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Most of the time, chemists can point at a pair of electrons and say where it lives. This pair is the O–H bond in water; that pair is a lone pair on nitrogen. The electrons are localised: tied to one atom or shared between exactly two.
Some electrons refuse to be pinned down like that. They spread across three, six, or even billions of atoms at once. These are delocalised electrons, and once you learn to spot them you will find them explaining things that otherwise look like coincidences: why vinegar is acidic but alcohol isn’t, why graphite conducts but diamond doesn’t, why carrots are orange, and why copper wire works.
A working definition
An electron is delocalised when it occupies an orbital that extends over more than two atoms. The usual cause is a row of p orbitals lying parallel to each other on neighbouring atoms. When three or more such orbitals line up, they overlap as one continuous system, and the electrons in it are shared by every atom in the row.
Think of a bus route with no fixed stops. A passenger in a taxi (a localised electron) goes from A to B and nowhere else. A passenger on a hop-on-hop-off loop can be found anywhere along the route. The loop isn’t longer for the fun of it: by having more room, the electron’s energy drops. This is the key payoff. Spreading electrons over more atoms lowers the energy of the system, which makes the particle more stable.
How to recognise delocalisation
You can usually spot a delocalised system in a structure by looking for one of these patterns:
- Alternating single and double bonds (a conjugated chain or ring), e.g. buta-1,3-diene or benzene.
- A double bond next to an atom with a lone pair, e.g. the C=O next to the O⁻ in a carboxylate ion, or the C=O next to N in an amide.
- A double bond next to a positive charge or an empty p orbital, e.g. an allyl cation.
- A double bond next to an unpaired electron, as in the allyl radical.
In every case, the key condition is that the p orbitals must be able to line up side by side. That generally needs the atoms involved to be roughly in one plane. Twist the molecule so the p orbitals are perpendicular and the delocalisation is switched off.
Another quick test: if you can draw more than one reasonable Lewis structure that differs only in where the electrons are, the molecule has delocalised electrons. Those alternative drawings are resonance structures, and the need for them is a sign that single Lewis structures can’t show the real bonding.
Where delocalised electrons occur
Benzene and aromatic rings
The textbook example. Six p electrons circle a flat ring of six carbon atoms, making every C–C bond equal and giving benzene roughly 150 kJ/mol of extra stability compared with a model with three separate double bonds. The full argument is in bonding in benzene.
Ions such as carbonate and nitrate
Draw a single Lewis structure for the carbonate ion, CO₃²⁻, and you get one C=O and two C–O⁻. That predicts one short bond and two long ones. Measurements show three identical C–O bonds and a flat triangular ion. The two negative charges and the π electrons are spread over all three oxygen atoms. Each bond has a bond order of about 1⅓. The nitrate ion, NO₃⁻, is exactly analogous.
Carboxylate ions and the acidity of carboxylic acids
Here delocalisation does something you can measure with a pH meter. Compare ethanol and ethanoic acid. Both have an O–H group, but ethanoic acid is roughly a hundred billion times more acidic (pKa about 4.8 compared with about 16 for ethanol).
The difference lies in the ions left behind after H⁺ leaves:
- In the ethoxide ion, CH₃CH₂O⁻, the negative charge sits entirely on one oxygen atom. It is localised and therefore high in energy.
- In the ethanoate ion, CH₃COO⁻, the negative charge is shared equally between two oxygen atoms through the O–C–O π system. The two C–O bonds become identical.
A charge shared over two atoms is more stable than the same charge crammed onto one. Stabilising the ion makes it easier for the acid to give up H⁺, so the acid is stronger. Phenol (pKa about 10) sits in between, because the phenoxide ion’s charge can spread partly into the benzene ring, though less effectively than in a carboxylate.
Amides and why they’re barely basic
An amide nitrogen has a lone pair, just like the nitrogen in an amine. Yet amides are almost neutral. The lone pair is delocalised onto the neighbouring C=O group, so it is far less available to pick up H⁺. That same delocalisation gives the C–N bond partial double-bond character, which is why the peptide bond in proteins is flat and doesn’t rotate freely. Protein folding depends on that rigidity.
Conjugated chains and colour
In a conjugated chain of alternating single and double bonds, the π electrons spread along the whole chain. The longer the chain, the smaller the energy gap between the highest filled π orbital and the lowest empty one.
