Chemical Bonding and Molecular Structure
What you must understand
Distinguish ionic, covalent, and metallic bonding, and within covalent bonding distinguish polar from nonpolar using electronegativity differences. Draw Lewis structures, count formal charges, and identify resonance. Apply VSEPR theory to predict electron-domain and molecular geometries, remembering that lone pairs occupy more space than bonding pairs, compressing bond angles. Use valence bond ideas of hybridisation (sp, sp2, sp3, and expanded sets) and molecular orbital theory for bond order and magnetism, especially for diatomics. Know the shapes and polarity consequences of common species.
Traps to watch for
- Reporting electron-domain geometry when the question asks for molecular geometry (water is bent, not tetrahedral).
- Assuming a molecule with polar bonds must be polar overall; symmetry can cancel dipoles, as in CO2 and CCl4.
- Miscounting valence electrons for ions: add electrons for anions, subtract for cations.
- Ignoring that lone pairs reduce bond angles below ideal values, for example 104.5 degrees in water versus 109.5 in methane.
- Using hybridisation labels for molecules better described by delocalised MOs, such as benzene.
Worked example
Predict the shape and polarity of SF4. Sulfur contributes six valence electrons and four fluorines contribute four, giving ten electrons, or five electron domains around sulfur. One domain is a lone pair, so the electron-domain geometry is trigonal bipyramidal and the molecular geometry is a see-saw. The lone pair sits equatorial, and the dipoles do not cancel, so SF4 is polar.
Rule of thumb
Count electron domains first, place lone pairs in the position that minimises repulsion (equatorial in a trigonal bipyramid), then read off the molecular shape and test for symmetry.