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WikiHarold moved page Chemistry:Pi bond to Physics:Quantum pi bond without leaving a redirect

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imported>WikiHarold: WikiHarold moved page Chemistry:Pi bond to Physics:Quantum pi bond without leaving a redirect

2026-05-04T16:36:53Z

WikiHarold moved page Chemistry:Pi bond to Physics:Quantum pi bond without leaving a redirect

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{{Short description|Type of chemical bond}}
{{distinguish|Phi bond}}
[[Image:Ethylene 3D.png|200px|thumb|[[Ethylene]] (ethene), a small organic molecule containing a pi bond, shown in green.]]

In [[HandWiki:Chemistry|chemistry]], '''pi bonds''' ('''π bonds''') are [[Chemistry:Covalent bond|covalent]] chemical [[Chemistry:Chemical bond|bond]]s, in each of which two lobes of an [[Physics:Atomic orbital|orbital]] on one [[Atom|atom]] overlap with two lobes of an orbital on another atom, and in which this overlap occurs laterally. Each of these atomic orbitals has an [[Physics:Electron density|electron density]] of zero at a shared [[Node|nodal plane]] that passes through the two bonded [[Physics:Atomic nucleus|nuclei]]. This plane also is a nodal plane for the [[Chemistry:Molecular orbital|molecular orbital]] of the pi bond. Pi bonds can form in [[Chemistry:Double bond|double]] and [[Chemistry:Triple bond|triple bond]]s but do not form in [[Chemistry:Single bond|single bond]]s in most cases.

The Greek letter '''π''' in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. One common form of this sort of bonding involves p orbitals themselves, though d orbitals also engage in pi bonding. This latter mode forms part of the basis for [[Chemistry:Quintuple bond|metal-metal multiple bonding]].

== Properties ==
[[File:Pi-Bond.svg|thumb|Two p-orbitals forming a π-bond.]]

Pi bonds are usually weaker than [[Chemistry:Sigma bond|sigma bond]]s. The [[Chemistry:Carbon–carbon bond|C–C]] double bond, composed of one sigma and one pi bond,<ref>{{Cite book|title=Introduction to organic chemistry.|last1=Streitwieser|first1=Andrew|last2=Heathcock|first2=Clayton H.|last3=Kosower|first3=Edward M.|publisher=Macmillan|others=Heathcock, Clayton H., Kosower, Edward M.|year=1992|isbn=978-0024181701|edition=4th|location=New York|pages=[https://archive.org/details/introductiontoor00stre_0/page/250 250]|oclc=24501305|url-access=registration|url=https://archive.org/details/introductiontoor00stre_0/page/250}}</ref> has a [[Physics:Bond energy|bond energy]] less than twice that of a C–C single bond, indicating that the stability added by the pi bond is less than the stability of a sigma bond. From the perspective of [[Physics:Quantum mechanics|quantum mechanics]], this bond's weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation. This is contrasted by sigma bonds which form bonding orbitals directly between the nuclei of the bonding atoms, resulting in greater overlap and a strong sigma bond.

Pi bonds result from overlap of atomic orbitals that are in contact through two areas of overlap. Most orbital overlaps that do not include the s-orbital, or have different internuclear axes (for example px + py overlap, which does not apply to an s-orbital) are generally all pi bonds. Pi bonds are more diffuse bonds than the sigma bonds. [[Physics:Electron|Electron]]s in pi bonds are sometimes referred to as '''pi electrons'''. Molecular fragments joined by a pi bond cannot rotate about that bond without breaking the pi bond, because rotation involves destroying the parallel orientation of the constituent p orbitals.

For [[Chemistry:Homonuclear molecule|homonuclear]] [[Chemistry:Diatomic molecule|diatomic molecule]]s, bonding π molecular orbitals have only the one nodal plane passing through the bonded atoms, and no nodal planes between the bonded atoms. The corresponding ''anti''bonding, or π* ("pi-star") molecular orbital, is defined by the presence of an additional nodal plane between these two bonded atoms.

