The difference in boiling points of molecules with similar molecular masses, such as methane CH4 and ammonia NH3 , can be attributed to the differences in the types and strengths of intermolecular forces IMFs present in these substances.Methane CH4 is a nonpolar molecule, which means that the electrons are distributed evenly, and there is no net dipole moment. The primary intermolecular force present in methane is London dispersion forces, also known as van der Waals forces. These forces are relatively weak and arise due to temporary fluctuations in electron distribution, leading to instantaneous dipoles that attract other molecules.On the other hand, ammonia NH3 is a polar molecule, as the nitrogen atom has a higher electronegativity than the hydrogen atoms, resulting in a net dipole moment. In addition to London dispersion forces, ammonia also exhibits stronger intermolecular forces called hydrogen bonding. Hydrogen bonding occurs when a hydrogen atom is covalently bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine and is attracted to another electronegative atom in a neighboring molecule.The presence of hydrogen bonding in ammonia leads to stronger intermolecular forces compared to the London dispersion forces in methane. As a result, more energy is required to overcome these forces and convert ammonia from a liquid to a gaseous state. This explains the higher boiling point of ammonia 33.34C compared to methane 161.5C .In summary, the difference in boiling points of molecules with similar molecular masses, such as methane and ammonia, is due to the differences in the types and strengths of intermolecular forces present in these substances. The stronger hydrogen bonding in ammonia leads to a higher boiling point compared to the weaker London dispersion forces in methane.