alcohol hydrogen bonding
Why does ethanol (boiling point 78 °C) boil so much hotter than propane (boiling point -42 °C), even though both are small molecules of similar weight? The answer is the -OH group and its ability to form hydrogen bonds. These extra sticky forces between alcohol molecules are the single biggest reason alcohols look and behave the way they do.
A hydrogen bond is an attraction between an O-H hydrogen on one molecule — left slightly positive because oxygen hogs the bonding electrons — and a lone pair on the oxygen of a neighbouring molecule. Picture a crowd of alcohol molecules holding hands through their -OH groups, each oxygen offering lone pairs and each O-H reaching for a neighbour's oxygen. To boil the liquid you must pull these hands apart, which costs extra energy and raises the boiling point well above that of a comparable hydrocarbon or ether.
The same hydrogen bonding explains other properties. Small alcohols (methanol, ethanol, propanol) mix with water in all proportions because their -OH groups hydrogen-bond to water just as readily as to each other. As the greasy carbon tail grows, though, water solubility drops — by about eight carbons an alcohol behaves more like an oil. Hydrogen bonding is thus the quiet force behind alcohols' high boiling points, water miscibility, and even biological roles like the structure of DNA and proteins.
Ethanol (46 g/mol) boils at 78 °C, while dimethyl ether (the same formula, C2H6O, but no O-H) boils at -24 °C. The 100-degree gap is entirely hydrogen bonding.
Same atoms, but only the alcohol's O-H can hydrogen-bond, so it boils 100° higher.
Ethers and alkyl halides have polar bonds too, but with no O-H (or N-H/F-H) hydrogen they cannot donate a hydrogen bond, only accept one — which is why they boil far lower than their isomeric alcohols.