Name: benzene π orbitals and electron delocalization
Uploaded: 2025-05-22T09:05:38.000Z
Duration: 1 min 18 s
Channel: JoVE

The molecular orbital model describes benzene as a regular planar hexagon with six sp2 hybridized carbons.
Adjacent carbon atoms form &#963; bonds via sp2&#8211;sp2 orbital overlap. Each carbon atom also forms a &#963; bond with hydrogen by sp2&#8211;1s orbital overlap. All the bond angles are 120&#176;.
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 &#960; electrons, while the antibonding orbitals are empty, giving benzene a closed-shell configuration that confers stability.
Consequently, the &#960; 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&#8211;carbon bond length is observed for benzene, with a value between typical carbon&#8211;carbon single and double bond lengths.

structure of benzene: molecular orbital model

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 1, each carbon is bonded to three other atoms with C&#8211;C&#8211;C and H&#8211;C&#8211;C bond angles of 120&#176;. The C&#8211;H bond length is 109 pm, and the C&#8211;C bond length is 139 pm which is midway between the single bond length of sp3 hybridized carbons (154 pm) and sp2 hybridized carbons (133 pm).
<img alt="Benzene structure diagram; shows bond angles, lengths, and hybridization; chemistry education." src="/files/ftp_upload/12451/12451_PT_Figure_1.svg" style="text-align: center; height: 190px;" data-altupdatedat="1754904545000" data-alt="Molecular geometry diagram, benzene hybridization, bond angles, and lengths, sp2, chemical structure." />
The carbon atoms also have an unhybridized 2p atomic orbital with one electron, perpendicular to the plane of the ring, which overlaps with the p orbitals of neighboring carbons, forming a continuous loop of &#960; electrons above and below the plane of the ring. According to molecular orbital (MO) theory, these six 2p orbitals combine to form six &#960; cyclic molecular orbitals (MOs) &#968;1, &#968;2, &#968;3, &#968;4, &#968;5, and &#968;6, as shown in the figure below. Among them, &#968;1, &#968;2, and &#968;3 are bonding, whereas &#968;4, &#968;5 and &#968;6 are antibonding MOs. Generally, the energy of the MOs and the number of nodes increases from &#968;1 to &#968;6, whereas the bonding interaction decreases. However, these MOs of cyclic systems differ from the linear systems such as 1,3-butadiene by having two degenerate MOs.
<img alt="Six p atomic orbital diagram, with energy levels and corresponding wave functions &#936;." src="/files/ftp_upload/12451/12451_PT_Figure_2.svg" style="text-align: center; height: 250px;" data-altupdatedat="1754904545000" data-alt="p Orbital Energy Diagram, atomic orbital splitting, quantum mechanics, orbital energy levels" />
The lowest energy bonding MO, &#968;1, has no nodes where all the orbitals are in phase. The next lowest MO has one node and can be represented in two ways, where the nodal plane passes through a bond or an atom. These two bonding MOs are degenerate and represented as &#968;2 and &#968;3. Similarly, the two nodal planes can pass through bonds or atoms, providing two degenerate antibonding MOs of &#968;4 and &#968;5. The final MO, &#968;6, has the highest energy, with three nodal planes representing the out-of-phase combination of all p orbitals.
The six &#960; electrons that form a closed shell of delocalized &#960; electron density above and below the plane of the ring occupy the three bonding MOs, &#968;1, &#968;2, and &#968;3, whose energies are lower than that of the isolated p orbital, thus leading to unusual stability of benzene.

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 1, ...

structure-of-benzene-molecular-orbital-model

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benzene π orbitals and electron delocalization

Benzene has a planar ring structure with delocalized π electrons. Its six carbon atoms form a regular hexagon, and each carbon is sp2 hybridized. The C–C–C and H–C–C bond angles are 120°, which matches the flat shape of the ring.

