Chemical Bonding and Molecular Structure
Scope
This domain tests Lewis structures and formal charge, VSEPR geometry, hybridization, resonance and delocalization, polarity and dipole moments, and molecular orbital (MO) theory for diatomics.
Core rules
- Count total valence electrons, then assign formal charge FC = V − N − B/2; the best structure minimizes formal charges and keeps negative charge on the more electronegative atom.
- VSEPR counts every electron domain, including lone pairs and multiple bonds (a double bond counts as one domain). Lone pairs repel more strongly, compressing angles: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°.
- Hybridization follows steric number: 2 → sp, 3 → sp², 4 → sp³, 5 → sp³d, 6 → sp³d².
- MO bond order = (bonding − antibonding electrons)/2. O₂ has two unpaired π* electrons, which explains its paramagnetism.
Worked example
Determine the geometry, hybridization, and polarity of SO₂. Total valence electrons = 6 + 2(6) = 18. Place S centrally with one S=O double bond and one S–O single bond; resonance spreads the double-bond character over both S–O bonds and keeps formal charges near zero. Sulfur has three electron domains - two oxygens and one lone pair - so it is sp² hybridized with a bent geometry near 119°. The two S–O bond dipoles do not cancel, so the molecule is polar.
Traps
- Expanding the octet on period 2 atoms (C, N, O, F); only period 3 and heavier have accessible d orbitals.
- Treating resonance as rapid interconversion rather than a single delocalized description.
- Ignoring lone pairs when predicting geometry, or forgetting that a lone pair can make a symmetric-looking molecule polar.
- Assuming a polar bond always produces a polar molecule; CO₂ and CCl₄ are nonpolar because dipoles cancel.