CH 106 - Lesson 2
Home Up Hybridization Covalent Bonding Sigma Bonds Types of Hybridization Double Bonds Pi Bonds Triple Bonds Comparing Hybrid Types Valence


Pi Bonds

Pi bonds involve the electrons in the leftover p orbital for each carbon atom. Those p orbitals are the electron clouds or orbitals that are shown going up above and below each carbon atom.

In the sketch at the top here, the sp2 hybrid orbitals of one atom are shown as sticks so that you can concentrate on the unhybridized, or leftover, p orbital. The bottom sketch shows you how the hybrid orbitals of two carbon atoms can come together, overlap, and form a regular sigma bond (shown as a stick). Notice that this also brings the leftover p orbitals so close that they can also overlap and form what we call a pi bond. Notice that the overlapping occurs in two places, above and below the sigma bond. The pi bond does not overlap in the region directly between the two carbon atoms where the sigma bond is formed.

Diagram of p orbitals forming a pi bond. [62orb07.JPG]


The distinction between a sigma bond and a pi bond is diagrammed here. The sigma bond has orbital overlap directly between the two nuclei. The pi bond has orbital overlap off to the sides of the line joining the two nuclei.

Diagram comparing sigma and pi bonds. [62orb08.JPG]

The combination of a sigma and a pi bond between the same two carbon atoms is a double bond. A double bond consists of a sigma bond (using hybrid orbitals) and a pi bond (using p orbitals).

Another way of showing how two carbon atoms can form a double bond is indicated in the bottom (part c) of this diagram. Here we have two carbon atoms with sp2 hybridization, and each of those carbon atoms is bonded to two hydrogen atoms and also to the other carbon atom. Notice that this diagram shows how the carbon atoms and the hydrogen atoms are all in a flat plane. This is drawn in perspective. Notice that there is a sigma bond between the carbon atoms and each of the hydrogen atoms. Notice also a sigma bond between the two carbon atoms. The sigma bonds (or sigma orbitals) are shown as the dark shaded areas in this drawing. Each one of those sigma bonds uses one of the hybrid orbitals. Remember that with sp2 hybridization, there are three hybrid orbitals, and those are the ones used to form the sigma bonds between all the atoms.

Remember also that we have a leftover p orbital for each carbon atom. Those p orbitals are the electron clouds or orbitals that are shown going up above and below each carbon atom. In this particular diagram, the shading between those p orbitals shows that the p orbitals overlap one another and allow the electrons in those p orbitals to be shared. This kind of bond is called a pi bond. The pi bond results when p orbitals overlap one another in this side-to-side fashion.


This diagram tries to represent all of the sigma and pi bonds in ethene (C2H4). This is perhaps the best way to represent that a double bond consists of a sigma bond and a pi bond. It may possibly look like three bonds to you, but it is not. There is a sigma bond in the center and a pi bond above and below that. Altogether that represents a double bond.

Diagram of bonding in ethene. [62orb09.JPG]


Let me point out some important structural consequences of a double bond.

The double bonded carbon atoms and the atoms bonded to them all lie in a flat plane.
The pi bond sticks out above and below that plane.
There is no rotation around a double bond. You could not twist and turn those two carbon atoms without breaking the pi bond. The pi part of a double bond does not allow for rotation.


Model kits are generally inadequate for showing this sigma-pi nature of a bouble bond. In the model kits available to you, the double bonds are represented using either springs or curved pieces of plastic.

These models might lead you to think that double bonds are formed by some kind of process, in which normal single bonds are bent around to form a curved bond, but that is just not true at all. It is just the simplest way to indicate a double bond using models like this. So keep that in mind when you use models that include double bonds.

The models are not all bad, however. When you make a model of this molecule as part of your lab work, you will note that it does show that all six atoms lie in a flat plane and that you cannot twist or turn or rotate this model like you could the alkanes that you made earlier, which had only single bonds.

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E-mail instructor: Eden Francis

Clackamas Community College
2001, 2003 Clackamas Community College, Hal Bender