Some Characteristic Features of π Bonds

The Geometry of the Double Bond

A carbon-carbon double bond consists of one sigma ($\sigma$) bond and one pi ($\pi$) bond. The sigma bond forms by the end-on overlap of orbitals along the internuclear axis. The pi bond, however, forms by the sideways overlap of two unhybridised p orbitals on adjacent carbon atoms. This sideways overlap has strict geometric requirements.

For the p orbitals to overlap effectively, they must be parallel to each other. This forces all four atoms directly attached to the double-bonded carbons to lie in the same plane. In ethene ($\text{H}_2\text{C}=\text{CH}_2$), the two carbon atoms and the four hydrogen atoms all lie in one flat plane. The two p orbitals that form the $\pi$ bond are perpendicular to this molecular plane and are parallel to each other.

Property I: Restricted Rotation

Because the $\pi$ bond depends on the parallel alignment of p orbitals, rotation about the carbon-carbon double bond is not free. If one $\text{CH}_2$ group rotates relative to the other, the p orbitals would no longer be parallel. Their overlap would decrease, and the $\pi$ bond would break.

The energy required to rotate about a double bond is roughly $250-270\ \text{kJ mol}^{-1}$, which is comparable to the bond enthalpy of a $\pi$ bond itself. This is far larger than the thermal energy available at room temperature ($\approx 2.5\ \text{kJ mol}^{-1}$ per molecule), so rotation is effectively frozen.

Property II: Location of the Electron Cloud

The sideways overlap of p orbitals produces a region of electron density that is not concentrated along the line joining the two nuclei. Instead, the $\pi$ electron cloud is located above and below the plane of the molecule. The sigma bond lies in the plane, while the $\pi$ bond occupies space on either side.

This spatial arrangement has a direct consequence for chemical reactivity. The $\pi$ electrons are not held tightly between the nuclei like sigma electrons. They are more exposed and more polarisable. An attacking reagent can approach the molecule from above or below the plane and interact with this exposed electron density.

Property III: Reactivity of the $\pi$ Bond

Because the $\pi$ electron cloud is located above and below the plane, the electrons are relatively loosely held. They are easily available to electrophilic reagents (electron-seeking species). In general, the $\pi$ bond provides the most reactive centre in any molecule that contains a multiple bond.

This explains why alkenes undergo addition reactions readily while alkanes do not. The $\pi$ bond acts as a source of electrons, attracting electrophiles. The sigma bond framework remains intact while the $\pi$ bond breaks, allowing new atoms or groups to attach to the carbon atoms. …