The boiling points of NH3, N2 and H2 are -33°C, -183°C and -252°C respectively. Note that boiling point is the same as condensation temperature. Their differences in boiling points are due to differences in intermolecular attractions or forces (IMFs). Diatomic nitrogen and diatomic hydrogen are both non-polar molecules because their bonds are non-polar. As such, they don’t experience dipole-dipole interactions. The only forces of attraction between molecules of N2 or H2 are London dispersion forces caused by temporary polarization of electron clouds. N2 has more electrons than H2 and is therefore more polarizable, accounting for the difference in their boiling points.
Ammonia, NH3, not only experiences London dispersion forces between molecules but also has hydrogen bonding. Hydrogen bonding is a bit of a misnomer as it is actually as strong IMF rather than a chemical bond. Hydrogen bonding occurs in molecules that have a hydrogen bonded to a fluorine, oxygen or nitrogen atom. Because these three elements are very electronegative they have polar bonds resulting in charge separation that creates a strong dipole moment. The slightly negative end of one molecule will be attracted to the slightly positive (hydrogen) end of another. This increase in intermolecular attraction results in ammonia having a higher boiling point the either N2 or H2.