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In chemistry, pi bonds ( π bonds) are covalent bond chemical chemical bond, in each of which two lobes of an atomic 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 nodal plane that passes through the two bonded atomic nucleus. This plane also is a nodal plane for the molecular orbital of the pi bond. Pi bonds can form in double bond and but do not form in in most cases.
The Greek letter π in their name refers to , 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 also engage in pi bonding. This latter mode forms part of the basis for quintuple 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. 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 only the one nodal plane passing through the bonded atoms, and no nodal planes between the bonded atoms. The corresponding antibond, or π* ("pi-star") molecular orbital, is defined by the presence of an additional nodal plane between these two bonded atoms.
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 are about 154 picometer in ethane, 134 pm in ethylene and 120 pm in acetylene. More bonds make the total bond length shorter and the bond becomes stronger.
| +Comparison of bond-lengths in simple structures | ||
| ethane (1 σ bond) | ethylene (1 σ bond + 1 π bond) | acetylene (1 σ bond + 2 π bonds) |
In certain , 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 (Fe2(CO)6), dicarbon (C2), and 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 the bond distances are much shorter than expected.
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