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.
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 are also assumed to engage in pi bonding but this is not necessarily the case in reality, although the concept of bonding d orbitals still accounts well for hypervalence.
Pi bonds are usually weaker than sigma bonds because their (negatively charged) electron density is farther from the positive charge of the atomic nucleus, which requires more energy. 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.
Although the pi bond by itself is weaker than a sigma bond, pi bonds are often components of multiple bonds, together with sigma bonds. The combination of pi and sigma bond is stronger than either bond by itself. The enhanced strength of a multiple bond vs. 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 ethane (154 pm), ethylene (133 pm) and acetylene (120 pm).
In addition to one sigma bond, a pair of atoms connected via double bond and triple bonds have one or two pi bonds, respectively. Pi bonds result from overlap of atomic orbitals that with 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.
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.