angle strain
A carbon atom with four single bonds 'wants' its bonds spread out at 109.5 degrees, the tetrahedral angle that keeps its four electron pairs as far apart as possible. Angle strain is the energy a molecule must store when its bonds are bent away from that preferred angle, squeezed too narrow or splayed too wide. It is the discomfort of forcing carbon into a shape its bonds dislike.
Where does this happen? Mostly in small rings, where geometry leaves no choice. Cyclopropane is a triangle: its internal angles are pinned near 60 degrees, far below 109.5, so its carbon-carbon bonds are badly bent — they can't point straight at each other and instead form weak, banana-shaped 'bent bonds'. This idea, that small rings are strained by their cramped angles, is the old Baeyer strain theory. Cyclobutane (about 90 degrees) is also strained, cyclopentane only mildly so.
Baeyer assumed all rings were flat, which led him to wrongly predict large rings would be strained too. The fix is that real rings pucker out of the plane: cyclohexane folds into the chair, where every angle is essentially the ideal 109.5 degrees and angle strain is gone. So angle strain dominates only the smallest rings; from cyclohexane onward, molecules dodge it by going three-dimensional. Where it does bite — in cyclopropane and epoxides — it makes those rings unusually reactive, eager to spring open and relieve the strain.
Cyclopropane's 60-degree internal angles force its bonds far from 109.5 degrees, loading it with angle strain that makes the three-membered ring eager to open in reactions.
The further the bond angle strays from 109.5 degrees, the more angle strain the ring stores.
Baeyer's theory assumed flat rings and wrongly predicted strain in large ones; in reality rings pucker. Cyclohexane and larger rings are essentially strain-free because they are non-planar.