# Pi bond
A pi bond (π bond) is a covalent bond formed by the side-by-side (sideways) overlap of p orbitals on two adjacent atoms, with bonding electron density lying above and below the internuclear axis rather than along it[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation) • [2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds). It is the second bond of a double bond and one of the two extra bonds of a triple bond, and it is the source of the restricted rotation and stereoisomerism that distinguish multiple bonds from single bonds[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation).
| Key fact | Detail |
|---|---|
| Formation | Sideways overlap of two parallel p orbitals on adjacent atoms; the bonding electron density lies above and below the internuclear axis[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation) • [2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds) |
| Defining symmetry | A π bond has a nodal plane that includes the internuclear bond axis; a σ bond has no such nodal plane (IUPAC definition)[3](https://goldbook.iupac.org/terms/view/S05434) |
| Bond counting | A single bond is one σ bond; a double bond is one σ plus one π; a triple bond is one σ plus two π[2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds) |
| Strength | Weaker than the σ bond between the same atoms, because the π electrons sit farther from the nuclei[4](https://sites.science.oregonstate.edu/~gablek/CH334/Chapter1/ethylene_trajectory/bare_Bonds.htm) |
| Rotation barrier | Rotation about a double bond requires breaking the π bond; estimates place the π bond energy of ethene at roughly 60 to 66 kcal/mol[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation) • [5](https://www.sciencedirect.com/org/science/article/pii/S0022326324014816) • [6](https://pmc.ncbi.nlm.nih.gov/articles/PMC10462764/) |
| Consequence | Restricted rotation makes cis–trans (E/Z) stereoisomers of alkenes isolable compounds[7](https://openstax.org/books/organic-chemistry/pages/7-4-cis-trans-isomerism-in-alkenes) |
| Delocalization | Adjacent aligned p orbitals overlap across intervening σ bonds, delocalizing π electrons in conjugated systems[8](https://chem.libretexts.org/Courses/Purdue/Chem_26505%3A_Organic_Chemistry_I_%28Lipton%29/Chapter_1._Electronic_Structure_and_Chemical_Bonding/1.10%3A_Pi_Conjugation) |
## Formation and geometry
In the valence bond picture, a π bond arises when two half-filled p orbitals on neighboring atoms are parallel and overlap laterally. In ethene (C₂H₄), each carbon is sp² hybridized: three sp² hybrids at angles of 120° form the four C–H σ bonds and the C–C σ bond, while the one remaining unhybridized 2p orbital on each carbon, perpendicular to the molecular plane, overlaps side by side to form the π bond[2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds). The molecule is planar because planarity is the arrangement in which the p orbitals overlap most efficiently; tilting the two sp² frameworks relative to each other would reduce the overlap[2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds).
The resulting electron distribution differs sharply from that of a σ bond. A σ bond concentrates electron density along the line between the two nuclei, whereas the π bond's electron cloud is split into two lobes, one above and one below the molecular plane, with a node along that plane[4](https://sites.science.oregonstate.edu/~gablek/CH334/Chapter1/ethylene_trajectory/bare_Bonds.htm). Because the electrons in the π orbital sit farther from the two nuclei than those in the σ orbital, the π interaction is weaker[4](https://sites.science.oregonstate.edu/~gablek/CH334/Chapter1/ethylene_trajectory/bare_Bonds.htm).
In molecular orbital language, the side-by-side combination of two p orbitals produces a π bonding molecular orbital and a π* antibonding molecular orbital; the bonding combination has no node between the nuclei, while the antibonding combination has a node there[9](https://openstax.org/books/chemistry/pages/8-4-molecular-orbital-theory). In a triple bond such as acetylene (ethyne), each sp-hybridized carbon has two unhybridized p orbitals, oriented perpendicular to each other, and each pair overlaps side by side to give two π bonds at right angles; the two π bonding orbitals are degenerate, differing only in orientation[2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds) • [9](https://openstax.org/books/chemistry/pages/8-4-molecular-orbital-theory).
