Molecular Geometry & Bonding Theories

Identical atoms → nonpolar bond; different electronegativity → polar bond. A small nonzero ΔEN (e.g. Cl–Br) still gives a real bond dipole even though the ΔEN classification bin calls it 'nonpolar covalent'.

On the central atom, count domains = bonding pairs (a single, double, or triple bond is ONE domain each) + lone pairs + any single-electron radical.

Names the arrangement of BONDS only (lone pairs are invisible to the shape). Drop lone-pair domains from the parent electron geometry: e.g. bent, trigonal pyramidal, see-saw, T-shaped, square planar, square pyramidal.

Lone pairs repel more than bonding pairs, compressing angles below ideal: H₂O 104.5° and NH₃ 107° (vs 109.5°). More lone pairs → more compression.

Trigonal bipyramidal: lone pairs take equatorial sites (fewer 90° neighbours) → see-saw, T-shaped, linear. Octahedral: two lone pairs go trans (opposite) → square planar.

Symmetric arrangement of identical polar bonds → dipoles cancel → nonpolar (CO₂, BF₃, CH₄, PF₅, SF₆). Asymmetric (usually a lone pair on the central atom) → polar (H₂O, NH₃, SO₂, SCl₄).

Single = 1 σ; double = 1 σ + 1 π; triple = 1 σ + 2 π. Every bonded connection (including each C–H) is one σ; π count comes only from double and triple bonds.

Bond order = (bonding e⁻ − antibonding e⁻) / 2. Zero means no stable bond (Ne₂). Half-integers occur for odd species: O₂⁻ = 1.5, O₂⁺ = 2.5.

For B₂, C₂, N₂ (and CO), s-p mixing puts π₂p BELOW σ₂p. For O₂, F₂, Ne₂ the order is conventional: σ₂p below π₂p.

All electrons paired → diamagnetic (weakly repelled). Any unpaired electron → paramagnetic (attracted). O₂ has 2 unpaired electrons in π*₂p → paramagnetic.