| 1.3.3 Shapes of Molecules |
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Electron-pair repulsion theory predicts the three-dimensional shape of a covalent molecule or polyatomic ion.
Electron pairs around a central atom repel each other and arrange themselves as far apart as possible to minimise repulsion.
A multiple bond (double or triple) counts as one region of electron density, because its bonding pairs all lie in the same direction in space.
This is why CO₂ is linear and each carbon in C₂H₄ is trigonal planar.
Lone pairs repel more strongly than bonding pairs, because they are held closer to the central atom.
The repulsion order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, and each lone pair compresses the bond angle by approximately 2.5°.
Key Definition Bond length is the distance between the nuclei of two bonded atoms.
Key Definition Bond angle is the angle between two covalent bonds that share a common atom.
Both bond length and bond angle depend on the arrangement of electron pairs around the central atom.
The ten specified species span the major geometries:
- Linear — BeCl₂ and CO₂; bond angle 180°
- Trigonal planar — BCl₃ and C₂H₄; bond angle 120°
- Tetrahedral — CH₄ and NH₄⁺; bond angle 109.5°
- Trigonal pyramidal — NH₃; bond angle 107°
- Bent (non-linear) — H₂O; bond angle 104.5°
- Trigonal bipyramidal — PCl₅; bond angles 90° and 120°
- Octahedral — SF₆; bond angle 90°
Exam questions frequently ask for the shapes of analogous species, and the method is always the same:
- Count the bonding regions, treating each multiple bond as a single region.
- Count the lone pairs on the central atom.
- Match the totals to the corresponding standard geometry.
- Adjust the predicted bond angle for lone-pair compression (about 2.5° per lone pair).
Precision matters in marking: exact bond angles must be stated, and shape names must use the full standard term, especially "trigonal pyramidal" rather than just "pyramidal".