Why do cyclic hydrocarbons have higher boiling points than their acyclic isomers?

As pointed out in the comments to this question, cyclic hydrocarbons have higher boiling points than their acyclic isomers. The major attractive force for hydrocarbons should be the London forces, which scale with surface area. My intuition has me thinking the cyclic molecules have less surface area, but they might enable more productive interactions over larger parts of their surface area due to their fixed conformations. I would love to see a study or some data that supports my intuition.

Boiling point data

Three carbons

• Cyclopropane - BP $=-33\ ^\circ\text{C}$
• Propene - BP $=-47.6\ ^\circ\text{C}$
• Propane - BP $=-42\ ^\circ\text{C}$

Four carbons

• Cyclobutane - BP $=12.5\ ^\circ\text{C}$
• Methylcyclopropane - BP $=4\ ^\circ\text{C}$
• 1-butene - BP $=-6.7\ ^\circ\text{C}$
• Butane - BP $=-1\ ^\circ\text{C}$

Five carbons

• Cyclopentane - BP $=49.2\ ^\circ\text{C}$
• 1-Pentene - BP $=30\ ^\circ\text{C}$
• Pentane - BP $=35.9\ ^\circ\text{C}$

Six carbons

• Cyclohexane - BP $=80.7\ ^\circ\text{C}$
• 1-Hexene - BP $=63\ ^\circ\text{C}$
• Hexane - BP $=68.5\ ^\circ\text{C}$