Chemical Bonding and Molecular Structure — Notes

Atoms combine to reach a noble-gas configuration: an octet (a duplet for H). A Lewis symbol shows valence electrons as dots. An ionic bond forms by electron transfer; it is favoured by a low ionisation enthalpy of the metal, a negative electron gain enthalpy of the non-metal and a high lattice enthalpy, which rises with higher ionic charge and smaller ionic size (NaCl about 788 kJ mol⁻¹; MgO about 3850 kJ mol⁻¹). A covalent bond is a shared pair.

Lewis structure steps: total valence electrons (adjust for charge); skeleton with the least electronegative atom central; single bonds; complete octets; add multiple bonds if electrons are short; check formal charges. Exceptions: incomplete octet (BeCl₂, BF₃), odd-electron (NO, NO₂), expanded octet (PCl₅, SF₆).

Formal charge = V − L − B/2. In nitrate: N +1, double-bonded O 0, single-bonded O −1 (two), total −1. In ozone (O=O⁺–O⁻): central O +1, terminal O 0 and −1. The preferred structure has the smallest formal charges and puts any negative charge on the more electronegative atom.

Bond length (nuclear distance), bond angle, bond enthalpy (energy to break one mole of bonds in the gas phase) and bond order. For the same atoms: higher order, shorter and stronger (C–C 154 pm and 347 kJ mol⁻¹; C=C 134 and 611; C≡C 120 and 837). Resonance: the real structure is a hybrid; O₃ bond order 1.5, CO₃²⁻ and NO₃⁻ 1.33.

Dipole moment μ = q × d (debye; 1 D = 3.336 × 10⁻³⁰ C m), a vector. HF 1.78 D, HCl 1.07, HBr 0.79, HI 0.38. CO₂, BF₃, CH₄ and CCl₄ are non-polar by symmetry; H₂O 1.85 D; NH₃ 1.47 D, much more than NF₃ 0.23 D because in NF₃ the lone-pair dipole opposes the N–F dipoles.

Electron-pair regions repel: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair; a multiple bond is one region. Lone pairs lower bond angles (CH₄ 109.5°, NH₃ 107°, H₂O 104.5°). In a trigonal bipyramid lone pairs go equatorial; in an octahedron the second lone pair goes trans to the first.

A bond forms by overlap of half-filled orbitals. σ bonds: head-on overlap along the axis; π bonds: sideways overlap of parallel p orbitals. A double bond = 1σ + 1π; a triple bond = 1σ + 2π. Hybridisation (number of hybrid orbitals = steric number): sp linear (BeCl₂, C₂H₂), sp² trigonal planar (BF₃, C₂H₄), sp³ tetrahedral (CH₄, NH₃, H₂O), sp³d trigonal bipyramidal (PCl₅), sp³d² octahedral (SF₆). Ethene has 5σ + 1π; ethyne 3σ + 2π.

Atomic orbitals combine to give bonding and antibonding molecular orbitals. Bond order = (Nb − Na)/2. For Li₂ to N₂ the order is σ2s, σ2s, π2p, σ2p, π2p, σ*2p (s–p mixing); for O₂ and F₂, σ2p lies below π2p. N₂ has bond order 3 and is diamagnetic; O₂ has bond order 2 and two unpaired electrons (paramagnetic). O₂⁺ 2.5, O₂⁻ 1.5, O₂²⁻ 1.

H bonded to F, O or N attracts a lone pair on another F, O or N. Intermolecular H bonding (H₂O, HF, NH₃, alcohols) raises boiling points (H₂O 373 K, H₂S 213 K) and gives ice an open structure, so ice is less dense than water. Intramolecular H bonding (o-nitrophenol) keeps molecules separate, so the compound is more volatile (steam-volatile) than the para isomer.