Chemical Bonding and Molecular Geometry
What this domain tests
This domain checks whether you can build a Lewis structure, translate it into a VSEPR shape and hybridisation, and then decide bond polarity, molecular polarity and bond order. Expect questions on formal charge, resonance, expanded octets and the paramagnetism of O₂.
Core rules
- Count total valence electrons, place the least electronegative atom centrally, then complete octets. Formal charge = (valence electrons) − (non-bonding electrons) − ½(bonding electrons); the best structure minimises formal charges and puts negative charge on the more electronegative atom.
- VSEPR: electron pairs repel in the order lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, so each lone pair closes the angles by roughly 2–2.5°.
- Steric number 2 → linear (sp), 3 → trigonal planar (sp²), 4 → tetrahedral (sp³), 5 → trigonal bipyramidal (sp³d), 6 → octahedral (sp³d²).
- Molecular orbital theory: O₂ has two unpaired electrons in π* orbitals, so it is paramagnetic; bond order = (bonding − antibonding)/2.
Worked example
Determine the shape of XeF₄. Xenon contributes 8 valence electrons, four fluorines contribute 28, giving 36. Form four Xe–F bonds (8 electrons) and place three lone pairs on each F (24 electrons): 32 used, so four electrons remain as two lone pairs on xenon. The steric number is 4 + 2 = 6, an AX₄E₂ system. The two lone pairs occupy axial positions to minimise repulsion, so the molecule is square planar with F–Xe–F angles of 90°, and xenon is sp³d² hybridised.
Traps that cost marks
- A molecule with polar bonds can be non-polar overall: CO₂, CCl₄, BF₃ and SF₆ all have zero dipole moment by symmetry.
- NH₃ (107°) and H₂O (104.5°) have smaller angles than CH₄ (109.5°) purely because of lone pairs, not because of electronegativity.
- Do not confuse formal charge (a bookkeeping device) with oxidation state (electrons assigned to the more electronegative atom).
- SF₄ is a seesaw, not tetrahedral, and ClF₃ is T-shaped.