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Electron orbitals

Electron atomic and molecular orbitals, showing a Pi-bond at the bottom right of the picture.

In chemistry, pi bonds (π bonds) are covalent chemical bonds where two lobes of one involved electron orbital overlap two lobes of the other involved electron orbital. Only one of the orbital's nodal planes passes through both of the involved nuclei.


Two p-orbitals forming a π-bond.

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. P orbitals usually engage in this sort of bonding. However, d orbitals can engage in pi bonding also.

Pi bonds are usually weaker than sigma bonds because their (negatively charged) electron density is further from the positive charge of the atomic nucleus, which requires more energy. From the perspective of quantum mechanics, this bond weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation.

Although the pi bond by itself is weaker than a sigma bond, pi bonds are most often found in multiple bonds together with sigma bonds and the combination is stronger than either bond by itself. This can be seen from comparison of the carbon-carbon bond lengths in ethane (154 pm), ethylene (133 pm) and acetylene (120 pm).

Pi bond3

Top: two parallel p-orbitals. Bottom: pi bond formed by overlap. Pink and gray represent a ball and stick model of the molecular fragment that contains the pi bond.


Pi bond breaking when bond rotates because parallel orientation is lost. Pink and gray represent a ball and stick model of the molecular fragment that contains the pi bond.


Two s-orbitals continue to overlap when bond rotates because orientation is along axis. Circles represent s orbitals. Ellipses represent merged sigma bond. Pink and gray represent a ball and stick model of the molecular fragment that contains the sigma bond.

Two atoms connected through double bonds or triple bonds usually have one sigma bond between them and one or more pi bonds. Pi bonds result from parallel orbital overlap: the two combined orbitals meet lengthwise and create more diffuse bonds than the sigma bonds. Electrons in pi bonds are sometimes referred to as pi electrons. Molecular fragments joined by a pi bond cannot rotate about that bond unless the pi bond breaks because rotation involves destroying the parallel orientation of the constituent p orbitals.

Special casesEdit

Pi bonds do not necessarily have to connect a pair of atoms that are also sigma-bonded. In agostic complexes, pi interactions between a metal atom and the sigma bond of molecular hydrogen play a critical role in the reduction of some organometallic compounds. Alkyne and alkene pi bonds often bond "side-on" with metal atoms in a bond that has significant pi character.

In some cases of multiple bonds between two atoms, there is no sigma bond at all, only pi bonds. Examples are diiron hexacarbonyl (Fe2(CO)6), dicarbon (C2) and the borane B2H2. In these compounds the central bond consists only of pi bonding, and in order to achieve maximum orbital overlap the bond distances are much shorter than expected.[1]

See also Edit

References Edit

  1. Bond length and bond multiplicity: σ-bond prevents short π-bonds Eluvathingal D. Jemmis, Biswarup Pathak, R. Bruce King, Henry F. Schaefer III Chemical Communications, 2006, 2164 - 2166 Abstract

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