Chemical bonding and intermolecular forces
This domain asks you to move from a molecular formula to a Lewis structure, then to geometry, hybridisation and polarity, and finally to macroscopic properties such as melting point, volatility and conductivity. Each step must be consistent with the previous one.
Lewis structures and formal charge
Count valence electrons, place the least electronegative atom centrally, form single bonds, complete octets, then add multiple bonds if electrons remain. Formal charge = (valence e⁻) − (lone-pair e⁻) − ½(bonding e⁻); the best structure minimises formal charges and puts negative charge on the more electronegative atom. Exceptions: BeCl₂ and BF₃ are electron-deficient, NO is odd-electron, and period-3 elements (PCl₅, SF₆) can expand the octet.
VSEPR geometry
Count electron domains (bonds + lone pairs): 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral. Lone pairs repel more strongly than bonding pairs, compressing angles: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°.
Hybridisation and polarity
Hybridisation follows the domain count: sp (2), sp² (3), sp³ (4). A molecule is polar when the vector sum of bond dipoles plus lone-pair contributions is non-zero. CO₂ (linear, opposed dipoles) is non-polar; H₂O (bent) is polar.
Intermolecular forces and solids
Ion–dipole > hydrogen bonding > dipole–dipole > London dispersion, though dispersion grows with polarisability and dominates in I₂ and Br₂. Hydrogen bonds require H bonded to N, O or F. Ionic solids melt high and conduct when molten; molecular solids are soft and volatile; network covalent solids (diamond, SiO₂) are extremely hard; metals conduct as solids via delocalised electrons.
Worked example
For SO₃²⁻, valence electrons = 6 + 3(6) + 2 = 26. Sulphur is central; three S–O bonds use 6 electrons and the remaining 20 sit as lone pairs, one on sulphur. Four domains → sp³ → trigonal pyramidal, and the dipoles do not cancel, so the ion is polar.
Traps
- Forgetting to add charge when counting electrons for ions.
- Counting the lone pair but calling the shape tetrahedral instead of trigonal pyramidal.
- Claiming C–H bonds are polar; they are essentially non-polar.
- Explaining boiling points with intramolecular bonds rather than intermolecular forces.