== Multiple bonds ==

A typical [[Chemistry:Double bond|double bond]] consists of one sigma bond and one pi bond; for example, the C=C double bond in [[Chemistry:Ethylene|ethylene]] (H2C=CH2). A typical [[Chemistry:Triple bond|triple bond]], for example in [[Chemistry:Acetylene|acetylene]] (HC≡CH), consists of one sigma bond and two pi bonds in two mutually perpendicular planes containing the bond axis. Two pi bonds are the maximum that can exist between a given pair of atoms. [[Chemistry:Quadruple bond|Quadruple bond]]s are extremely rare and can be formed only between [[Chemistry:Transition metal|transition metal]] atoms, and consist of one sigma bond, two pi bonds and one [[Chemistry:Delta bond|delta bond]].

A pi bond is weaker than a sigma bond, but the combination of pi and sigma bond is stronger than either bond by itself. The enhanced strength of a multiple bond versus a single (sigma bond) is indicated in many ways, but most obviously by a contraction in bond lengths. For example, in organic chemistry, carbon–carbon [[Bond length|bond length]]s are about 154 pm in [[Chemistry:Ethane|ethane]],<ref>{{cite journal |journal= Theoretica Chimica Acta |year= 1970 |volume= 18 |issue= 1 |pages= 21–33 |title= Relaxation during internal rotation ethane and hydrogen peroxyde |first1= A. |last1= Veillard |doi= 10.1007/BF00533694|s2cid= 94310101 }}</ref><ref>{{cite journal |title= The equilibrium carbon–carbon single-bond length in ethane |first1= Marlin D. |last1= Harmony |journal= J. Chem. Phys. |volume= 93 |issue= 10 |pages= 7522–7523 |year= 1990 |doi= 10.1063/1.459380 |bibcode= 1990JChPh..93.7522H}}</ref> 134 pm in ethylene and 120 pm in acetylene. More bonds make the total bond length shorter and the bond becomes stronger.

{| class="wikitable" style="margin:1em auto; text-align:center;"
|+Comparison of bond-lengths in simple structures
|-
||170px
||160px
||150px
|-
|| [[Chemistry:Ethane|ethane]] (1 σ bond)
|| [[Chemistry:Ethylene|ethylene]] (1 σ bond + 1 π bond)
|| [[Chemistry:Acetylene|acetylene]] (1 σ bond + 2 π bonds)
|}

==Special cases==
A pi bond can exist between two atoms that do not have a net sigma-bonding effect between them.

In certain metal complexes, pi interactions between a metal atom and [[Chemistry:Alkyne|alkyne]] and [[Chemistry:Alkene|alkene]] pi antibonding orbitals form pi-bonds.

In some cases of multiple bonds between two atoms, there is no net sigma-bonding at all, only pi bonds. Examples include diiron hexacarbonyl (Fe2(CO)6), dicarbon (C2), and [[Chemistry:Diborane(2)|diborane(2)]] (B2H2). In these compounds the central bond consists only of pi bonding because of a sigma antibond accompanying the sigma bond itself. These compounds have been used as computational models for analysis of pi bonding itself, revealing that in order to achieve maximum [[Orbital overlap|orbital overlap]] the bond distances are much shorter than expected.<ref>{{cite journal |last1=Jemmis |first1=E. D. |last2=Pathak |first2=Biswarup |last3=King |first3=R. Bruce |last4=Schaefer III |first4=Henry F. |year=2006 |title=Bond length and bond multiplicity: σ-bond prevents short π-bonds |journal=Chemical Communications |issue=20 |pages=2164–2166 |doi=10.1039/b602116f |pmid=16703142 |authorlink3=R. Bruce King |authorlink4=Henry F. Schaefer, III}}</ref>

== See also ==
* Aromatic interaction
* [[Chemistry:Delta bond|Delta bond]]
* [[Physics:Molecular geometry|Molecular geometry]]
* [[Chemistry:Pi backbonding|Pi backbonding]]
* [[Physics:Pi interaction|Pi interaction]]

== References ==
{{reflist}}

{{Chemical bonding theory}}

[[Category:Chemical bonding]]

{{Sourceattribution|Pi bond}}

Physics:Quantum pi bond - Revision history

imported>WikiHarold: WikiHarold moved page Chemistry:Pi bond to Physics:Quantum pi bond without leaving a redirect

imported>WikiHarold: WikiHarold moved page Chemistry:Pi bond to Physics:Quantum pi bond without leaving a redirect