Chemical Bonding and Molecular Structure: From Lewis to VSEPR and MO
Core concepts
Candidates must move comfortably between four models. Lewis structures give connectivity and formal charge; check that the sum of formal charges equals the overall charge. VSEPR predicts geometry from electron-domain count: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral, with lone pairs compressing bond angles (CH₄ 109.5° → NH₃ 107° → H₂O 104.5°). Hybridisation follows from the steric number (sp, sp², sp³, sp³d, sp³d²), and questions often ask for the hybridisation of a specific atom in an unusual molecule. Molecular orbital theory explains what Lewis theory cannot: O₂ is paramagnetic because of two unpaired electrons in π* orbitals, and bond order = (bonding − antibonding)/2 predicts stability, as in He₂⁺ (order 0.5) versus He₂ (order 0).
Also revise bond polarity versus molecular polarity (CO₂ is nonpolar despite polar bonds; H₂O is polar), hydrogen bonding and its consequences for boiling points (H₂O, HF, NH₃), and the distinction between sigma and pi bonds.
Traps to watch
- Counting lone pairs as if they were atoms when predicting shape.
- Assuming a molecule with polar bonds must be polar overall.
- Forgetting that bond order can be fractional.
- Treating dative bonds as a separate bond type when they are ordinary covalent bonds.
Worked example
Predict the shape and polarity of SF₄. Sulfur has 6 valence electrons; four S–F bonds use four, leaving one lone pair. Steric number = 5, so the electron geometry is trigonal bipyramidal; the lone pair occupies an equatorial position, giving a see-saw shape. The dipoles do not cancel, so SF₄ is polar. Rule of thumb: count domains first, then place lone pairs equatorially to minimise repulsion.