Final answer:
A localized π bond has electrons confined between two atoms while a delocalized π bond's electrons are spread over three or more atoms, as seen in benzene, leading to greater stability and is visualized as a circle within the hexagon.
Step-by-step explanation:
The difference between a localized and a delocalized π bond involves the distribution of π (pi) electrons across different atoms within a molecule. A localized π bond is one where the pi electrons are confined to the space between two adjacent atoms. In contrast, a delocalized π bond involves pi electrons that are spread out over three or more atoms, allowing for an extended area of electron density. This delocalization leads to additional stabilization of the molecule, often referred to as resonance stabilization.
A classic example of delocalization can be seen in benzene, where electrons are free to move around the aromatic ring, depicted with a circle inside the hexagon symbol. This results in a molecule that is more stable and less reactive than its counterpart with just localized bonds. The concept of delocalized bonds is also described by Molecular Orbital Theory and shown graphically with dashed lines or circles in resonance structures like in CO₂.