Chemical Bonding and Molecular Structure
Unit 4 tests Lewis structures, VSEPR geometry, hybridization, bond parameters and a qualitative use of molecular orbital theory.
Core rules
- A covalent bond forms by sharing electrons; Lewis structures are checked using formal charge, and the most stable structure has formal charges closest to zero.
- VSEPR counts bonding pairs plus lone pairs around the central atom; repulsion order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
- Common geometries: 2 domains linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral; lone pairs convert these into bent, pyramidal, see-saw, T-shaped or square pyramidal shapes.
- Hybridization from steric number: 2 → sp, 3 → sp², 4 → sp³, 5 → sp³d, 6 → sp³d².
- Molecular orbital theory gives bond order = (bonding − antibonding electrons)/2 and explains paramagnetism of O₂.
- Hydrogen bonds (F, O, N) raise boiling points and affect solubility.
Worked example
Determine the shape of ClF₃. Chlorine contributes 7 valence electrons; three are used in bonds to fluorine, leaving 4 as two lone pairs. Steric number = 3 + 2 = 5, so hybridization is sp³d. The two lone pairs occupy equatorial positions to minimise repulsion, giving a T-shaped molecule with bond angles slightly below 90°.
Traps
- CO₂ is linear (sp) while SO₂ is bent (sp²) because of the lone pair on sulfur.
- NH₃ (107°) has a smaller bond angle than NF₃? No - NF₃ is smaller, about 102°, because N–F bond pairs lie farther from nitrogen.
- Bond order of O₂⁻ is 1.5 and of O₂²⁻ is 1.
- Resonance structures are not real species; the hybrid is the actual structure.
- Count all valence electrons before drawing a Lewis structure.