Lone pairs of electrons exert a stronger repulsive force than bonding pairs, resulting in distortions of the idealized molecular geometry.

In a molecule, both lone pairs and bonding pairs of electrons occupy space around the central atom, but they influence the molecular shape differently. The fundamental distinction is that lone pairs repel more strongly than bonding pairs. This increased repulsion arises because lone pairs are situated closer to the central atom and occupy more spatial volume, leading to greater repulsive interactions.

The Valence Shell Electron Pair Repulsion (VSEPR) theory provides a framework for predicting the geometry of individual molecules based on the number of electron pairs surrounding their central atoms. According to VSEPR theory, electron pairs arrange themselves to minimize repulsion. Bonding pairs, which are shared between two atoms, are more spread out and thus exert less repulsion. In contrast, lone pairs, which are localized around a single atom, create more significant repulsive forces. To delve deeper into this concept, one can examine the relationship between electron arrangements and their corresponding repulsions.

This variation in repulsion leads to distortions in the idealized molecular shapes. For instance, in ammonia ($\text{NH}_3$), there are three bonding pairs and one lone pair of electrons around the nitrogen atom. While VSEPR theory suggests that the molecule should adopt a tetrahedral shape, the stronger repulsion from the lone pair distorts the geometry into a trigonal pyramid.

Moreover, the presence of lone pairs also influences the bond angles in a molecule. Taking water ($\text{H}_2\text{O}$) as an example, the ideal bond angle for a tetrahedral arrangement is $109.5^\circ$. However, due to the presence of two lone pairs on the oxygen atom, the bond angle is reduced to approximately $104.5^\circ$. This alteration in bond angles is a significant aspect of molecular geometry.

In summary, lone pairs of electrons create stronger repulsive forces compared to bonding pairs, leading to observable distortions in a molecule’s shape. This enhanced repulsion, stemming from the proximity of lone pairs to the central atom, ultimately affects both the overall molecular geometry and the bond angles, as illustrated by the differences seen in molecules such as ammonia and water.