Question

Why does water (or any solvent) boil when its vapour pressure equals the atmospheric pressure?

Answer 1

because the water (or solvent) molecules have sufficient energy to overcome the force (pressure) keeping them from "escaping" from the liquid phase.

Answer 2

thats true,,but my doubt goes like this,,, atmospheric pressure exerts force on the solvent downwards preventing them to escape,,,and moreover vapor pressure is defined as the pressure exerted by the vapors of the solvent ON THE SOLVENT itself,, so both forces are downwards , preventing the escape,,,,then how do they escape?

Answer 3

hm im not sure that the vapour is exerting force downwards on the solvent. Yes, solvent molecules are in an equilibrium between the two phases (gas and liquid). when this equilibrium is driven forward by the evaporation (or escape) of more liquid molecules into the gas phase as their kinetic energy increases there is a net movement of the solvent.

Answer 4

but why EXACTLY at the point when vapour pressure equals atmospheric pressure?

Answer 5

i think it's because it breaks the equilibrium at that point. There is a net movement of solvent molecules into the gas phase. I mean, it's kind of odd. I don't know if you've ever left a glass of water sitting in a room (at mild temperature) for a long time. The water evaporates. Does that mean the pressure of the room decreased, or that the vapour pressure of the water increased for no reason?

nope. There is always a gaussian distribution of kinetic energy in a group of molecules, some always have sufficient energy to overcome the atmospheric pressure, but they may not be at the surface of the container.

Answer 6

I think the statement "water (or any solvent) boils when its vapour pressure equals the atmospheric pressure" it's referring to a process that is short, temporally. (i could be wrong about this last point, but it's something to think about).