Vapor Pressure and Boiling Point

In any liquid, molecules are constantly moving with a range of kinetic energies. At the surface, the fastest ones can break free of their neighbors and escape into the gas phase. If the liquid is in a closed container, some of those escaped molecules condense back, and eventually the rate of escape equals the rate of return. The pressure of vapor at that balance point is the vapor pressure. Weaker intermolecular forces (think dispersion forces in a nonpolar liquid versus hydrogen bonding in water) mean less energy is needed to escape, so more molecules vaporize at a given temperature, and vapor pressure is higher.

Boiling point is the temperature at which vapor pressure rises to meet the surrounding atmospheric pressure, letting bubbles of vapor form throughout the liquid, not just at the surface. Atop a tall mountain, atmospheric pressure is lower than at sea level, so the liquid's vapor pressure needs to climb only a little before it matches that reduced external pressure. Water reaches that match point below its normal 100 degrees C, which is why it can boil around 70 degrees C high in the mountains. The intermolecular forces inside the water have not changed; only the outside pressure it has to overcome has.

This is why high vapor pressure and high boiling point pull in opposite directions. A liquid with weak intermolecular forces evaporates easily (high vapor pressure at room temperature) but needs very little added heat to push that vapor pressure up to atmospheric pressure, so it boils at a low temperature. This connects directly to the gas laws and intermolecular force ranking elsewhere in the unit: whatever force you identify as strongest (hydrogen bonding, dipole-dipole, or dispersion) predicts both the lowest vapor pressure and the highest boiling point in a comparison.