Source: Organic Chemistry II Post Bacc, Penn State
Difficulty: Intermediate | Prerequisites: General Chemistry (atomic orbitals, electron configuration), Organic Chemistry I (Lewis structures, hybridisation, resonance basics, functional groups).
This topic sits at the heart of Organic Chemistry II. You are moving from drawing Lewis structures and counting bonds into understanding why certain molecules are unusually stable or unstable, based on their electronic structure. Aromaticity explains why benzene resists addition reactions that other alkenes undergo readily, and why nitrogen-containing rings such as pyrrole and pyridine behave so differently from one another.
You should already be comfortable with hybridisation (sp, sp2, sp3), resonance structures, and the idea that electrons in p orbitals can overlap sideways to form pi bonds. If any of that feels shaky, revisit your Organic Chemistry I notes on conjugation and resonance before going further.
Benzene is not three double bonds flipping back and forth; it is one delocalised structure whose six pi electrons sit in an especially favourable set of molecular orbitals, giving it roughly 34 kcal/mol of extra stability. Molecular orbital theory shows that cyclic conjugation creates a unique energy pattern (one very low bonding MO plus two degenerate bonding MOs) that explains this stability far better than resonance drawings alone.
Aromatic
A cyclic, planar, continuously conjugated pi-electron system containing 4n + 2 pi electrons (n = 0, 1, 2, ...). In simple terms, a ring of overlapping p orbitals whose electron count hits a magic number, making it unusually stable.
Resonance hybrid
The actual electronic structure of a molecule, represented by the weighted combination of its resonance structures. Think of it as the real molecule, which is a blend of all the drawings you can write for it, not any single one.
Resonance structures
Different Lewis drawings of the same molecule that differ only in the placement of electrons and pi bonds. They are not separate molecules and do not interconvert.
Kekulé structures
The two alternating-single-and-double-bond drawings historically proposed for benzene. They are the most familiar resonance structures of benzene.
Molecular orbital (MO)
An orbital formed by combining atomic orbitals across multiple atoms. Combining two atomic orbitals always produces two molecular orbitals: one bonding, one antibonding.
Bonding molecular orbital
A molecular orbital lower in energy than the contributing atomic orbitals, formed by constructive (in-phase) overlap. In simple terms, electrons here hold the atoms together.
Antibonding molecular orbital
A higher-energy molecular orbital formed by destructive (out-of-phase) overlap, characterised by at least one node between the atoms. Electrons here work against bonding.
Node
A region or plane where the probability of finding an electron is zero and where the orbital phase changes sign. More nodes means higher energy.
Degenerate orbitals
Orbitals that have the same energy. In benzene, the second and third bonding MOs are degenerate.
Pi bond
A bond formed by sideways overlap of p orbitals, with electron density above and below the plane of the atoms rather than along the bond axis.
Delocalisation
Distribution of pi electrons over several adjacent atoms rather than confinement to one bond. Think of it as electrons spread across the whole ring instead of stuck between two carbons.
Continuous conjugation
A sequence of adjacent atoms with aligned p orbitals capable of overlapping throughout a ring or chain. If one atom breaks the chain (e.g. an sp3 carbon), conjugation stops.
Resonance stabilisation
The extra stability a molecule gains from electron delocalisation across a conjugated system. For benzene, this is approximately 34 kcal/mol.
Kekulé proposed benzene as cyclohexatriene: a six-membered ring with three alternating double bonds.
Ladenburg (Kekulé's student) challenged this model. He pointed out that Kekulé's structure should produce four distinct disubstituted isomers (e.g. xylenes), because positions next to a double bond and next to a single bond would be inequivalent.
Experimentally, only three disubstituted benzene isomers exist: substituents adjacent (ortho), separated by one carbon (meta), or opposite (para).
Kekulé's fix: benzene rapidly equilibrates between two equivalent cyclohexatriene forms, making all positions equivalent on the experimental timescale. This is the origin of the famous "snake biting its tail" story.
Linus Pauling (c. 1920s) reframed the picture. Benzene is not flipping between two structures. The two Kekulé drawings are resonance structures of one single, delocalised molecule.
