Why do cyclic ethers have higher boiling points than their acyclic isomers? TL;DR version is the question title. Some context and data follow.

I was creating an assignment for my organic chemistry students in which they would need to draw as many isomers as they could from a simple formula, for example $\ce{C3H6O}$ (which fits 7 structures - 9 if you count minor enol tautomers and 11 if your count stereo). In this set (ignoring stereo and enols) there are:


  Two alcohols



Cyclopropanol - boiling point $=101-102\ ^\circ\text{C}$
Allyl alcohol - boiling point $=97\ ^\circ\text{C}$



  Two carbonyl compounds



Propanal - boiling point $=46\ ^\circ\text{C}$ 
Acetone - boiling point $=56\ ^\circ\text{C}$



  Three ethers



Propylene oxide - boiling point $=34\ ^\circ\text{C}$ 
Oxetane - boiling point
$=49-50\ ^\circ\text{C}$
Methyl vinyl ether - boiling point $=6\ ^\circ\text{C}$


The boiling point of methyl vinyl ether immediately stood out as anomalous. The fully saturated ethyl methyl ether has a boiling point of $7.4\ ^\circ\text{C}$.

Here are some more data comparing cyclic ethers to their acyclic isomers (and some analogs). The gap appears to be closing as the number of carbons increases. 


  Four carbons



Tetrahydrofuran - boiling point $=66\ ^\circ\text{C}$
1-Butene oxide - boiling point $=61-65\ ^\circ\text{C}$
Ethyl vinyl ether - boiling point $=33\ ^\circ\text{C}$
Allyl methyl ether - boiling point $=42-43\ ^\circ\text{C}$
Diethyl ether - boiling point $=34.6\ ^\circ\text{C}$
Methyl propyl ether - boiling point $=38.8\ ^\circ\text{C}$



  Five carbons



Tetrahydropyran - boiling point $=88\ ^\circ\text{C}$
2-Methyltetrahydrofuran - boiling point $=80.3\ ^\circ\text{C}$
Allyl ethyl ether - boiling point $=65-66\ ^\circ\text{C}$
Butyl methyl ether - boiling point $=70-71\ ^\circ\text{C}$
The data I could find suggests that cyclic ethers have higher dipole moments than their acyclic counterparts. 

$$\mathbf{Four~carbons}$$
\begin{array}{c @{} c} \hline
\text{THF} &amp; \mathrm{1.63~ D ^1, 1.75~ D ^2} \\
\text{1-butene oxide} &amp; \mathrm{1.89~ D^2} \\
\text{ethyl vinyl ether} &amp; \mathrm{1.26~ D ^2} \\
\text{diethyl ether} &amp; \mathrm{1.15 ~D ^3, 1.15 D ^2} \\
\text{methyl propyl ether} &amp; \mathrm{1.11 ~D ^2} \\\hline
\end{array}

$$\mathbf{Five~carbons}$$
\begin{array}{c @{} c} \hline
\text{Tetrahydropyran} &amp; \mathrm{1.58~D ^2, 1.87~D ^4} \\
\text{2-Methyltetrahydrofuran} &amp; \mathrm{1.38~ D ^5} \\
\text{dipropyl ether} &amp; \mathrm{1.21~D^2, 1.00~D^6} \\\hline
\end{array}

This seems consistent with intuitive expectations based on conformational models where, in the cyclic ethers where the molecular geometry is constrained by the ring, the lone pairs are pointing in the opposite direction from the carbon skeleton. Whereas the lone pairs and carbon chains in the conformationally more mobile acyclic isomers are less geometrically fixed and therefore less "directed" in space. 

It seems likely that the higher dipole moments in the cyclic compounds would lead to greater alignment / ordering in the liquid phase, which in turn would lead to higher boiling points.

References:


  1: http://en.wikipedia.org/wiki/Tetrahydrofuran
  2: https://physicalchemistryrosamonte.wordpress.com/material-balances/material-balances-on-a-crystallizer/physical-properties-of-pure-methanol/dipole-moment/ 
  3: http://en.wikipedia.org/wiki/Diethyl_ether
  4: http://www.drugfuture.com/chemdata/tetrahydropyran.html
  5: http://www.stenutz.eu/chem/solv28.php?s=1&amp;p=20
  6: https://books.google.com/books?id=G6jaBwAAQBAJ&amp;pg=PA50&amp;lpg=PA50&amp;dq=propyl%20ethyl%20ether%20dipole%20moment&amp;source=bl&amp;ots=VA52gqJ0kn&amp;sig=GM8tua_QpXccFW3VXpy9qiNf7rk&amp;hl=en&amp;sa=X&amp;ei=rzRvVZ__Cc32yQT9gIPACQ&amp;ved=0CDsQ6AEwBQ#v=onepage&amp;q=propyl%20ethyl%20ether%20dipole%20moment&amp;f=false