Part 3: Shapes of molecules & ions

1. Electron pair repulsion

The shape of a molecule or ion depends on the electron pairs around the central atom. These pairs repel each other and spread out as far apart as possible to minimise repulsion. It is the electron pairs that repel, not the atoms.

A lone pair occupies more space around the central atom than a bonding pair, so its repulsion is stronger:

This squeezes bonding pairs closer together, so bond angles become smaller than the ideal angle.

2. Counting regions

Count the regions of electron density around the central atom: each single, double or triple bond counts as one region, and each lone pair counts as one.

Then separate two ideas:

• The electron arrangement uses all regions (bonds and lone pairs).
• The molecular shape describes only where the atoms are.

For example, CO₂ has two C=O double bonds, so two regions around carbon, giving a linear shape. Each double bond is two shared pairs, so CO₂ has four shared pairs in total, but only two regions.

3. The core shapes

4. A counting shortcut (and its limits)

When the central atom is joined only by single bonds to monovalent atoms (H, F, Cl, Br, I), you can count electron pairs like this:

Bonding pairs = number of bonded atoms; lone pairs = total pairs − bonding pairs.

This shortcut does not work for molecules with double bonds (such as CO₂) or with oxygen bonded to the central atom (such as SO₄²⁻). For those, draw a dot-and-cross or Lewis structure and count regions directly.

5. Five and six regions with lone pairs

You may meet unfamiliar molecules. Work them out from the number of bonding pairs and lone pairs. The names are useful but you can always describe the shape in words.

With five regions, lone pairs go in the equatorial positions, where they have more room. With six regions and two lone pairs, the lone pairs go opposite each other to keep as far apart as possible.

6. Worked predictions

XeF₄

Xenon has 8 outer electrons. Four are used in bonds with four F atoms; the remaining four form two lone pairs. Total: 6 regions (4 bonding, 2 lone). The lone pairs sit opposite each other, so the four F atoms lie in a square around Xe: square planar, 90°.

KrF₂

Krypton has 8 outer electrons. Two are used in bonds with F; six form three lone pairs. Total: 5 regions (2 bonding, 3 lone). The three lone pairs take the equatorial positions, leaving the F atoms at the two axial positions: linear, 180°. Show the three lone pairs on Kr in your diagram.

IF₄⁺

Iodine has 7 outer electrons; the + charge removes one, leaving 6. Four are used in bonds; two form one lone pair. Total: 5 regions (4 bonding, 1 lone): seesaw, like SF₄.

7. Drawing shapes

• Solid lines: bonds in the plane of the page.
• Wedges: bonds coming towards you. Dashed wedges: bonds going away.
• Draw lone pairs on the central atom (two dots or a lobe). Leaving them out of a shape drawing is a common error.
• For an ion, put the structure in square brackets with the charge outside.
• Label the bond angle.

8. Quick checks

Finesse practice: indicative answers to check your reasoning, not official mark allocations.

9. Sources