Pi bonds

Double and Triple (π) Bonds

We think of double and triple bonds being formed by edge on overlap of p-orbitals. Below are representations of the orbitals formed by the edge on overlap of the p-orbitals in ethylene and acetylene.

Figure 1: The π-bond in ethylene. The sticks connecting the atoms represent the bonds formed by the overlap of the sp² hybrid orbitals (the σ-bonds or σ-framework).

The carbon atoms in ethylene have three electron groups around them (the double bond to the other carbon plus the two single bonds to the hydrogens), which means they are formally sp² hybridized. This means that each carbon atom has a single p-orbital left sticking above and below the plane of the molecule. In figure 1 we see that when the positive lobes (indicated in red) of the two carbons overlap they form half of the electron density of the π-bond. The overlap of the negative lobes (indicated in blue) form the other half of the π-bond.

Figure 2: The π-bond in acetylene. The sticks connecting the atoms represent the bonds formed by the overlap of the sp hybrid orbitals (the σ-bonds or σ-framework). One π-bond (orbital) is colored green and yellow, the other is colored red and blue.

The carbon atoms in acetylene have two electron groups around them (the triple bond to the other carbon plus the single bond to the hydrogen), which means they are formally sp hybridized. This means that each carbon atom has two p-orbital left sticking out from the sides of the molecule. In figure 2 we see that when the positive lobes (indicated in green and red) of the two carbons overlap they form half of the electron density of each π-bond. The overlap of the negative lobes (indicated in yellow and blue) form the other half of each π-bond.

Based on template by A. Herráez as modified by J. Gutow

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