Carbon-carbon single bonds enable essentially free rotation about their axes. There is actually a small torsional barrier of about 12 kJ/mol (compared to the C-C bond energy of 350 kJ/mol), which slightly favors the staggered over the eclipsed conformation of the two methyl groups in ethane. To a good approximation, each carbon atom forms four
hybrid orbitals directed towards the vertices of a regular tetrahedron.
Each carbon atom involved in a C=C double bond forms three
hybrids, in the same plane, approximately 120° apart. The remaining
-orbitals on the two carbon atoms form a
-bond, which together with the
constitutes a C=C double bond. In contrast to a single bond, a double bond forms a rigid planar structure, with only a small amplitude of torsional motion allowed. If two of the hydrogen atoms in ethylene (ethene) are replaced by chlorine atoms, cis and trans isomers, molecules with distinct physical properties become possible. (These isomeric forms can be interconverted at higher temperatures or by UV radiation.)
A C≡C triple bond, which has two orthogonal
-bonds between the two carbon atoms, is, like a single bond, cylindrically symmetrical and allows free rotation. This is most evident in a molecule such as dimethylacetylene.
The central C-C bond in the biphenyl molecule represents an instance in which a single bond exhibits restricted rotation. In this case, the cause is steric hindrance between adjacent hydrogen atoms on the two rings. The most stable dihedral angle between the rings is about 39°. Rotation becomes even more restricted when larger groups are substituted for hydrogens.
The ball-and-stick models of the molecules illustrated in the graphics are intended to optimize display of rotational effects. Some distortions of atomic sizes have been introduced.