# Why does water have an unusually low vapor pressure for a liquid?

Feb 25, 2016

Because of (you guessed it!) hydrogen bonding.

#### Explanation:

You are right to remark on the unusual volatility of water. If you look at water's $G r o u p$ $V I$ analogues, ${H}_{2} S$, ${H}_{2} S e$, and ${H}_{2} T e$, ALL of them are $\text{FOUL-SMELLING, POISONOUS GASES}$ at room temperature and pressure.

Water is, however, a fairly involatile liquid. Water has a great propensity to hydrogen-bond because of the polarity of the $O - H$ bond, and these hydrogen bonds, which act intermolecularly (i.e. BETWEEN molecules), account for the unusually high boiling point. The other Group VI elements are electronegative, but not so much as oxygen, and intermolecular H-bonding in water seems to be the determinant.

A similar situation exists for $H F$, whose boiling point I forget, but it is higher than those of $H C l$, $H B r$ etc. And the same situation again pertains for ammonia ($N {H}_{3}$) versus phosphine ($P {H}_{3}$) versus arsine ($A s {H}_{3}$). Can you account for the boiling point in this series? You'll have to find the boiling points first!

Note that I have used boiling points and vapour pressures interchangeably here, in that at room temperature and pressure when the vapour pressure of the liquid is $1$ $a t m$, the liquid is at its boiling point by definition.