I'm currently studying relative lowering of vapor pressure with the addition of a solid non volatile solute.
The book explains it with the notion that some of the non volatile solute also appears at the surface between the liquid solution and the vapor. However, i don't understand why don't the molecules of the solute at the surface just escape? Obviously, a solid can't vaporize (Unless you raise the temperature, (And maybe reducing the pressure?)) however the molecules of the solute that are at the surface between the solution and the vapor and should not be held by the strong intermolecular bonds of the solid.
Hence, if i understand this correctly, they, just like the water (Or other solvent) molecules should be constantly moving around in the solution and sometimes 'bump' into other molecules to gain enough velocity to escape the solution. The few molecules that are are the surface between the solution and the air should be getting frequently hit with high energy molecules from below, which should constantly knock such molecules into the vapor. This would mean two things:
- The solute also has a small partial vapor pressure
- The vapor pressure should still lower, as once solute particles are knocked out into vapor, a few of the solvent particles should escape through the 'opening' before another solute can take it's place. Rinse and repeat, like a valve being opened and closed. Hence, while there would still be a lowering of vapor pressure, there shouldn't be any permanent layer being formed between the solution and the vapor.
But both of these are in contradiction to what's in the book. So where did i go wrong in my reasoning?