# Double bond
A double bond is a covalent bond in which two pairs of electrons are shared between two atoms, giving a bond order of 2[1](https://openstax.org/books/chemistry/pages/7-key-terms) • [2](https://www.chem.ucla.edu/~harding/IGOC/D/double_bond.html). In the usual valence-bond picture it consists of one sigma (σ) bond, whose electron density lies along the internuclear axis, and one pi (π) bond, formed by side-by-side overlap of p orbitals above and below that axis[3](https://openstax.org/books/chemistry-atoms-first-2e/pages/5-1-valence-bond-theory). Double bonds occur between carbon atoms in alkenes, between carbon and oxygen in carbonyl groups, and between many other element pairs, including the O=O bond of molecular oxygen[3](https://openstax.org/books/chemistry-atoms-first-2e/pages/5-1-valence-bond-theory).
| Key fact | Detail |
|---|---|
| Definition | Covalent bond in which two electron pairs are shared; bond order = 2, four electrons total[1](https://openstax.org/books/chemistry/pages/7-key-terms) • [2](https://www.chem.ucla.edu/~harding/IGOC/D/double_bond.html) |
| Composition | One σ bond plus one π bond; the first bond between two atoms is always σ, and only one σ bond is possible in a given location[3](https://openstax.org/books/chemistry-atoms-first-2e/pages/5-1-valence-bond-theory) |
| π-bond geometry | Side-by-side overlap of two p orbitals; a nodal plane includes the internuclear axis, unlike a σ bond, which has no such nodal plane[3](https://openstax.org/books/chemistry-atoms-first-2e/pages/5-1-valence-bond-theory) • [4](https://old.goldbook.iupac.org/html/S/S05434.html) |
| Rotation barrier | About 350 kJ/mol (84 kcal/mol) for a C=C double bond, versus 12 kJ/mol for rotation in ethane[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes) |
| Bond length and strength | Ethylene C=C: 134 pm, 728 kJ/mol; ethane C–C: 153 pm, 377 kJ/mol[6](https://openstax.org/books/organic-chemistry/pages/1-8-sp2-hybrid-orbitals-and-the-structure-of-ethylene) |
| Hybridization | Each carbon of an alkene double bond is sp² hybridized; the π bond uses the remaining unhybridized 2p orbital on each carbon[7](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds) |
| Stereochemical consequence | Restricted rotation makes cis–trans (E/Z) isomerism possible when both double-bond carbons carry two different groups[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes) |
## Structure of the π bond
The σ component of a double bond forms like any single bond, from end-to-end orbital overlap along the internuclear axis. The π component is different: it arises from the side-by-side overlap of two parallel p orbitals, so the regions of electron density lie on opposite sides of the axis, and along the axis itself there is a node, a plane where the probability of finding an electron is zero[3](https://openstax.org/books/chemistry-atoms-first-2e/pages/5-1-valence-bond-theory). IUPAC defines the two labels as symmetry designations: a two-center π-bond has a nodal plane that includes the internuclear bond axis, whereas a σ-bond has no such nodal plane[4](https://old.goldbook.iupac.org/html/S/S05434.html).
In ethene (C₂H₄), each carbon uses three sp² hybrid orbitals, formed from the 2s orbital and two of the 2p orbitals, to make the two C–H σ bonds and the C–C σ bond. The third, unhybridized 2p orbital on each carbon, perpendicular to the plane of the hybrid orbitals, overlaps sideways with its partner to form the π bond[7](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds). All six atoms of ethene lie in one plane, because tilting the two sp² frameworks would misalign the p orbitals and weaken the π overlap; the planar arrangement is the most stable bonding configuration[7](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds).
## Strength, length, and rotation
A double bond is both shorter and stronger than the corresponding single bond between the same atoms, because four electrons bond the nuclei rather than two. Ethylene has a C=C length of 134 pm and a bond strength of 728 kJ/mol, against 153 pm and 377 kJ/mol for the C–C single bond of ethane[6](https://openstax.org/books/organic-chemistry/pages/1-8-sp2-hybrid-orbitals-and-the-structure-of-ethylene). The π bond itself adds 267 kJ/mol of strength to a carbon–carbon double bond over the single bond, and adding a second π bond (as in a triple bond) adds a further 225 kJ/mol; each individual π bond is generally weaker than the σ bond between the same two atoms because side-by-side overlap is less effective than end-to-end overlap[8](https://openstax.org/books/chemistry-2e/pages/8-1-valence-bond-theory).
**Restricted rotation.** Rotation around a single σ bond is essentially free, because end-to-end overlap does not depend on the relative orientation of the orbitals on the two atoms. Rotation around a double bond is different: turning the two halves of the molecule destroys the parallel alignment of the p orbitals and breaks the π bond, so the barrier to rotation must be at least as great as the π-bond strength, estimated at 350 kJ/mol (84 kcal/mol)[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes) • [7](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds). For comparison, the rotation barrier in ethane is only 12 kJ/mol[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes).
## Cis–trans isomerism
Because the two halves of a double bond cannot rotate relative to each other at ordinary conditions, the substituents on a double bond hold fixed spatial relationships. In 2-butene, the two methyl groups can sit on the same side of the double bond (cis-2-butene) or on opposite sides (trans-2-butene); since interconversion would require breaking the π bond, the two forms are distinct, isolable compounds called cis–trans stereoisomers[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes).
