Intermolecular forces of attraction between positional isomers of alcohols This seemingly trivial question is as follows: Why is the boiling point of 1-butanol ($\ce {117.7 ^\circ C}$) higher than that of 2-butanol ($\ce {99 ^\circ C}$)?
The only reason I can think of is that the former molecule has a &quot;straighter&quot; chain than the other and a &quot;straighter&quot; chain has a greater surface area allowing for stronger van der Waals forces to exist between the molecules. As molecules in the liquid state do not pack themselves close to each other, ideas of &quot;packing efficiency&quot; would not be useful in explaining these boiling point differences. Perhaps, it is instructive for us to first observe if the trend is persistent for longer chains?
\begin{array}{|c|c|c|c|} \hline
\text {No. of C} &amp; \text{1-alcohol} &amp; \text{2-alcohol} &amp; \text{3-alcohol}\\ \hline
\ce{5} &amp;  138 &amp; 119.3 &amp; 116\\ \hline
\ce{6} &amp; 157 &amp; 140 &amp; 135\\ \hline
\ce{7} &amp; 175 &amp; 159 &amp; 156\\ \hline
\ce{8} &amp; 195 &amp; 178.5 &amp; 173\\ \hline
\ce{9} &amp; 214 &amp; 194 &amp; -\\ \hline
\end{array}
As can be seen from the above table, where the boiling points are all sourced from Wikipedia, with units of $\ce { _^\circ C}$, the trend is indeed persistent. The alcohol isomer with the $\ce {-OH}$ group attached to the 2-position consistently exhibits lower boiling points of around $\ce {16}$ to $\ce {20}$ $\ce {^\circ C}$. Shifting the hydroxyl group further along the chain by one more position also lowers the boiling point of the molecule but to a lesser extent.
The explanation I can come up with for this is the same as aforementioned: By placing the hydroxyl group at the end, there is the maximum surface area for interaction of the alcohol molecules with each other. Furthermore, the hydroxyl groups, being placed at the ends, would cause the least amount of disruption to the hydrophobic interactions between the alkyl chains. Does anyone have better explanations for the above trends?
It does seem that placing the hydroxyl group at the terminal carbon maximises the hydrogen bonding interactions between alcohol molecules and also the van der Waal forces between them. This is best illustrated with a crude diagram I have made:

Finer differences due to the position of the hydroxyl group along the chain, such as the differences between the 2- isomer and the 3- isomer, are not explained here. My explanation does explain the markedly higher boiling point of the 1- isomer.