Chemical Bonding and Molecular Structure
This domain tests how electrons are shared or transferred, how the resulting Lewis structure is converted into a three-dimensional shape by VSEPR theory, and how bonding and shape together explain polarity and physical properties.
Bond types
- Ionic bonding transfers electrons to form a lattice; melting points are high and conduction occurs only when molten or aqueous. Lattice strength rises with ionic charge and falls with ion radius.
- Covalent bonding shares electron pairs. Increasing bond order shortens the bond and increases bond enthalpy, so C≡C is shorter and stronger than C=C.
- Metallic bonding uses delocalised electrons: malleable, ductile, good conductor. A dative bond (as in NH4+) becomes an ordinary covalent bond once formed.
- Resonance delocalises electrons over equivalent positions, giving bond orders intermediate between the drawn structures, as in CO3^2− and benzene.
VSEPR shapes
- Count electron domains around the central atom; double and triple bonds count as one domain, lone pairs count as one.
- Repulsion order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, so lone pairs compress adjacent angles.
- 2 domains linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 5 trigonal bipyramidal, 6 octahedral. In a trigonal bipyramid, lone pairs occupy equatorial positions.
Polarity
- A molecule is polar only if polar bonds are arranged asymmetrically. CO2 and CCl4 cancel; H2O, NH3 and SO2 do not.
- Electronegativity difference indicates bond polarity; molecular shape decides whether the dipoles add.
Worked example
SF4: sulfur has six valence electrons, four are used in S–F bonds leaving one lone pair, so there are 5 domains - trigonal bipyramidal electron geometry. The lone pair takes an equatorial site, giving a seesaw shape with F–S–F angles near 89° and 117°, both below the ideal 90° and 120° because of lone pair–bond pair repulsion.
Traps that cost marks
- Forgetting the lone pairs on the central atom when counting domains.
- Expanding the octet on a period 2 atom; only period 3 and below can do this.
- Confusing electron-pair geometry with molecular shape when a lone pair is present.
- Assuming a molecule with polar bonds is automatically polar.
- Describing ionic compounds as discrete molecules rather than a lattice.