A double bond between 2 carbon atoms means there is a $\pi$ bond with the side orbital overlaps in 1 plane, in parallel to the axial $\sigma$ bond.
The consequence is, free rotation would need to break this $\pi$ bond, making the double bond just a single bond, what would require a strong torque and a lot of energy.
By other words, particular rotation positions around a single bond have ( almost ) the same energies. But for double bonds, the molecule energy strongly depends on on this orientation. The $\pi$ bond can exist, only if the bystander groups on both carbon atoms are about in the same plane. If they wete on about perpendicular planes, the orbitals of both atoms cannot overlap to form the bond.