Lone Pair Repulsion and Molecular Shapes
Lone pair repulsion is the effect by which non-bonding electron pairs repel other electron pairs more strongly than bonding pairs do, which makes bond angles smaller than ideal and determines the final shape of a molecule. In brief: the repulsion order is lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair, which is why the bond angle is 109.5° in CH4, 106.7° in NH3 and 104.5° in H2O.
Predicting molecular shapes in VSEPR theory requires two steps: finding the arrangement of all electron domains, and then deciding which domains are lone pairs. This article explains why lone pairs are larger, how they affect angles and positions, and gives a complete table of the common shapes with up to six electron domains.
Why do lone pairs repel more strongly?
A bonding pair is attracted by two nuclei, so its electron cloud is stretched along the axis between the atoms and is relatively narrow. A lone pair is attracted by only one nucleus, so its cloud is broader and spreads over a wider angle around the central atom. A broader cloud takes up more space and pushes neighboring pairs away more efficiently. Hence the order of repulsion is:
LP–LP > LP–BP > BP–BP
The lone pairs are also the reason for the preferred positions in trigonal bipyramidal molecules, where they occupy the equatorial positions with fewer close contacts at 90°.
How do lone pairs change bond angles?
In a tetrahedral electron geometry, the angle is 109.5° when there are no lone pairs. Each lone pair compresses the angle between the bonded atoms by about 2–3°:
- CH4 (0 lone pairs): 109.5°.
- NH3 (1 lone pair): 106.7°.
- H2O (2 lone pairs): 104.5°.
In ClF3, the two lone pairs in equatorial positions push the two axial fluorines slightly away from 90° to give an angle of about 87.5° between the axial and equatorial bonds. In SF4 the axial F–S–F angle is 173° instead of 180°.
What else affects bond angles?
- Multiple bonds: a double or triple bond has more electron density than a single bond, so it repels more strongly. In H2C=O the H–C–H angle is 116°, and in COCl2 the Cl–C–Cl angle is 111°.
- Electronegativity of the outer atoms: bonding pairs are pulled toward highly electronegative atoms and occupy less space near the central atom. For this reason the F–N–F angle in NF3 (about 102°) is smaller than the H–N–H angle in NH3.
- Size of the central atom: for heavier atoms, such as in PH3 (93.5°) and H2S (92.1°), the angles are close to 90°.
- Size of the outer atoms: large atoms such as I or Br crowd each other and can widen the angle between them.
Shapes with two to four electron domains
| Notation | Domains | Electron geometry | Molecular shape | Example |
|---|---|---|---|---|
| AX2 | 2 | Linear | Linear | CO2 |
| AX3 | 3 | Trigonal planar | Trigonal planar | BF3 |
| AX2E | 3 | Trigonal planar | Bent | SO2 |
| AX4 | 4 | Tetrahedral | Tetrahedral | CH4 |
| AX3E | 4 | Tetrahedral | Trigonal pyramidal | NH3 |
| AX2E2 | 4 | Tetrahedral | Bent | H2O |
Shapes with five and six electron domains
| Notation | Domains | Electron geometry | Molecular shape | Example |
|---|---|---|---|---|
| AX5 | 5 | Trigonal bipyramidal | Trigonal bipyramidal | PCl5 |
| AX4E | 5 | Trigonal bipyramidal | Seesaw | SF4 |
| AX3E2 | 5 | Trigonal bipyramidal | T-shaped | ClF3 |
| AX2E3 | 5 | Trigonal bipyramidal | Linear | XeF2 |
| AX6 | 6 | Octahedral | Octahedral | SF6 |
| AX5E | 6 | Octahedral | Square pyramidal | BrF5 |
| AX4E2 | 6 | Octahedral | Square planar | XeF4 |
How do you determine the shape of a molecule?
- Draw the Lewis structure and count the valence electrons.
- Count the domains on the central atom, with each multiple bond counted once.
- Find the number of lone pairs: E = (valence electrons of A − electrons used in bonds) / 2, adjusted for charge.
- Name the notation AXnEm and look up the shape.
- Adjust the angle by 2–3° per lone pair and consider the effect of multiple bonds.
Worked examples
Example 1: the chlorine trifluoride molecule
ClF3 has 7 + 21 = 28 valence electrons. Three Cl–F bonds use 6, the fluorines take 18 and the remaining 4 form two lone pairs on chlorine. The notation is AX3E2, the electron geometry is trigonal bipyramidal and both lone pairs sit in the equatorial plane. The shape is T-shaped.
Example 2: the chlorite ion
ClO2− has 7 + 12 + 1 = 20 valence electrons. In the structure with two single bonds, the two Cl–O bonds use 4 electrons, each oxygen takes three lone pairs (12 electrons) and chlorine keeps the last 4 as two lone pairs. Chlorine has four domains: AX2E2. The shape is bent with an angle of about 111°.
Example 3: the iodine pentafluoride molecule
IF5 has 7 + 35 = 42 electrons. Five I–F bonds use 10, the fluorines take 30 and the remaining 2 are a lone pair on iodine: AX5E. The shape is square pyramidal.
Which mistakes are common?
- Ignoring lone pairs when naming the shape: the shape describes only the positions of atoms.
- Using 109.5° for ammonia and water: the measured values are lower.
- Forgetting that the electron geometry and shape differ: they are the same only when there are no lone pairs.
- Counting double bonds as two domains: a multiple bond is one domain.
- Placing lone pairs at random: follow the rules for equatorial and opposite positions.