The familiar σ-plus-π description is a model, not a unique decomposition. The alternative bent-bond picture, in which three equivalent curved bonds describe a double bond, divides the same total electron density differently; the two drawings are dot-for-dot equivalent, and either may be used as convenient[10](https://chem.libretexts.org/Bookshelves/General_Chemistry/ChemPRIME_%28Moore_et_al.%29/07%3A_Further_Aspects_of_Covalent_Bonding/7.08%3A_Sigma_and_Pi_Bonds).
## Restricted rotation and stereoisomerism
Rotation about a single (σ) bond is nearly free because the end-to-end [orbital overlap](https://www.edgepedia.org/orbital-overlap) does not depend on the relative orientation of the orbitals; the bonding electron density is symmetric about the axis, so turning one end changes nothing[2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds). Rotation about a double bond is different: turning one end twists one p orbital out of alignment with the other, destroying the sideways overlap and breaking the π bond[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation). The barrier to rotation must therefore be at least as great as the strength of the π bond itself[7](https://openstax.org/books/organic-chemistry/pages/7-4-cis-trans-isomerism-in-alkenes).
The magnitude of that barrier has been estimated in several ways. The physical organic chemistry literature gives the π bond energy of ethene as about 60 kcal/mol, enough that random thermal energy cannot supply it, so rotation around double bonds ordinarily does not occur[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation). A bond dissociation energy analysis gives Eπ = 65.2 ± 0.1 kcal/mol for ethene, obtained as the difference between the C–H bond dissociation energies of ethane (101.0 ± 0.1 kcal/mol) and ethene (35.8 ± 0.1 kcal/mol), and corroborated by the singlet–triplet gap of 66.0 kcal/mol, since triplet ethene has a fully twisted D₂d geometry[5](https://www.sciencedirect.com/org/science/article/pii/S0022326324014816). An experimental determination of the rotational barrier gives 65.9 ± 0.9 kcal/mol, with the singlet–triplet gap measured at 58 ± 3 kcal/mol[6](https://pmc.ncbi.nlm.nih.gov/articles/PMC10462764/). For comparison, the barrier to rotation in ethane is only 12 kJ/mol[7](https://openstax.org/books/organic-chemistry/pages/7-4-cis-trans-isomerism-in-alkenes).
Because rotation does not occur, alkenes with suitable substitution exist as distinct, isolable stereoisomers. In 2-butene the two methyl groups can sit on the same side of the double bond (cis) or on opposite sides (trans); the two forms cannot interconvert spontaneously and are different compounds[7](https://openstax.org/books/organic-chemistry/pages/7-4-cis-trans-isomerism-in-alkenes). [Cis–trans isomerism](https://www.edgepedia.org/cis-trans-isomerism) is possible whenever both double-bond carbons are each attached to two different groups; if either carbon bears two identical groups, the isomerism disappears[7](https://openstax.org/books/organic-chemistry/pages/7-4-cis-trans-isomerism-in-alkenes). The E/Z designations used for more general cases rely on the [Cahn–Ingold–Prelog priority rules](https://www.edgepedia.org/cahn-ingold-prelog-priority-rules)[1](https://www.sciencedirect.com/topics/chemistry/bond-rotation). Rotation can be forced photochemically, and the light-induced isomerization of cis alkenes to their trans isomers plays an important role in the chemistry of vision[11](https://s3.lite.msu.edu/res/msu/botonl/b_online/library/newton/Chy251_253/Lectures/Geometric_Isomers/GeometricIsomers.html).
## Delocalization and conjugation
A π bond need not be confined to two atoms. When p orbitals on three or more adjacent atoms are aligned, they overlap side by side across the intervening σ bonds, and the π electrons delocalize over the whole array; such a system is called conjugated, and delocalization in general lowers the molecule's energy and increases its stability[8](https://chem.libretexts.org/Courses/Purdue/Chem_26505%3A_Organic_Chemistry_I_%28Lipton%29/Chapter_1._Electronic_Structure_and_Chemical_Bonding/1.10%3A_Pi_Conjugation). Lone pairs, radicals, and carbenium ions can participate alongside ordinary π bonds[8](https://chem.libretexts.org/Courses/Purdue/Chem_26505%3A_Organic_Chemistry_I_%28Lipton%29/Chapter_1._Electronic_Structure_and_Chemical_Bonding/1.10%3A_Pi_Conjugation). In benzene, the six p orbitals of the ring overlap cyclically, and the six π electrons are delocalized over the entire ring rather than located in three localized double bonds[2](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds).