Each carbon is bonded to three other atoms. The C–H bond length is 109 pm, and the C–C bond length is 139 pm. That C–C distance is midway between the single bond length of sp3 hybridized carbons, 154 pm, and the sp2 value, 133 pm.

Beneath and above the ring plane, each carbon also has one unhybridized 2p atomic orbital with one electron. These p orbitals stand perpendicular to the ring and overlap with neighboring p orbitals. Together, they form a continuous loop of π electron density above and below the ring.

In molecular orbital theory, the six 2p orbitals combine to make six cyclic π molecular orbitals: ψ1, ψ2, ψ3, ψ4, ψ5, and ψ6. The first three, ψ1 through ψ3, are bonding molecular orbitals. The last three, ψ4 through ψ6, are antibonding molecular orbitals.

The energy of the molecular orbitals rises from ψ1 to ψ6, and the number of nodes also increases. As the nodes increase, the bonding interaction decreases. Cyclic systems differ from linear systems such as 1,3-butadiene because they have two degenerate molecular orbitals, or orbitals with the same energy.

The lowest-energy bonding orbital, ψ1, has no nodes and places all orbitals in phase. The next bonding level has one node and can be drawn in two ways, with the nodal plane passing through a bond or through an atom. These two orbitals are degenerate and are labeled ψ2 and ψ3.

The same idea applies to the antibonding set. Two nodal planes can pass through bonds or atoms, giving two degenerate antibonding orbitals, ψ4 and ψ5. The final orbital, ψ6, has the highest energy and three nodal planes. It represents the out-of-phase combination of all p orbitals.

Benzene has six π electrons, and they occupy the three bonding orbitals, ψ1, ψ2, and ψ3. This creates a closed shell of delocalized π electron density above and below the ring. Because these electrons are in orbitals lower in energy than an isolated p orbital, benzene shows unusual stability.

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Aromatic compounds are linked to both pleasant smells and important chemical structures. The word aromatic first came from fragrant natural sources such as fresh flowers and spices. Early organic chemists isolated benzene derivatives from oils of plants, including vanillin from vanilla, methyl salicylate from wintergreen, and cinnamaldehyde from cinnamon. The history of benzene also helped shape the idea of aromatic compounds. In 1825, Faraday isolated benzene from compressed illuminating gas,...

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Kekule’s benzene model describes benzene as a six-carbon ring with alternating pi bonds. August Kekule proposed this structure in 1865 from the formula C6H6, the idea that all benzene hydrogens are equivalent, and the rule that each carbon has four bonds because of tetravalency. In this model, each carbon is attached to one hydrogen atom. It also predicts two isomers for 1,2-disubstituted benzene derivatives. In one form, the substituted carbons are linked by a double bond, and in the other...

Benzene has a planar ring structure with delocalized π electrons. Its six carbon atoms form a regular hexagon, and each carbon is sp2 hybridized. The C–C–C and H–C–C bond angles are 120°, which matches the flat shape of the ring. Each carbon is bonded to three other atoms. The C–H bond length is 109 pm, and the C–C bond length is 139 pm. That C–C distance is midway between the single bond length of sp3 hybridized carbons, 154 pm, and the sp2 value, 133 pm. Beneath and above the ring plane,...

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Aromatic rings follow H&amp;uuml;ckel’s rule, also called the 4n + 2 rule. This rule helps identify which unsaturated cyclic compounds are aromatic and why some ring systems are unstable instead. Benzene is the classic example, while cyclobutadiene and cyclooctatetraene show different behavior and are not aromatic. Eric H&amp;uuml;ckel, a German chemical physicist, first described the structural features needed for aromaticity. An aromatic compound must be cyclic, planar, and continuously...

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key terms and concepts

Explore benzene’s planar ring, sp2 carbons, and six π molecular orbitals. Learn how delocalized electrons and bonding orbitals give benzene its stability.

Benzene π Orbitals and Electron Delocalization | JoVE K12

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Benzene π Orbitals and Electron Delocalization

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