A small gap means a low-energy photon can be absorbed. Short conjugated systems absorb only in the ultraviolet and look colourless. Long ones absorb visible light and look coloured. β-Carotene, with eleven conjugated C=C bonds, absorbs blue light and appears orange, colouring carrots and pumpkins. Many food colourings, textile dyes and indicators work on the same principle. If you have met the photon energy calculator, you can check how absorbed wavelength and energy are linked.
A useful analogy is a guitar string. A longer string vibrates at a lower pitch. A longer conjugated chain “vibrates” at a lower energy, shifting absorption from ultraviolet towards visible light.
Graphite
In graphite, each carbon is bonded to three others in flat hexagonal sheets. The fourth electron of each carbon goes into a p orbital perpendicular to the sheet, and these overlap across the whole layer, giving a delocalised electron system spread over the entire sheet. Those electrons can move when a voltage is applied, so graphite conducts electricity along its layers. In diamond, every carbon uses all four electrons in localised σ bonds, and diamond is an electrical insulator. Same element, different bonding, opposite behaviour (see allotropes of carbon).
Metals
Metallic bonding is delocalisation at its largest scale. The outer electrons of the metal atoms belong to the whole lattice. This “sea” of electrons explains electrical and thermal conductivity, lustre and the ability of metals to bend without shattering. A copper wire is a very long, very well-behaved delocalised electron system.
Summary table
| System | Delocalised over | Visible consequence |
|---|---|---|
| Benzene | 6 carbon atoms in a ring | Equal bond lengths; substitution not addition |
| Carbonate / nitrate | 3 oxygen atoms and the central atom | Three identical bonds, flat ion |
| Carboxylate ion | O–C–O unit | Carboxylic acids much stronger acids than alcohols |
| Amide | O–C–N unit | Weak basicity; flat, rigid peptide bond |
| Conjugated chain | Whole chain | Colour when the chain is long enough |
| Graphite | Whole layer | Conducts along layers |
| Metal | Whole crystal | Conducts electricity and heat; malleable |
Why delocalisation stabilises: two ways to see it
The quantum picture. An electron confined to a small space has a higher minimum kinetic energy than one given more room. This is the same result as the “particle in a box” in physical chemistry: a bigger box gives lower energy levels. Delocalisation gives electrons a bigger box.
The molecular orbital picture. When several p orbitals combine, they form a set of molecular orbitals. The lowest of them spreads over every atom and is lower in energy than an isolated two-atom π bond. Filling the low-lying orbitals of an extended system typically gives a lower total energy than filling several separate two-atom bonds.
Both views agree: delocalisation is not a decoration. It changes energies, and energies drive chemistry.
Common mistakes
- Treating resonance forms as real, separate molecules. The carbonate ion is not switching between three structures. It is one structure all the time, with the electrons spread evenly.
- Thinking delocalised electrons are “free” to leave. In benzene or carbonate, delocalised electrons are still bound to the molecule. Only in metals and graphite do they move through the solid under an applied voltage.
- Moving atoms when drawing resonance. Resonance forms differ only in where electrons are placed. If a hydrogen moves, you have drawn a different compound (a tautomer), not a resonance form.
- Forgetting the geometry requirement. Delocalisation needs overlapping p orbitals, so the atoms must be roughly coplanar. A C=C bond separated from another by two single bonds is not conjugated.
- Saying delocalisation always makes things unreactive. It stabilises, but a delocalised system can still be very reactive. The electron-rich π cloud of benzene is exactly what electrophiles attack.
- Mixing up σ and π electrons. In organic examples the delocalised electrons are π electrons; the σ framework stays localised.
Key takeaways
- Delocalised electrons occupy orbitals spread over three or more atoms, usually through side-by-side overlap of parallel p orbitals.
- Look for alternating single and double bonds, or a double bond next to a lone pair, charge or radical; a need for several resonance structures is another sign.
- Spreading electrons over more atoms lowers their energy and stabilises the molecule or ion.
- Delocalisation explains equal bond lengths in benzene and carbonate, the acidity of carboxylic acids, the weak basicity of amides, colour in conjugated dyes, and conductivity in graphite and metals.
- The atoms involved must be roughly planar for the p orbitals to overlap.
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