Molecular shape depends on electron pair arrangement

The three-dimensional shape of a molecule or ion depends on the number and arrangement of electron pairs surrounding its central atom. These electron pairs fall into two categories:

• Bonding pairs - Involved in forming covalent bonds with other atoms.
• Lone pairs - Not involved in bonding and remain on the central atom.

For instance, the ammonia molecule (NH₃) has four electron pairs around its central nitrogen atom: three are bonding pairs connected to hydrogen atoms, and one is a lone pair.

Electron pair repulsion theory

A molecule adopts a shape that minimises the repulsion between its electron pairs due to their negative charge:

1. Electron pairs, whether bonding or lone, repel each other.
2. Lone pairs cause more repulsion than bonding pairs because they are closer to the nucleus.
3. Electron pairs arrange themselves as far apart as possible to minimise repulsion.

Repulsion decreases in this order:

Lone pair-lone pair > Lone pair-bonding pair > Bonding pair-bonding pair.

Consequently, lone pairs occupy more space, affecting the molecule's shape by compressing bond angles. This principle underpins the electron pair repulsion theory, which helps predict molecular geometry.

Examples include methane (CH₄), ammonia (NH₃), and water (H₂O), which differ in bond angle and shape despite each having four electron pairs around the central atom due to the varying number of lone pairs.

Shapes of species with multiple bonds

When determining the overall shape of molecules or ions containing multiple bonds, treat each multiple bond as if it were a single electron pair.

Example 1: Carbon dioxide (O=C=O)

• Each C=O double bond is treated as one electron pair.
• The central carbon atom therefore has 2 bonding pairs and 0 lone pairs.
• The CO₂ molecule is linear with a 180° bond angle.

Example 2: Ethene (H₂C=CH₂)

• The C=C double bond is treated as one electron pair.
• Each carbon atom therefore has 3 bonding pairs and 0 lone pairs
• The ethene molecule is trigonal planar with 120° bond angles.

Worked example 1 - Predicting the shape of the OF2 molecule

Let's apply the steps to predict the shape of a oxygen difluoride (OF₂) molecule.

Step 1: Identify the central atom

Oxygen (O) is the central atom.

Step 2: Calculate total outer shell electrons

Oxygen is in group 6, hence it has 6 outer shell electrons

Step 3: Add electrons for each fluorine atom

Each fluorine atom contributes one electron, adding 2 electrons for 2 fluorines

Total electrons = 6 (from O) + 2 (from 2 F) = 8 electrons

Step 4: Calculate total electron pairs

$\text{Total electron pairs }=\frac{8}{2}=4\text{ electron pairs}$

Step 5: Deduce bonding and lone pairs

2 fluorine atoms mean 2 bonding pairs, so 4 - 2 = 2 lone pairs.

Step 6: Predict the molecular shape

With 2 bonding and 2 lone pairs, OF₂ has a bent or non-linear shape with an approximate bond angle of 104.5°.

Worked example 2 - Predicting the shape of the NH4+ ion

Predicting the shape of an ammonium (NH_4_^+^) ion involves similar steps.

Step 1: Identify the central atom

Nitrogen (N) is the central atom

Step 2: Calculate total outer shell electrons

Nitrogen is in group 5, hence it has 5 outer shell electrons

Step 3: Add electrons for each hydrogen atom

Each hydrogen atom contributes one electron, adding 4 electrons for 4 hydrogens

Total electrons = 5 (from N) + 4 (from 4 H) = 9 electrons

Step 4: Adjust for charge on the ion

The ammonium ion has a +1 charge, meaning we subtract one electron

Total electrons after charge adjustment = 9 $-$ 1 = 8 electrons

Step 5: Calculate total electron pairs

$\text{Total electron pairs }=\frac{8}{2}=4\text{ electron pairs}$

Step 6: Deduce bonding and lone pairs

All 4 electron pairs are used for bonding with hydrogen atoms, leaving 0 lone pairs.

Step 7: Predict the molecular shape

With 4 bonding pairs and 0 lone pairs, the NH_4_^+^ ion has a tetrahedral shape with bond angles of approximately 109.5°.

Determining the number of electron pairs

To predict the shape of a molecule or ion, first determine the total number of electron pairs (bonding + lone) on the central atom using these steps:

1. Identify the central atom bonded to all other atoms.
2. Find the number of outer shell electrons of the central atom using its group number.
3. Add one electron for each bonded atom.
4. If the species is an ion, add one electron for each negative charge or subtract one electron for each positive charge.
5. Divide the total number of electrons by two to get the number of electron pairs.
6. Subtract the number of bonds from the number of electron pairs to determine the number of lone pairs.

Based on the electron pair arrangement, the molecular shape can be predicted as follows: