Colligative properties

I have some problems with these properties. 1) Lowering vapor pressure and boiling point. It happens because the solute molecules take space in the surface of the liquid so a lower number of the liquid molecules change to the gas form. Also the vapor pressure increases when you increase the temperature of the liquid because more molecules get to escape into a form of gas.

So the real question here, What makes boiling point related to vapor pressure? I have read a lot of articles about that.. The definition of boiling point is the temperature at which the vapor pressure of a liquid is equal to the pressure of the atmosphere on the liquid. Which will rapidly increase the rate of forming gas but why does that happen? It says because of the vapor pressure balanced the atmospheric pressure which makes bubbles form. How does it exactly balance out? Isnt pressure in all directions? How can you be sure that it balances out? Also, Bubbles are created at the bottom of the liquid, How does atmospheric pressure affect it?

2) Same applies to Freezing points.

I need a microscopic explanation. Also without some advanced chemistry complexity.

Here's an answer without complexity which I read in my book. When you heat a liquid in open vessel, its molecules turn to vapours, but the atmospheric pressure above is pushing these vapours down. Some vapours would have enough energy to tackle this atmospheric pressure and would get away from the liquid leading to evaporation. Only some particles at the surface would have enough energy but when vapour pressure reaches atmospheric pressure, all liquid molecules have enough energy to counteract the pressure on them and thus liquid can form vapours from the whole bulk of the liquid leading to boiling.

• Thanks! but you still didnt tell me how vapor pressure cancel out the force that the atmospheric pressure exerts on the liquid. Also another point is the bubbles, How does atmospheric pressure affect them? Like it is at the bottom of the liquid – Biker Mar 19 '16 at 9:45
• Pressure exerted by a gas is the force exerted by the the collisions of gas particles per unit area. So basically air particles above a liquid are constantly colliding with the vapour particles and pushing it down, kind of. Surface molecules have the highest energy and that's why they are the only liquid particles that could escape the pressure of air above them. Liquid particles at the bottom don't have enough energy to be converted to vapour. At boiling temperature all the liquid particles have enough energy, including those at bootom. As vapour is a gas, we see bubbles. – jatin Mar 19 '16 at 10:32
• So I can just assume that the vapor pressure is pushing upward and the atmospheric pressure is pushing downward. So they cancel out at the Boiling point. Now most of the molecules will have enough energy to escape the intermolecular force and there wont be a problem with atmospheric pressure, Right? Also another question if you don't, What happens if the vapor pressure is larger that atmospheric pressure like in cooking? My book says that the boiling point of water increases...? ِAnd why the temperature changes in solutions during boiling? – Biker Mar 19 '16 at 11:09
• I think you meant cooking under pressure. When we cook under pressure water molecules just cant vapourize at boiling point because now its facing much greater pressure which water molecules have to overcome to boil. So we have to heat water more and more till it has enough energy to overcome the increased pressure, giving us the utility to cook with water above 100 degrees Celsius. – jatin Mar 19 '16 at 11:22
• Yea that is what I thought. So as you said the pressure was created by vapor pressure then why it helps to boil the liquid instead of decreasing the number of the molecules that changes to gas. Boiling is a point where vapor and atmosphere pressure are equal. Shouldn't that just increase the pressure on the molecules on the surface preventing it to change to gas That was my initial question. – Biker Mar 19 '16 at 11:28