Shapes of molecules and polarity · Grade 12
The shape of a molecule is set by the electron pairs round its central atom. Bonding pairs and lone pairs repel each other and spread out as far apart as possible: this is electron-pair repulsion theory. Lone pairs repel more strongly than bonding pairs, so they squeeze the bond angles. Whether a molecule is polar depends on both its bonds and its shape.
Key facts
- Count the electron pairs round the central atom: take its outer electrons (add one for each negative charge, take one away for each positive charge); one is used for each bond, and the rest pair up as lone pairs.
- 2 pairs: linear, 180°. 3 pairs: trigonal planar, 120°. 4 pairs: tetrahedral, 109.5°. 5 pairs: trigonal bipyramidal, 120° and 90°. 6 pairs: octahedral, 90°.
- Repulsion: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. Each lone pair closes the angle by about 2.5°: NH₃ is trigonal pyramidal (107°) and H₂O is bent (104.5°).
- Name the shape from where the atoms are, not the lone pairs: XeF₄ has 4 bonding pairs and 2 lone pairs and is square planar.
- A double or triple bond counts as one region of electrons: CO₂ is linear.
- A bond is polar when the two atoms have different electronegativities: the more electronegative atom is δ−.
- A molecule is polar if it has polar bonds and its shape does not let the dipoles cancel: H₂O and CHCl₃ are polar; CO₂, BF₃ and CCl₄ have polar bonds but are non-polar.