Surprising values of intermolecular forces of ethane, oxygen and xenon I know that gases with stronger intermolecular forces should have a larger value of the van der Waals constant $a$. So I looked up a table on van der Waals constants and tried to reason about the relative values of that constant. Though I was able to explain most of the values, there was one thing I found absurd. Here:

$$
\begin{array}{cc}
\hline
\text{Substance} &amp; a~(\pu{atm {dm}^6 mol^{-2}}) \\
\hline
\text{Oxygen} &amp; 1.364 \\
\text{Xenon} &amp; 4.137 \\
\text{Ammonia} &amp; 4.225 \\
\text{Ethane} &amp; 5.507 \\
\hline
\end{array}
$$

My intution says that dioxygen, which has a slightly higher molar mass (32 vs. 30) than ethane, but significantly less molecular size or volume, should have a only slightly lower value of $a$ than that of ethane, but here we find a significant difference! Another example is xenon; its molar mass (131) is much larger than ethane (30), but its value of $a$ is smaller. Does it have to do with the dominance of molecular size over mass? Or atomicity or atomic dipoles? I suspect that it's probably the large and branched structure of ethane that's causing this.

Why does the molecular size/volume seem to play such a dominating role in the value of the constant $a$? I thought it was primarily influenced by molecular mass.
Van der Waals forces are weak and depend on polarizability. There is a correspondence between the boiling points of oxygen (−183 °C), xenon (−108 °C), and ethane (−88.5 °C) and their constants (1.364, 4.137, 5.507). These gases are polarizable to varying extents, but not polar with charge separation. The size of the molecule and its mass will affect polarizability - perhaps a way to look at this is to say electron density in the outer regions of the molecule will determine how easily dipole-dipole interactions can evolve.

But ammonia stands out with a boiling point (−33 °C) which is too high to be explained by van der Waals forces alone. Of all the gases listed, ammonia is the only one with significant hydrogen bonding capability. Perhaps if $\ce{H2O}$ were included, it would be an even greater outlier.