Resonance structures are not distinct compounds. They are different drawings used to represent the same electronic structure.
The resonance hybrid is the real molecule.
All six C–C bonds in benzene are identical in length: approximately 1.39 angstroms.
For comparison: a typical C–C single bond is about 1.54 angstroms; a typical C=C double bond is about 1.33 angstroms.
Benzene's bond length sits between these values, consistent with each bond having roughly 1.5-bond character.
The circle drawn inside a benzene ring represents this resonance hybrid. In practice, most instructors prefer drawing one Kekulé structure when working through reaction mechanisms.
Cyclohexene (one double bond) releases a known amount of energy when hydrogenated.
If benzene were simply three independent cyclohexene-like double bonds, hydrogenating all three should release roughly three times that energy.
In practice, benzene releases approximately 34 kcal/mol less energy than predicted.
That 34 kcal/mol gap is the resonance stabilisation energy: the extra stability benzene gains from its delocalised electronic structure.
This is a large stabilisation. It is the experimental proof that benzene is not an ordinary cyclohexatriene.
Two atomic orbitals combine to form two molecular orbitals. This is always true: the number of MOs equals the number of AOs combined.
In-phase (constructive) overlap produces a bonding MO: lower energy, increased electron density between atoms, no node between them.
Out-of-phase (destructive) overlap produces an antibonding MO: higher energy, a node between the atoms where electron density is zero.
Key rule: as the number of nodes increases, the orbital's energy increases.
The "top-down method" simplifies drawings. Instead of drawing full 3D p orbitals, you look down on the ring and represent phases as shaded or unshaded circles.
Six p orbitals in a chain produce six MOs: three bonding, three antibonding.
The three bonding MOs are filled by the six pi electrons.
Lowest bonding MO: no nodes, strongly bonding across the entire chain.
Second bonding MO: one node, still overall bonding but less stabilising.
Third bonding MO: two nodes, weakly bonding.
The three antibonding MOs sit above and were not covered in detail.
Benzene also has six p orbitals producing six MOs, but the ring arrangement changes the energy pattern significantly.
Each p orbital interacts with two neighbours (not just one), and the ring closure adds an extra stabilising interaction.
The benzene MO pattern:
One very low-energy bonding MO with no nodes (the "super deluxe" orbital). All six p orbitals overlap constructively around the ring.
Two degenerate bonding MOs at the next energy level, each with one nodal pattern. Still overall bonding.
Three antibonding MOs at higher energy (unoccupied in ground-state benzene).
The six pi electrons fill the three bonding MOs: two electrons in each.
"Degenerate" in this context simply means the two orbitals have identical energy.
The MO picture explains benzene's stability more fundamentally than resonance structures. The especially low-energy first bonding orbital, created by the ring closure, is the key feature that sets cyclic conjugation apart from acyclic conjugation.
Hückel's rule (aromatic): 4n + 2 pi electrons, where n = 0, 1, 2, ... giving counts of 2, 6, 10, 14, ...
Anti-aromatic electron count: 4n pi electrons, where n = 1, 2, 3, ... giving counts of 4, 8, 12, ...
Benzene resonance stabilisation: approximately 34 kcal/mol (from heats of hydrogenation)
C–C single bond length: approximately 1.54 angstroms
C=C double bond length: approximately 1.33 angstroms
Benzene C–C bond length: approximately 1.39 angstroms
Number of MOs = number of AOs combined
Benzene: 6 p orbitals produce 3 bonding MOs + 3 antibonding MOs
Benzene's aromatic stability is the reason it undergoes substitution rather than addition reactions, which is the foundation of much industrial organic synthesis (pharmaceuticals, dyes, polymers). The concept of resonance stabilisation also explains why the circle-in-a-ring notation appears throughout biochemistry, from the aromatic amino acids (phenylalanine, tyrosine, tryptophan) to the nucleotide bases in DNA and RNA.
Students often think the two Kekulé structures of benzene are two separate molecules in rapid equilibrium. They are not. There is only one molecule; the two drawings are resonance structures of the same species.
Students sometimes believe that resonance structures "average out" to give single bonds everywhere. Benzene's bonds are not single bonds; they are intermediate (1.5-bond character), which is different from a single bond.
