Atomic Structure and Chemical Bonding
This section covers Atomic Structure and Chemical Bonding: quantum numbers and orbital shapes, electronic configuration, periodic trends, and the bonding models used to predict shape, polarity and magnetic behaviour. WBCHSE questions often ask you to apply these ideas rather than merely recall them.
Core rules
- Fill orbitals by Aufbau, Pauli and Hund; use the (n+l) rule when orbital energies are close. Cr is [Ar]3d⁵4s¹ and Cu is [Ar]3d¹⁰4s¹.
- VSEPR counts bond pairs plus lone pairs. Lone-pair repulsion is greater than bond-pair repulsion, so NH₃ (107°) has a smaller angle than tetrahedral CH₄ (109.5°).
- Hybridisation follows steric number: 2 = sp, 3 = sp², 4 = sp³. π bonds use unhybridised p orbitals.
- Molecular orbital theory gives bond order = (N_b − N_a)/2 and predicts paramagnetism.
Worked example
Find the bond order and magnetic character of O₂⁻. O₂ has 16 electrons; O₂⁻ has 17. In the MO diagram the extra electron enters a π* orbital. Bonding electrons = 8, antibonding = 5, so bond order = (8 − 5)/2 = 1.5. One unpaired electron remains, so O₂⁻ is paramagnetic. Compare O₂ (bond order 2, two unpaired) and O₂²⁻ (bond order 1, diamagnetic).
Traps
Do not fill 4s before 3d without checking the Cr and Cu exceptions; do not confuse hybridisation with geometry when lone pairs are present; do not use bond order alone to decide polarity; remember that a σ bond is stronger than a π bond, but π bonds often control reactivity.