1. The problem statement, all variables and given/known data If the boiling point is the point at which vapor pressure > atmospheric pressure, so all of the water molecules can break free and fly into the atmosphere (i.e. overcoming the atmospheric pressure), then why is it that when atmospheric pressure > vapor pressure, the atmospheric pressure is pushing down on the water such that no vapor can escape too? I guess the question im asking is: Is vapor pressure independent of external pressure? 2. Relevant equations 3. The attempt at a solution So Vapor pressure is the result of the phenomenon when the surface molecules of a liquid gain enough kinetic energy to leave the liquid and become a gas. The force per unit area of these molecules is the vapor pressure. So it makes sense as T goes up, Vapor Pressure should increase since on average more molecules have the energy to break free from their intermolecular bonds. From an intermolecular bond perspective this makes sense. But form a pressure perspective im still a bit confused. If the boiling point is the point at which the vapor pressure > the atmospheric pressure so all the molecules can keep leaving the surface, then how is it that molecules can leave the surface via evaporation when pAtmosphere > pVapor Pressure? wouldnt the atmospheric pressure push all of the water molecules down into liquid? Just as all of the water molecules push against the atmospheric gas molecules to free themselves when pvapor pressure > patmosphere? ____________________________________ Also if im comparing the vapor pressure of a solid vs that of a liquid (sublimation vs evaporation), then if I decrease the temperature of both solid and liquid at the same rate BELOW the melting point of the solid (but at different pressures), apparently the vapor pressure of the liquid will decrease faster than that of the solid, to the extent where Vapor Pressure of the Solid > Vapor PRessure of the Liquid. What's going on here? I assume its a different case because we're in two differnet phases?