Atomic Structure and Periodicity
Scope
The exam probes how electron configuration follows from quantum numbers and how that configuration generates periodic trends: atomic and ionic radius, ionization energy, electron affinity, electronegativity, and the well-known anomalies that break the smooth trends.
Core rules
- Quantum numbers: n (shell), l (subshell, 0 to n−1), m_l (−l to +l), m_s (±½). Pauli's principle forbids two electrons in an atom sharing all four.
- Aufbau filling order, with Hund's rule maximizing parallel spins inside a subshell. Cr ([Ar]3d⁵4s¹) and Cu ([Ar]3d¹⁰4s¹) are the standard exceptions.
- Effective nuclear charge rises left to right while shielding stays roughly constant, so radius falls and ionization energy, electron affinity, and electronegativity rise.
- Down a group, added shells dominate: radius increases, ionization energy falls.
Worked example
Why is the first ionization energy of oxygen (1314 kJ/mol) lower than that of nitrogen (1402 kJ/mol), even though O sits to the right? N is [He]2s²2p³: three p electrons occupy separate orbitals, a half-filled, exchange-stabilized set. O is [He]2s²2p⁴: the fourth p electron must pair in an occupied orbital, adding electron–electron repulsion. Losing an electron relieves that repulsion and restores the stable half-filled 2p³ arrangement, so O ionizes more easily than the trend predicts.
Traps
- Ionizing transition metals removes 4s electrons before 3d, even though 4s fills first.
- Quoting electronegativity trends for noble gases, where the value is not defined.
- Forgetting that ionic radius tracks charge: Fe³⁺ is much smaller than Fe²⁺.
- Missing the electron affinity anomaly: F is smaller than Cl, and its compact 2p orbitals make added-electron repulsion unusually large.