[Cis–trans isomerism](https://www.edgepedia.org/cis-trans-isomerism) occurs whenever both double-bond carbons are attached to two different groups; if either carbon bears two identical groups, only one configuration exists[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes). When the cis/trans labels become ambiguous, the E/Z system is used instead: the configuration is E if the two groups of highest Cahn–Ingold–Prelog priority (ranking of substituents by atomic number to label E/Z) lie on opposite sides of the double bond[9](https://www.chem.ucla.edu/~harding/IGOC/E/e.html). The geometry of a double bond can matter biologically as well as chemically; the sex attractant of the common housefly is the alkene cis-9-tricosene[5](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes).
## Double bonds beyond carbon–carbon
The carbon–oxygen double bond of a carbonyl group (C=O) resembles the alkene C=C bond in many respects. The carbonyl carbon is sp² hybridized and forms three σ bonds; its fourth valence electron occupies a p orbital that overlaps with an oxygen p orbital to form the π bond, while the oxygen carries two nonbonding electron pairs[10](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/18%3A_Ethers_and_Epoxides_Thiols_and_Sulfides/18.14%3A_Preview_of_Carbonyl_Chemistry). Representative values place the C=O bond at 119.7 pm and 745 kJ/mol, compared with 140.1 pm and 358 kJ/mol for a C–O single bond[11](https://openstax.org/books/chemistry/pages/8-1-valence-bond-theory).
**Delocalized double bonds.** In molecules with alternating single and double bonds, the π electrons need not stay localized between one pair of atoms. In benzene, all six carbons are sp² hybridized with bond angles of 120°, and each p orbital overlaps equally with both neighbors, so the six π electrons are delocalized over the entire ring; this delocalization stabilizes benzene by 150 kJ/mol relative to a localized arrangement[12](https://openstax.org/books/organic-chemistry/pages/15-2-structure-and-stability-of-benzene). IUPAC classifies double bonds that are neither conjugated nor cumulative as isolated double bonds[13](https://goldbook.iupac.org/terms/view/I03287).
## Reactivity
The π bond is the reactive part of a double bond in many reactions. In additions to unsaturated hydrocarbons, the only bond broken in the hydrocarbon is the π bond; its electrons form a bond to one atom of an added reagent such as bromine, while the reagent's own bond electrons form the second new bond to the other carbon[14](https://openstax.org/books/chemistry/pages/chapter-20). Nearby σ bonds can also interact with a π system: in propene, group orbitals of the methyl C–H bonds interact with the π orbital of the double bond, an effect known as hyperconjugation[4](https://old.goldbook.iupac.org/html/S/S05434.html).
## References
1. [Ch. 7 Key Terms, Chemistry, OpenStax](https://openstax.org/books/chemistry/pages/7-key-terms)
2. [Double bond, Illustrated Glossary of Organic Chemistry, UCLA](https://www.chem.ucla.edu/~harding/IGOC/D/double_bond.html)
3. [5.1 Valence Bond Theory, Chemistry: Atoms First 2e, OpenStax](https://openstax.org/books/chemistry-atoms-first-2e/pages/5-1-valence-bond-theory)
4. [σ, π (sigma, pi), IUPAC Gold Book, entry S05434](https://old.goldbook.iupac.org/html/S/S05434.html)
5. [7.4 Cis–Trans Isomerism in Alkenes, Organic Chemistry (OpenStax), LibreTexts](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/07%3A_Alkenes-_Structure_and_Reactivity/7.04%3A_Cis-Trans_Isomerism_in_Alkenes)
6. [1.8 sp² Hybrid Orbitals and the Structure of Ethylene, Organic Chemistry, OpenStax](https://openstax.org/books/organic-chemistry/pages/1-8-sp2-hybrid-orbitals-and-the-structure-of-ethylene)
7. [5.3 Multiple Bonds, Chemistry: Atoms First, OpenStax](https://openstax.org/books/chemistry-atoms-first/pages/5-3-multiple-bonds)
8. [8.1 Valence Bond Theory, Chemistry 2e, OpenStax](https://openstax.org/books/chemistry-2e/pages/8-1-valence-bond-theory)
9. [E, Illustrated Glossary of Organic Chemistry, UCLA](https://www.chem.ucla.edu/~harding/IGOC/E/e.html)
10. [18.14 Preview of Carbonyl Chemistry, Organic Chemistry (OpenStax), LibreTexts](https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_%28OpenStax%29/18%3A_Ethers_and_Epoxides_Thiols_and_Sulfides/18.14%3A_Preview_of_Carbonyl_Chemistry)
11. [8.1 Valence Bond Theory, Chemistry, OpenStax](https://openstax.org/books/chemistry/pages/8-1-valence-bond-theory)
12. [15.2 Structure and Stability of Benzene, Organic Chemistry, OpenStax](https://openstax.org/books/organic-chemistry/pages/15-2-structure-and-stability-of-benzene)
13. [Isolated double bonds, IUPAC Compendium of Chemical Terminology, entry I03287](https://goldbook.iupac.org/terms/view/I03287)
14. [Answer Key Chapter 20, Chemistry, OpenStax](https://openstax.org/books/chemistry/pages/chapter-20)
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*Topic: Encyclopedia › Physical world and mathematics › Chemistry › Chemical principles and methods › Chemical bonding and intermolecular forces*
*Initially written Sep 17, 2026 · Reviewed: Oct 15, 2026 · Edited: Oct 15, 2026 · Last review: Oct 15, 2026*
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