Final answer:
Lone pairs occupy more space around the central atom than bonding pairs due to increased electrostatic repulsion, as they are not shared and hence do not experience the mitigating effects of bonding. This difference affects molecular geometries and is the reason why electron-pair geometries account for all pairs of electrons, both bonding and non-bonding (lone pairs).
Step-by-step explanation:
The question why are lone pairs spread over a larger region of space than bonding pairs? is related to the concept of electron-pair repulsion in molecular geometry. Lone pairs, because they are not shared by atoms, exert repulsion on other electron pairs without the mitigating effect of bonding with other atoms. This results in lone pairs being spread out over a larger region of space to minimize the repulsion between them and bonding pairs. Lone pairs simply occupy more space around the central atom compared to bonding pairs, due to the electrostatic repulsion that is more significant for lone pairs. Additionally, the electron-pair geometry of a molecule takes into account all electrons, showcasing how lone pairs affect shapes of molecules differently compared to bonding pairs.
Considering the order of repulsion, which determines how much space different regions of electrons occupy, the sequence from largest to smallest is: lone pair > triple bond > double bond > single bond. For instance, in a compound like ammonia (NH3), the lone pair on the nitrogen atom occupies a larger region of space than the bonding pairs shared with hydrogen atoms, which can be visualized in electron distribution and electrostatic potential models.