Isomers and boiling points

Isomers have the same molecular formula but different structures. Structural isomers join the atoms in a different order: C₄H₁₀ is both butane (a straight chain) and 2-methylpropane (branched). C₂H₆O is both ethanol, a liquid alcohol, and dimethyl ether, a gas. Same atoms, different substances.

A C=C double bond can't twist, so groups on it are stuck on one side or the other. When each carbon of the double bond has two different groups, you get cis (same side) and trans (opposite sides) isomers. Cis- and trans-but-2-ene have different boiling points, and in food, trans fats behave differently in the body from cis fats.

A carbon with four different groups attached is a chiral centre. Its molecule comes in two mirror-image forms that can't be laid on top of each other, like your left and right hands (enantiomers). They have the same boiling point and most properties, but the body tells them apart: one form of a drug can work while its mirror image does nothing or harms. That's why drug makers care which one they make.

Molecules attract each other through intermolecular forces, which are far weaker than bonds but decide boiling points. All molecules feel London dispersion forces, which grow with size: methane boils at −162 °C, pentane at 36 °C, octane at 126 °C. Polar molecules also attract end to end.

The strongest is hydrogen bonding, when H is attached to O, N or F and meets another O, N or F. Ethanol (46 g/mol) boils at 78 °C while propane (44 g/mol) boils at −42 °C: the same size, but ethanol's –OH groups hydrogen-bond. Branching lowers boiling points, because compact molecules touch less: 2,2-dimethylpropane boils at 10 °C, pentane at 36 °C.