Final answer:
The vapor pressure of CH4 is greater than that of NH3 due to the presence of hydrogen bonding in NH3, which is a stronger intermolecular force than the dispersion forces that both CH4 and NH3 exhibit. This stronger force in ammonia means its molecules are more tightly held together, decreasing its tendency to vaporize compared to methane.
Step-by-step explanation:
To explain why the vapor pressure of CH4 (methane) is greater than that of NH3 (ammonia), one must consider the intermolecular forces present in these molecules. Methane only experiences dispersion forces (also known as London forces), which are comparatively weak and arise due to the momentary distribution of electron density within a molecule. In contrast, ammonia not only experiences dispersion forces but also hydrogen bonding, which is much stronger due to the electrostatic attraction between the hydrogen atom of one molecule and the nitrogen atom of another.
Since dispersion forces increase with molecular mass or size, and both CH4 and NH3 are of similar size, it is not the mass that affects their vapor pressure significantly but rather the types of intermolecular forces. The stronger hydrogen bonding in NH3 results in higher attractive forces between its molecules, making it harder for them to escape into the vapor phase. Thus, ammonia has a lower vapor pressure compared to methane, which with its weaker dispersion forces has molecules that can escape more readily into the vapor phase, resulting in a higher vapor pressure.