Shapes of molecules and ionsEdexcel International A Level Chemistry: Revision notes

The shape of a molecule or ion is decided by the electron pairs around the central atom. Pairs of electrons are negatively charged, so they repel each other and take up positions that are as far apart as possible, giving minimum repulsion. The shape is named from the arrangement of the atoms, not the lone pairs.

To apply the theory:

1. Count the outer electrons on the central atom (adjust for any charge on an ion).
2. Decide how many are used in bonding pairs and how many remain as lone pairs.
3. A double bond counts as one region of electron density.
4. Arrange the pairs as far apart as possible, then name the shape from the atoms.

Repulsion strength: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. A lone pair is held closer to the central atom, so it is more spread out and repels more.

Bond length is the distance between the nuclei of two bonded atoms. Bond angle is the angle between two bonds at the central atom.

When the central atom has only bonding pairs, the shape follows from the number of pairs:

• 2 pairs: linear, 180°: BeCl₂ (gaseous) and CO₂ (each C=O double bond counts as one region)
• 3 pairs: trigonal planar, 120°: BCl₃, and each carbon atom in C₂H₄
• 4 pairs: tetrahedral, 109.5°: CH₄ and NH₄⁺
• 5 pairs: trigonal bipyramidal, 90° and 120°: gaseous PCl₅
• 6 pairs: octahedral, 90°: SF₆

BeCl₂ and BCl₃ are electron-deficient: beryllium has only four and boron only six electrons in its outer shell after bonding. In PCl₅ and SF₆ the central atom has more than eight electrons (an expanded octet), which is possible because phosphorus and sulfur have empty d orbitals available.

In NH₄⁺ one bond is a dative covalent bond, but once formed all four bonds are identical, so the ion is a regular tetrahedron.

Lone pairs take up space but are not named in the shape. They push the bonding pairs closer together, so the angle is smaller than 109.5°.

• NH₃: three bonding pairs and one lone pair, trigonal pyramidal, H–N–H 107°
• H₂O: two bonding pairs and two lone pairs, bent (V-shaped), H–O–H 104.5°

Both have four electron pairs arranged tetrahedrally, but lone pair–bonding pair repulsion is greater than bonding pair–bonding pair repulsion, so the angle falls by about 2.5° for each lone pair.

The angle rises to 109.5° in the ammonium ion, because the lone pair on NH₃ is used to bond to H⁺ and there is no longer a lone pair to squeeze the bonds.

The same method works for molecules not listed by name. Match the number of bonding pairs and lone pairs to a known shape.

• PH₃ and H₃O⁺ (3 bonding pairs, 1 lone pair): trigonal pyramidal, about 107°
• H₂S and SCl₂ (2 bonding pairs, 2 lone pairs): bent, about 104.5°
• BF₃ (3 pairs, no lone pairs): trigonal planar, 120°; BF₄⁻ and CCl₄ (4 pairs): tetrahedral, 109.5°
• PF₆⁻ (6 pairs): octahedral, 90°

Worked example: the ion NH₂⁻. Nitrogen contributes 5 outer electrons, the two hydrogen atoms 1 each and the negative charge 1, making 8 electrons, which is 4 pairs. Two pairs are bonding and two are lone pairs, so the ion is bent with an angle of about 104.5°.

A full answer has three parts: count the bonding pairs and lone pairs, state that pairs repel and are as far apart as possible, then name the shape and the bond angle.

• State the angle as a number: 180°, 120°, 109.5°, 107°, 104.5°, 90°.
• For molecules with lone pairs, compare lone pair and bonding pair repulsion.
• Explain differences between angles by comparing the number of lone pairs.