17.5
Structure of Benzene: Molecular Orbital Model
According to the molecular orbital (MO) model, benzene has a planar structure with a regular hexagon of six sp2 hybridized carbons. As shown in Figure…
The molecular orbital model describes benzene as a regular planar hexagon with six sp2 hybridized carbons.
Adjacent carbon atoms form σ bonds via sp2–sp2 orbital overlap. Each carbon atom also forms a σ bond with hydrogen by sp2–1s orbital overlap. All the bond angles are 120°.
In addition, each carbon has a half-filled unhybridized 2p atomic orbital perpendicular to the plane of the ring.
These six 2p orbitals combine to form three bonding and three antibonding molecular orbitals.
The bonding orbitals are completely filled by the six π electrons, while the antibonding orbitals are empty, giving benzene a closed-shell configuration that confers stability.
Consequently, the π electron density is delocalized in the form of doughnut-shaped regions above and below the plane of the ring.
Delocalization also explains why just one carbon–carbon bond length is observed for benzene, with a value between typical carbon–carbon single and double bond lengths.
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Q1: Why does benzene have a planar hexagonal structure?
Benzene adopts a planar hexagonal structure because all six carbon atoms are sp2 hybridized. Each carbon forms three σ bonds at 120° angles: two with adjacent carbons and one with hydrogen. The unhybridized 2p orbitals on each carbon remain perpendicular to the ring plane, allowing them to overlap and form π molecular orbitals that stabilize the planar geometry.
Q2: What is the significance of benzene's carbon-carbon bond length?
Benzene's C–C bond length is 139 pm, intermediate between single bonds (154 pm) and double bonds (133 pm). This uniform bond length reflects π electron delocalization across all six carbons. Rather than alternating single and double bonds, the π electrons are distributed equally, creating equivalent bonds throughout the ring and explaining benzene's unusual stability.
Q3: How do the six 2p orbitals in benzene combine to form molecular orbitals?
The six 2p orbitals combine to form six cyclic molecular orbitals: three bonding and three antibonding. The lowest-energy bonding orbital has no nodes with all orbitals in phase. The next two bonding orbitals are degenerate, each containing one nodal plane. The six π electrons completely fill the three bonding orbitals, creating a closed-shell configuration that confers exceptional stability.
Q4: What does delocalization mean in the context of benzene's π electrons?
Delocalization means the six π electrons are not confined to specific bonds but distributed across the entire ring. This creates doughnut-shaped electron density regions above and below the molecular plane. The delocalized π electron system lowers the overall energy of benzene, making it more stable than a structure with localized double bonds would be.
Q5: Why are benzene's bonding molecular orbitals completely filled while antibonding orbitals are empty?
Benzene has six π electrons that occupy the three bonding molecular orbitals (ψ1, ψ2, and ψ3), with each orbital holding two electrons. The three antibonding orbitals (ψ4, ψ5, and ψ6) remain empty. This closed-shell electron configuration, where all bonding orbitals are filled and antibonding orbitals are vacant, results in exceptional stability and is a key feature of aromatic compounds meeting the criteria for aromaticity and the Hückel 4n + 2 rule.
Q6: How do degenerate molecular orbitals differ in benzene compared to linear conjugated systems?
In benzene's cyclic system, degenerate molecular orbitals arise naturally due to the ring's symmetry. The second and third bonding orbitals (ψ2 and ψ3) are degenerate, as are the fourth and fifth antibonding orbitals (ψ4 and ψ5), with nodal planes passing through either bonds or atoms. Linear conjugated systems like 1,3-butadiene lack this degeneracy because they lack the cyclic symmetry that produces equivalent orbital combinations.
Q7: What role does sp2 hybridization play in benzene's structure and bonding?
sp2 hybridization of all six carbons enables benzene's planar geometry and σ-bonding framework. Each sp2 carbon forms three σ bonds at 120° angles, creating the hexagonal ring structure. The unhybridized 2p orbital on each carbon is perpendicular to the molecular plane, allowing overlap with neighboring 2p orbitals to form the delocalized π electron system that stabilizes the molecule.