How do non-bonding (lone) electron pairs influence molecular shapes?

Non-bonding electron pairs influence molecular shapes by repelling bonding pairs, causing the molecule to adopt a specific geometry.

Non-bonding electron pairs, also known as lone pairs, play a significant role in determining the shape of a molecule. This is due to the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron pairs around a central atom will arrange themselves in a way that minimises repulsion, thus determining the shape of the molecule.

Lone pairs are particularly influential because they occupy more space than bonding pairs. This is because they are located closer to the central atom and are not shared between two atoms, unlike bonding pairs. As a result, they exert a greater repulsive force on the other electron pairs. This repulsion causes the bonding pairs to be pushed closer together, altering the shape of the molecule.

For example, consider a molecule with two bonding pairs and two lone pairs, such as water (H2O). According to the VSEPR theory, the four electron pairs should arrange themselves in a tetrahedral shape for minimal repulsion. However, because the lone pairs exert a greater repulsive force, the two bonding pairs are pushed closer together, resulting in a bent or V-shaped molecule.

Similarly, in ammonia (NH3), which has three bonding pairs and one lone pair, the lone pair pushes the bonding pairs closer together, resulting in a trigonal pyramidal shape instead of the expected tetrahedral shape.

In summary, non-bonding electron pairs significantly influence the shape of a molecule by exerting a greater repulsive force on the bonding pairs, causing them to be pushed closer together. This results in the molecule adopting a specific geometry that minimises electron pair repulsion. Understanding this concept is crucial for predicting the shape and properties of molecules.