In chemistry, pi bonds are covalent chemical bonds where two lobes of one
involved atomic orbital overlap two lobes of the other involved atomic
orbital. Each of these atomic orbitals is zero at a shared nodal plane, passing
through the two bonded nuclei. The same plane is also a nodal plane for the
molecular orbital of the pi 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 often 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; the C-C double bond has a bond
energy less than twice the C-C single bond bond energy; which leads to the
conclusion that the p orbital overlap to form molecular orbitals is a weaker bond
than when s orbitals overlap to form molecular orbitals. 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. 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 only
the one nodal plane passing through the bonded atoms, and no nodal planes
between the bonded atoms. The corresponding antibonding, or π*
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. 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 is indicated in many
ways, but most obviously by a contraction in bond lengths. For example
in organic chemistry, carbon–carbon bond lengths are about 154 pm in ethane, 134
pm in ethylene and 120 pm in acetylene. More bonds make the total bond shorter
and stronger. Special cases
A pi bond can exist between two atoms that do not have a net sigma-bonding
effect between them. 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 net sigma-bonding at all, only pi bonds. Examples include
diiron hexacarbonyl6), dicarbon, and diborane(2). 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 that in order to achieve
maximum orbital overlap the bond distances are much shorter than
expected. See also
Aromatic interaction Delta bond
Molecular geometry Pi backbonding
Pi interaction References