Why does rotating a double bond break the bond?
Consider a carbon-carbon ($\ce{C=C}$) bond.

A double bond consists of a $\sigma$ bond and a $\pi$ bond. A $\pi$ bond is formed by sideways overlapping of unhybridized p-orbitals of two carbon atoms, above and below the plane of carbon atoms.

If now one of the carbon atoms of the double bond is rotated with respect to the other, the p-orbitals will no longer overlap, and the $\pi$ bond should break. But the breaking of the $\pi$ bond requires 251 kJ/mol of energy, which is not provided by the collision of the molecules at room temperature. Consequently the rotation about a carbon-carbon double bond is not free, but it is strongly hindered or restricted.

In carbon-carbon single bond ($\ce{C-C}$) only 12.55 kJ/mol of energy is required (this data is for ethane molecule). At room temperature, the collision of molecules supply sufficient kinetic energy to overcome this energy barrier.