In chemistry, pi bonds (π bonds) are covalent chemical bonds where two lobes of one involved atomic orbital overlap two lobes of the other involved atomic orbital. These orbitals share a nodal plane which passes through both of the involved nuclei.
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. D orbitals also engage in pi bonding, and form part of the basis for metal-metal multiple bonding.
Pi bonds are usually weaker than sigma bonds. 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.
Pi bonds result from overlap of atomic orbitals that are in contact through two areas of overlap. Pi-bonds are 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 without breaking the pi bond, because rotation involves destroying the parallel orientation of the constituent p orbitals.
For homonuclear diatomic molecules, bonding π molecular orbitals have no nodal planes that pass between the bonded atoms. The corresponding antibonding, or π* ("pi-star") molecular orbital, is defined by the presence of an additional nodal plane between these two bonded atoms.
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A typical double bond consists of one sigma bond and one pi bond, for example the C=C double bond in ethylene. A typical triple bond, for example in acetylene, 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 bonds 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 lengths are 154 pm) in ethane, 134 pm in ethylene and 120 pm in acetylene. More bonds make the total bond shorter and stronger.
| ethane | ethylene | acetylene |
Pi bonds do not necessarily connect a pair of atoms that are also sigma-bonded.
In certain metal complexes, 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 sigma bond at all, only pi bonds. Examples include 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]
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