View source for Pi bond

{{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 [[chemistry]], '''pi bonds''' ('''π bonds''') are [[covalent bond|covalent]] chemical [[chemical bond|bond]]s, in each of which two lobes of an [[atomic orbital|orbital]] on one [[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 [[electron density]] of zero at a shared [[node (physics)|nodal plane]] that passes through the two bonded [[atomic nucleus|nuclei]]. This plane also is a nodal plane for the [[molecular orbital]] of the pi bond. Pi bonds can form in [[double bond|double]] and [[triple bond]]s but do not form in [[single bond]]s in most cases.
The Greek letter '''π''' in their name refers to [[p orbital]]s, 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 orbital]]s also engage in pi bonding. This latter mode forms part of the basis for [[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 [[sigma bond]]s. The [[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 [[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 [[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 p<sub>x</sub> + p<sub>y</sub> overlap, which does not apply to an s-orbital) are generally all pi bonds. Pi bonds are more diffuse bonds than the sigma bonds. [[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 [[homonuclear molecule|homonuclear]] [[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 [[antibond|''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 [[double bond]] consists of one sigma bond and one pi bond; for example, the C=C double bond in [[ethylene]] (H<sub>2</sub>C=CH<sub>2</sub>). A typical [[triple bond]], for example in [[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. [[Quadruple bond]]s are extremely rare and can be formed only between [[transition metal]] atoms, and consist of one sigma bond, two pi bonds and one [[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]]s are about 154&nbsp;[[picometer|pm]] in [[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&nbsp;pm in ethylene and 120&nbsp;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
|-
||[[File:Ethane-staggered-CRC-MW-dimensions-2D.png|170px]]
||[[File:Ethylene-CRC-MW-dimensions-2D.png|160px]]
||[[File:Acetylene-CRC-IR-dimensions-2D.svg|150px]]
|-
|| [[ethane]] (1 σ bond)
|| [[ethylene]] (1 σ  bond + 1 π bond)
|| [[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 complex]]es, pi interactions between a metal atom and [[alkyne]] and [[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 (Fe<sub>2</sub>(CO)<sub>6</sub>), [[dicarbon]] (C<sub>2</sub>), and [[diborane(2)]] (B<sub>2</sub>H<sub>2</sub>). 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]] 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 |authorlink1=Eluvathingal Devassy Jemmis |authorlink3=R. Bruce King |authorlink4=Henry F. Schaefer, III}}</ref>
== See also ==
* [[Aromatic interaction]]
* [[Delta bond]]
* [[Molecular geometry]]
* [[Pi backbonding]]
* [[Pi interaction]]
== References ==
{{reflist}}
{{Chemical bonding theory}}
[[Category:Chemical bonding]]