Units 2-3 cover how atoms bond together and the resulting molecular shapes and properties. Understanding IMFs is crucial for explaining physical properties like boiling points and solubility.
Ionic: metal + non-metal, electron transfer, lattice structure. Covalent: non-metals, electron sharing. Metallic: sea of delocalised electrons. Lewis structures: show valence electrons, follow octet rule (exceptions: H=2, expanded octets for Period 3+). Formal charge = valence e⁻ - lone pair e⁻ - ½ bonding e⁻. Best structure: minimize formal charges. Resonance: multiple valid Lewis structures → electrons delocalised.
VSEPR: electron groups repel maximally. Count electron domains (bonds+lone pairs). Linear (2), trigonal planar (3), tetrahedral (4), trigonal bipyramidal (5), octahedral (6). Lone pairs reduce bond angles. Molecular polarity: molecular geometry + bond polarity. Symmetric molecules with polar bonds can be non-polar (CO₂, BF₃, CCl₄).
London dispersion forces (LDF): all molecules, increases with size/surface area. Dipole-dipole: polar molecules. Hydrogen bonding: H bonded to N, O, F interacting with lone pair on another N, O, F. Ion-dipole: ions in polar solvent (dissolving). Solid types: ionic (high MP, brittle), molecular (low MP), metallic (variable MP, conductive), network covalent (very high MP — diamond, SiO₂). Solubility: like dissolves like.
Boiling occurs when molecules gain enough energy to overcome intermolecular forces (IMFs) and escape to the gas phase. Stronger IMFs → more energy needed → higher boiling point. Hierarchy: ion-ion > hydrogen bonding > dipole-dipole > London dispersion (but LDF can dominate for very large molecules). When comparing similar-sized molecules: H-bonding molecules have higher BPs than those with just dipole-dipole or LDF. When molecules have different sizes: larger molecules have stronger LDF. This is why I₂ (large atoms, strong LDF, solid) has a higher melting point than F₂ (small, gas) despite both being non-polar.
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