Pi bond

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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 top 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. From the perspective ofquantum 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.

Contents [hide] 1 Multiple bonds 2 Effect of bond rotation 3 Special cases 4 See also 5 References

[edit]Multiple bonds

Atoms connected via a double bond or triple bond have, in addition to one sigma bond, one or two pi bonds, respectively.

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 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 lengthsare in ethane (154 pm), in ethylene (133 pm) and in acetylene (120 pm). More bonds make the total bond shorter and stronger.

[edit]Effect of bond rotation

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.