonal bipyramidal (5), octahedral (6). Lone pairs repel more strongly than bonding pairs, compressing bond angles (e.g. water is bent at about 104.5°, ammonia pyramidal at about 107°).
Intermolecular forces - London dispersion, dipole–dipole, and hydrogen bonding - explain boiling points and solubility. Stronger forces mean higher boiling points. Hydrogen bonding requires H bonded to N, O, or F.
Traps to watch
- Judging polarity from bond dipoles alone while ignoring molecular geometry (CO₂ has polar bonds but is non-polar overall).
- Confusing intramolecular bonds with intermolecular forces when explaining physical properties.
- Assuming all covalent compounds are molecular; giant covalent structures (diamond, silica) behave very differently.
- Forgetting that ionic compounds conduct only when molten or dissolved, not solid.
Worked example
Predict the shape of BF₃. Boron has three bonding pairs and no lone pairs, so three electron domains give a trigonal planar shape with 120° angles. Rule of thumb: count domains first, then adjust for lone pairs.