Shapes of molecules for A-level Chemistry

The shapes of molecules at A-level Chemistry come from electron pair repulsion: the pairs of electrons around a central atom repel each other and spread out as far apart as possible. Count the bonding pairs and lone pairs, and the total tells you the arrangement, from linear (2 pairs, 180°) to octahedral (6 pairs, 90°). Lone pairs repel more strongly than bonding pairs, so each one closes the bond angles by about 2.5°.

Electron pair repulsion theory

AQA's A-level specification treats bonding pairs and lone (non-bonding) pairs of electrons as charge clouds that repel each other. Pairs of electrons in the outer shell of the central atom arrange themselves as far apart as possible to minimise repulsion. The theory is often called VSEPR, short for valence shell electron pair repulsion.

The pairs don't all repel equally. Lone pairs sit closer to the central atom and spread out more, so:

lone pair to lone pair repulsion > lone pair to bond pair > bond pair to bond pair

That order is why molecules with lone pairs have smaller bond angles than the basic shapes. AQA asks you to explain the shapes and bond angles of simple molecules and ions with up to six electron pairs around the central atom.

Shapes of molecules table

The first six rows and octahedral are the ones you will use most. With five pairs, lone pairs always go in the equatorial positions (around the middle), where there is more room, which is what makes SF₄ a seesaw and ClF₃ T-shaped. With six pairs and two lone pairs, the lone pairs sit opposite each other, leaving a square planar molecule.

How to work out the shape of a molecule

1. Find the central atom and its number of outer electrons (its group number for groups 1 to 7, or 8 for xenon).
2. Add one electron for each atom bonded to it by a single bond.
3. For an ion, add one electron for each negative charge, or take one away for each positive charge.
4. Divide by 2 to get the number of electron pairs.
5. The number of bonded atoms is the number of bonding pairs; the rest are lone pairs.
6. Use the table to name the shape, then knock about 2.5° off the basic angle for each lone pair (for 4 pairs).

Worked examples

Example 1: SCl₂

Sulfur has 6 outer electrons, plus 2 from the chlorines = 8, so 4 pairs. Two are bonding and two are lone pairs. The shape is bent, with a bond angle of about 104.5°, like water.

Example 2: the ion PCl₄⁺

Phosphorus has 5 outer electrons, plus 4 from the chlorines = 9, minus 1 for the positive charge = 8, so 4 pairs, all bonding. The shape is tetrahedral, with bond angles of 109.5°.

Example 3: the ion ICl₄⁻

Iodine has 7 outer electrons, plus 4 = 11, plus 1 for the negative charge = 12, so 6 pairs: 4 bonding and 2 lone. The lone pairs sit opposite each other, so the shape is square planar with bond angles of 90°.

Example 4: BrF₅

Bromine has 7 outer electrons, plus 5 = 12, so 6 pairs: 5 bonding and 1 lone. The shape is square pyramidal. The lone pair pushes the four fluorines at the base slightly away from it, so their angles to the top fluorine are a little under 90°.

Try one: predict the shape of a molecule, then press Mark to check it in the Chemical Bonding course.

Bond angles

A bond angle is the angle between two bonds from the central atom. The starting angle comes from the number of electron pairs, which spread as far apart as possible: 180° for two pairs, 120° for three, 109.5° for four, 90° and 120° for five, and 90° for six. Lone pairs then squeeze the angles between the bonds.

Why each lone pair takes off about 2.5°

A lone pair is held closer to the central atom than a bonding pair, so it spreads out more and repels more strongly. It pushes the bonding pairs closer together, which shrinks the angle between the bonds. With four electron pairs, each lone pair reduces the angle by about 2.5°: from 109.5° with none, to about 107° with one, to about 104.5° with two. That rule of thumb gives the values A-level mark schemes expect.

Worked example: the bond angle in NH₂⁻

Nitrogen has 5 outer electrons, plus 1 from each of the two hydrogens, plus 1 for the negative charge: 8 electrons, so 4 pairs. Two are bonding pairs and two are lone pairs. The four pairs start tetrahedral at 109.5°, and the two lone pairs take off about 2 × 2.5° = 5°, so NH₂⁻ is bent with a bond angle of about 104.5°, the same as water.

Explaining a bond angle in the exam

A typical question asks why NH₃ has a smaller bond angle than CH₄. A full answer has three parts:

• Both have four electron pairs around the central atom, arranged roughly tetrahedrally.
• CH₄ has four bonding pairs, so the angle is 109.5°. NH₃ has three bonding pairs and one lone pair.
• The lone pair repels more strongly than a bonding pair, so it pushes the N-H bonds closer together and the angle falls to about 107°.

Water has two lone pairs, so the bonds are squeezed further, to about 104.5°. Real molecules don't always land exactly on these values, but these are the angles A-level mark schemes expect.

Common mistakes

• Ignoring lone pairs. NH₃ isn't trigonal planar: its lone pair makes it trigonal pyramidal.
• Naming the shape from the electron pairs. The shape describes where the atoms are. Water has four pairs in a tetrahedral arrangement, but its shape is bent.
• Counting a double bond as two pairs. It counts as one bonding region.
• Forgetting the charge on an ion. Add electrons for negative charges and remove them for positive ones.
• Writing "about 109°" for tetrahedral. The tetrahedral angle is 109.5°.

Build any of these molecules in 3D in the molecule builder, and see how they bond in the covalent bonding guide. For practice in a short timed paper, try today's A-level Chemistry: Year 1 paper.

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Chemistry · Bonding and structureMolecule builder: build molecules in 3D, check shapesFree 3D molecule builder: build molecules atom by atom, then read the formula, VSEPR shape, bond angles, polarity and a stability check with the reason.