Delocalized π systems also govern light absorption. When an electron in a conjugated system absorbs a photon of the right wavelength, it is promoted to a higher energy level, typically a π-to-π* transition; systems with fewer than eight conjugated double bonds absorb only in the ultraviolet and are colorless, while longer systems absorb at longer wavelengths and appear yellow to red[8](https://chem.libretexts.org/Courses/Purdue/Chem_26505%3A_Organic_Chemistry_I_%28Lipton%29/Chapter_1._Electronic_Structure_and_Chemical_Bonding/1.10%3A_Pi_Conjugation). The largest conjugated π systems occur in graphene, graphite, conductive polymers, and carbon nanotubes[8](https://chem.libretexts.org/Courses/Purdue/Chem_26505%3A_Organic_Chemistry_I_%28Lipton%29/Chapter_1._Electronic_Structure_and_Chemical_Bonding/1.10%3A_Pi_Conjugation).
## Reactivity
The π bond is the reactive face of a multiple bond. Because its electron density lies off the bond axis and is more loosely held than σ electron density, π bonds are the sites where electrophilic attack begins; essentially every addition reaction of organic chemistry, including hydrogenation, halogenation, hydration, and hydrohalogenation, starts with an electrophile attacking the π bond[12](https://premiermcatprep.com/mcat-books/organic-chemistry/bonding/pi-bonds).
## References
1. [Isomerism about Double Bonds, in *Encyclopedia of Physical Science and Technology* (Charles L. Perrin), ScienceDirect Topics](https://www.sciencedirect.com/topics/chemistry/bond-rotation)
2. [5.3 Multiple Bonds, *Chemistry: Atoms First*, OpenStax](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds)
3. [σ, π (sigma, pi), IUPAC Gold Book, S05434](https://goldbook.iupac.org/terms/view/S05434)
4. [Hybrid Atomic Orbitals: Bond Formation in Ethylene, Oregon State University CH334](https://sites.science.oregonstate.edu/~gablek/CH334/Chapter1/ethylene_trajectory/bare_Bonds.htm)
5. [π-Bond Dissociation Energies: C–C, C–N, and C–O, The Journal of Organic Chemistry (ScienceDirect)](https://www.sciencedirect.com/org/science/article/pii/S0022326324014816)
6. [Synthesis, structural analysis, and properties of highly twisted alkenes, PMC](https://pmc.ncbi.nlm.nih.gov/articles/PMC10462764/)
7. [7.4 Cis–Trans Isomerism in Alkenes, *Organic Chemistry*, OpenStax](https://openstax.org/books/organic-chemistry/pages/7-4-cis-trans-isomerism-in-alkenes)
8. [1.10: Pi Conjugation, Chemistry LibreTexts](https://chem.libretexts.org/Courses/Purdue/Chem_26505%3A_Organic_Chemistry_I_%28Lipton%29/Chapter_1._Electronic_Structure_and_Chemical_Bonding/1.10%3A_Pi_Conjugation)
9. [8.4 Molecular Orbital Theory, *Chemistry*, OpenStax](https://openstax.org/books/chemistry/pages/8-4-molecular-orbital-theory)
10. [7.8: Sigma and Pi Bonds, Chemistry LibreTexts](https://chem.libretexts.org/Bookshelves/General_Chemistry/ChemPRIME_%28Moore_et_al.%29/07%3A_Further_Aspects_of_Covalent_Bonding/7.08%3A_Sigma_and_Pi_Bonds)
11. [Geometric Isomers, LON-CAPA, Michigan State University](https://s3.lite.msu.edu/res/msu/botonl/b_online/library/newton/Chy251_253/Lectures/Geometric_Isomers/GeometricIsomers.html)
12. [Pi Bonds, MCAT Organic Chemistry, Premier MCAT Prep](https://premiermcatprep.com/mcat-books/organic-chemistry/bonding/pi-bonds)
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