A common error is assuming that more resonance structures always means more stabilisation. What matters for aromaticity is whether the molecule is cyclic, planar, continuously conjugated, and has the right electron count.
Students frequently confuse the number of MOs with the number of electrons. Six p orbitals give six MOs; the six pi electrons fill only the three bonding MOs.
Be prepared to draw the lowest-energy bonding MO of benzene (all p orbitals in phase, no nodes, electron density distributed constructively around the ring).
You may also be asked to draw one or both of the degenerate bonding MOs (each with one nodal pattern).
Antibonding MOs are lower priority, but know they exist and that they are unoccupied in ground-state benzene.
Expect to calculate resonance stabilisation from heats of hydrogenation data. Know the 34 kcal/mol figure and what it represents.
Be able to explain why benzene's bond lengths are all 1.39 angstroms and what this implies about its electronic structure.
True or false: Benzene's two Kekulé structures represent two different molecules in equilibrium. (False: they are resonance structures of one molecule.)
Fill in the blank: Benzene releases approximately ______ kcal/mol less energy upon hydrogenation than predicted for three isolated double bonds. (34)
True or false: The number of molecular orbitals formed always equals the number of atomic orbitals combined. (True)
Fill in the blank: All six carbon-carbon bonds in benzene have a length of approximately ______ angstroms. (1.39)
True or false: In benzene, the three antibonding MOs are occupied by electrons. (False: only the three bonding MOs are filled.)
Q: Explain why benzene is more stable than a hypothetical cyclohexatriene with three isolated double bonds. What experimental evidence supports this?
A: Benzene's pi electrons are delocalised across the entire ring rather than confined to three separate double bonds. Heats of hydrogenation show that benzene releases approximately 34 kcal/mol less energy than expected for three independent double bonds, indicating substantial stabilisation from delocalisation.
Q: Draw and describe the lowest-energy bonding molecular orbital of benzene.
A: All six p orbitals overlap constructively (in phase) around the ring with no nodes. Electron density is distributed evenly above and below the plane of the ring. This orbital is especially low in energy because the ring closure allows every p orbital to interact with both of its neighbours.
Q: How many molecular orbitals does benzene have, and how are its six pi electrons distributed among them?
A: Benzene has six molecular orbitals (from six p orbitals): three bonding and three antibonding. The six pi electrons fill the three bonding MOs (two electrons each). The antibonding MOs remain empty.
Q: What is the difference between a resonance structure and a resonance hybrid?
A: Resonance structures are individual Lewis drawings that differ in electron placement. The resonance hybrid is the actual molecule, a single species with delocalised electrons that cannot be fully represented by any one drawing. Benzene's true structure is the hybrid; neither Kekulé structure alone is correct.
Q: Why did Ladenburg argue that Kekulé's original benzene structure should give four disubstituted isomers instead of three?
A: In Kekulé's fixed alternating-double-bond structure, two adjacent substituents could sit between a double bond or between a single bond, making those arrangements inequivalent. Experimentally only three isomers exist (ortho, meta, para), which showed that all ring positions are equivalent.
This material connects directly to electrophilic aromatic substitution (EAS), the next major reaction type in most Organic Chemistry II courses: benzene's stability is the reason it undergoes substitution rather than the addition reactions typical of alkenes.
Molecular orbital theory here also sets up the Diels-Alder reaction, where the frontier MOs of 1,3-butadiene become central.
The concept of delocalisation and resonance stabilisation carries into biochemistry, where the aromatic amino acids and nucleotide bases rely on the same electronic principles discussed in this lecture.
Aromaticity, Hückel's rule, 4n+2 rule, benzene structure, resonance hybrid, resonance structures, Kekulé structures, cyclohexatriene, molecular orbitals, bonding MO, antibonding MO, degenerate orbitals, pi bond, pi electrons, delocalisation, electron delocalisation, continuous conjugation, resonance stabilisation, resonance energy, heats of hydrogenation, 1,3,5-hexatriene, node, orbital energy, constructive overlap, destructive overlap, organic chemistry II, OChem II, aromatic stability