Chemical Bonding and Molecular Structure short notes

Formulas

Formal charge

$\displaystyle \text{FC} = V - L - \tfrac12 B$

V valence, L lone-pair and B bonding electrons.

Bond order

$\displaystyle \text{BO} = \tfrac12(N_b - N_a)$

Nᵦ bonding and Nₐ antibonding electrons.

Dipole moment

$\displaystyle \mu = q\times d$

1 D = 3.336 × 10⁻³⁰ C m.

Hybridisation number

$\displaystyle H = \tfrac12(V + M - C + A)$

V valence electrons, M monovalent atoms, C cation charge, A anion charge.

Shapes (VSEPR)

• BeCl₂: linear (sp, 180°).
• BF₃: trigonal planar (sp², 120°).
• CH₄: tetrahedral (sp³, 109.5°).
• NH₃: pyramidal (sp³, 107°), with one lone pair.
• H₂O: bent (sp³, 104.5°), with two lone pairs.
• PCl₅: trigonal bipyramidal (sp³d).
• SF₆: octahedral (sp³d²).
• Lone pair repulsion is stronger than bond pair repulsion: lp-lp > lp-bp > bp-bp.

Molecular orbital theory

• Up to N₂, the order is σ1s, σ*1s, σ2s, σ*2s, π2p (both), σ2p, π*2p, σ*2p.
• For O₂ and F₂, σ2p comes before the π2p orbitals.
• Bond orders: H₂ is 1, He₂ is 0, N₂ is 3, O₂ is 2, and F₂ is 1.
• O₂ is paramagnetic, with two unpaired electrons.
• Higher bond order means a shorter, stronger bond.

Hydrogen bonding

• It forms when H is bonded to F, O or N.
• Intermolecular hydrogen bonds raise boiling points, as in water.
• Intramolecular ones, as in o-nitrophenol, lower them.