Describe the differences between the enthalpy of vaporization of different molecules and the strength of their intermolecular forces.

The enthalpy of vaporization varies with the strength of intermolecular forces between molecules.

The enthalpy of vaporization is the amount of energy required to convert a liquid into a gas at a constant temperature. It is directly related to the strength of intermolecular forces between molecules. Molecules with strong intermolecular forces require more energy to overcome these forces and vaporize, resulting in a higher enthalpy of vaporization. For example, water has a high enthalpy of vaporization due to its strong hydrogen bonding between molecules.

On the other hand, molecules with weak intermolecular forces require less energy to vaporize, resulting in a lower enthalpy of vaporization. For example, methane has a low enthalpy of vaporization due to its weak London dispersion forces between molecules.

The strength of intermolecular forces is determined by the type of bonding between atoms in a molecule. Molecules with polar covalent bonds, such as water, have strong intermolecular forces due to the attraction between the partially positive and negative ends of the molecule. Molecules with nonpolar covalent bonds, such as methane, have weak intermolecular forces due to the lack of polarity in the molecule.

In summary, the enthalpy of vaporization is directly related to the strength of intermolecular forces between molecules. Molecules with strong intermolecular forces require more energy to vaporize, resulting in a higher enthalpy of vaporization, while molecules with weak intermolecular forces require less energy to vaporize, resulting in a lower enthalpy of vaporization.