Chapter 9: Problem 60
How does a delocalized molecular orbital differ from a molecular orbital such as that found in \(\mathrm{H}_{2}\) or \(\mathrm{C}_{2} \mathrm{H}_{4}\) ? What do you think are the minimum conditions (e.g. number of atoms and types of orbitals) for forming a delocalized molecular orbital?
Short Answer
Expert verified
Delocalized orbitals extend over multiple atoms and require conjugated systems. Localized orbitals are confined between two atoms.
Step by step solution
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Localized Molecular Orbitals
Localized molecular orbitals are those where electrons are confined mainly between two specific atoms. This confinement forms strong covalent bonds, such as the sigma (\(\sigma\)) bond found in diatomic hydrogen (\(\text{H}_{2}\)). In these molecules, electron clouds are well-defined and do not spread beyond the immediate vicinity of the bonded atoms.
Localized orbitals are simple and occur in smaller, less complex molecules. In ethylene (\(\text{C}_{2}\text{H}_{4}\)), for instance, the bonds are tightly held between specific pairs of atoms, like the \(\pi\) bond formed by the side-to-side overlap of \(p\) orbitals. This keeps electrons localized to the region between the bonded atoms, creating a stable structure.
Localized orbitals are simple and occur in smaller, less complex molecules. In ethylene (\(\text{C}_{2}\text{H}_{4}\)), for instance, the bonds are tightly held between specific pairs of atoms, like the \(\pi\) bond formed by the side-to-side overlap of \(p\) orbitals. This keeps electrons localized to the region between the bonded atoms, creating a stable structure.
Conjugated Systems
Conjugated systems are fascinating structures where alternating single and double bonds allow for electron sharing across multiple atoms.
These systems typically consist of three or more adjacent atoms, each having a \(p\) orbital capable of overlapping with its neighbors. This arrangement allows for the formation of delocalized molecular orbitals, spreading out electron density over the entire chain or ring structure.
In molecules like benzene (\(C_6H_6\)), the arrangement of alternating single and double bonds creates a continuous path for electrons to travel. The resulting stability from such arrangements is often associated with lower energy, making these structures very unique and integral to many aromatic compounds.
These systems typically consist of three or more adjacent atoms, each having a \(p\) orbital capable of overlapping with its neighbors. This arrangement allows for the formation of delocalized molecular orbitals, spreading out electron density over the entire chain or ring structure.
In molecules like benzene (\(C_6H_6\)), the arrangement of alternating single and double bonds creates a continuous path for electrons to travel. The resulting stability from such arrangements is often associated with lower energy, making these structures very unique and integral to many aromatic compounds.
p Orbital Overlap
For molecular orbitals to become delocalized, effective \(p\) orbital overlap is crucial.
When \(p\) orbitals on adjacent atoms align correctly, they can interact and overlap side-to-side, forming \(\pi\) bonds and eventually delocalized \(\pi\) systems. These overlaps allow electrons to be shared and move freely across multiple atoms.
A continuous overlap across a chain of \(p\) orbitals is essential for delocalization. Without this continuous overlap, electrons remain localized, confined between specific atoms, and do not contribute to the conjugated system's unique properties. For example, in benzene, each carbon atom's \(p\) orbitals overlap, creating a circular array of overlapping \(\pi\) orbitals.
When \(p\) orbitals on adjacent atoms align correctly, they can interact and overlap side-to-side, forming \(\pi\) bonds and eventually delocalized \(\pi\) systems. These overlaps allow electrons to be shared and move freely across multiple atoms.
A continuous overlap across a chain of \(p\) orbitals is essential for delocalization. Without this continuous overlap, electrons remain localized, confined between specific atoms, and do not contribute to the conjugated system's unique properties. For example, in benzene, each carbon atom's \(p\) orbitals overlap, creating a circular array of overlapping \(\pi\) orbitals.
Sigma and Pi Bonds
Understanding sigma and pi bonds is fundamental in grasping the concept of molecular orbitals. Sigma (\(\sigma\)) bonds form when atomic orbitals overlap head-on, which results in a robust and circular bonding region. These bonds are generally localized between two atoms and are the first bonds to form between interacting atoms.
Pi (\(\pi\)) bonds, on the other hand, are formed from the side-by-side overlap of \(p\) orbitals. These bonds can spread over several atoms under the right conditions, contributing to delocalization. When conjugated systems are present, \(\pi\) bonds facilitate electron movement across a molecule, enhancing stability and lowering energy. Both types of bonds are essential for understanding not only localized molecular structures but also what makes delocalized systems like benzene unique.
Pi (\(\pi\)) bonds, on the other hand, are formed from the side-by-side overlap of \(p\) orbitals. These bonds can spread over several atoms under the right conditions, contributing to delocalization. When conjugated systems are present, \(\pi\) bonds facilitate electron movement across a molecule, enhancing stability and lowering energy. Both types of bonds are essential for understanding not only localized molecular structures but also what makes delocalized systems like benzene unique.
