Atomic and Molecular Structure
What the exam expects
You must connect atomic structure to the periodic table and to bonding. Key concepts: subatomic particles and isotopes; electron configuration and orbital filling (Aufbau, Pauli, Hund); periodic trends (atomic radius, ionization energy, electronegativity, electron affinity); ionic vs. covalent vs. metallic bonding; Lewis structures, VSEPR geometry, and polarity; and intermolecular forces.
Traps to watch for
- Isotopes differ in neutrons, not protons or electrons. Average atomic mass is a weighted average of isotope masses.
- Electron configuration exceptions (e.g., Cr, Cu) and the order of orbital filling.
- Periodic trends: atomic radius decreases across a period and increases down a group; ionization energy and electronegativity increase across and decrease down.
- Polarity: a molecule can have polar bonds but be nonpolar overall if geometry cancels dipoles (CO₂ vs. H₂O).
- VSEPR shapes - lone pairs repel more strongly than bonding pairs, compressing bond angles.
- IMF strength ordering: hydrogen bonding > dipole-dipole > London dispersion (for comparable size).
Rule of thumb
Count valence electrons, draw the Lewis structure, count electron domains (bonding + lone pairs), then assign geometry. Polarity requires both polar bonds and asymmetric geometry.
Worked example
Water: O has 6 valence electrons, plus 2 H = 8. Lewis structure gives O with two bonds and two lone pairs - four electron domains → tetrahedral electron geometry, bent molecular geometry. The O–H bonds are polar and the bent shape does not cancel dipoles, so H₂O is polar and forms hydrogen bonds, giving it a high boiling point.