Single-Bond Rotation and the Alkene Double Bond

Rotation changes arrangement without changing connectivity

Hold one carbon of an ethane model still and rotate the other around the carbon–carbon axis. The same atoms remain bonded, but the positions of front hydrogens relative to rear hydrogens change continuously. These spatial arrangements are conformations, also called conformers or rotamers. In ethane, they interconvert rapidly at ordinary temperatures rather than existing as separately isolable substances.

A sigma bond has an electron distribution approximately symmetric about the internuclear axis. Rotation therefore does not destroy its head-on overlap. Rotation still has a small energy cost: interactions between adjacent bonds make some arrangements less favourable. In the simple bond-repulsion model used here, this barrier is described through torsional strain.

When front and rear carbon–hydrogen bonds line up in the viewing direction, ethane is eclipsed. Rotating by sixty degrees gives the staggered arrangement, in which the bonds are as far apart in angle as possible. Intermediate arrangements are skew conformations. Bond lengths and the approximately tetrahedral bond angles at each carbon remain essentially unchanged.

Two projections describe the same molecule

A Newman projection looks directly along the carbon–carbon bond. The nearer carbon is a central point and the rear carbon a circle. Three bonds from each appear 120 degrees apart in this projection. This does not mean the actual tetrahedral bond angle has become 120 degrees. A sawhorse projection instead shows the same bond obliquely, usually with the front carbon lower and rear carbon higher.

Staggering lowers torsional energy

In the simple bond-repulsion picture, eclipsing brings adjacent bonding clouds into less favourable alignments. Staggered ethane has less torsional strain and lower energy. The energy difference between the extreme arrangements is about 12.5 kilojoules per mole. Thermal motion and molecular collisions readily allow interconversion at ordinary temperatures, so rotation is nearly free for many practical purposes, though not literally without a barrier.

A pi bond changes the rotation problem

An acyclic hydrocarbon with one double bond has two fewer hydrogens than its alkane counterpart and follows ${\mathrm{C}}_n{\mathrm{H}}_{2n}$ for $n\ge 2$. Ethene is the first stable alkene. Each carbon uses three $s{p}^2$ hybrid orbitals for its sigma bonds in a plane. An unhybridised p orbital remains perpendicular to that plane and overlaps sideways with the p orbital on the other carbon.

The double bond contains one sigma and one pi bond. Pi density lies above and below the plane of the atoms. Twisting one carbon relative to the other disrupts the parallel p orbitals, so rotation costs much more energy than it does in ethane. This restriction allows certain alkenes to have distinct geometrical configurations.

The ethene carbon–carbon distance is about $134\mathrm{pm}$ and its C–H distance about $110\mathrm{pm}$. Around each carbon the geometry is approximately trigonal planar; measured H–C–H and H–C–C angles are about $116.6^{\circ }$ and $121.7^{\circ }$. The total double-bond enthalpy, about $681\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$, exceeds the ethane single-bond value of about $348\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$. The added pi component is nevertheless relatively exposed and can be consumed during addition.