Hydrogen sulfide, H₂S, boils at -60 °C. Water, which has a molecule only half as heavy, boils at…

(a) Force in $\text{H}_2\text{S}$. Sulfur is not small and not very electronegative, so $\text{H}_2\text{S}$ forms only weak hydrogen bonds. Its main forces are London dispersion forces, with a small dipole-dipole contribution.

(b) The high boiling points. In $\text{H}_2\text{O}$ and $\text{HF}$, hydrogen is bonded to the very electronegative oxygen or fluorine. The hydrogen carries a strong positive charge. It is attracted to a lone pair on a neighbouring molecule. These hydrogen bonds are much stronger than dispersion forces, so more thermal energy is needed to separate the molecules. The extra attraction outweighs the lower mass.

(c) Water against hydrogen fluoride. Each $\text{H}_2\text{O}$ molecule has two hydrogen atoms and two lone pairs, so it can make up to four hydrogen bonds. Each $\text{HF}$ molecule has one hydrogen and three lone pairs, so it can make only about two, one as donor and one as acceptor, and it forms chains. Water holds a three-dimensional network, so water has the stronger hydrogen bonding per molecule and boils higher.