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Name: carol
Status: educator
Grade: 9-12
Location: CO
Country: N/A
Date: 1/24/2006


Question:
I know salt, sugar, etc. with lower the vapor pressure of a liquid, but I don't know exactly why. Since it is how many particles and not the nature of the particles, it is not about attraction. How does it interfere with the equilibrium?


Replies:
Dear Carol,

Basically, this happens because some of the solute molecules are at the surface of the solution, and they get in the way of solvent molecules which might otherwise be moving fast enough to escape the liquid and go into the vapor phase. This lowers the vapor pressure (and also increases the boiling point).

Technically, the nature of the particles does matter a bit. This can be taken into account using the solute's activity coefficient. However, this is a college-level topic.

best regards,
dr. topper


Carol,

Here is a good site with a good explanation:

http://neon.chem.uidaho.edu/~honors/collig.html



Regards,

Todd Clark, Office of Science
US Department of Energy


Carol,

The colligative property: vapor pressure lowering is best explained from the perspective of entropy. We can imagine that the driving force for a liquid to go into the vapor phase is controlled by the difference in micro-states or entropy of the liquid and gas phases -there is an entropy gain in going into the gas phase. However, adding a non-volatile solute will increase the entropy of the liquid phase without equivalently increasing the entropy of the gas phase (since the solute is non-volatile) - it can be seen that the number of micro-states for a solution is much higher than that of a pure solvent. This means that the entropy gain in the solvent particles of the solution going from liquid to gas phase is comparatively smaller than that in the pure solvent. Thus there is less of a driving force to go into the gas phase and therefore a lower number of particles in the gas phase and the consequent lower vapor pressure.

I have seen text and heard some instructors explaining this phenomenon as resulting from the solute particles physically blocking the escape of solvent particles. This is wrong - as is clearly evident when the vapor pressure of a pure solvent does not change when we place a porous cover over the solvent. I suspect this explanation is given as a simplified explanation - but I think this just introduces misconceptions and should be avoided.

Greg (Roberto Gregorius)


This was also my reply to a question regarding boiling point elevation: We will focus mostly on the water molecules and a little bit on the salt ions...

This phenomenon involves the law of entropy or the fact that the water molecules want to move toward a state of disorder, or higher entropy. Just imagine room temperature plain water in a pot, without solute, being in a very ordered state, one molecule aligned right next to the other, a very ordered liquid. At this time, there are also water molecules right above the surface that have escaped the water from the pot, creating a vapor pressure, moving from the more ordered liquid to a less ordered gas or vapor...this to reach a certain level of entropy, or disorder.

Now, in another pot of water, if you place some sodium and chloride ions (salt) into the ordered water, the water molecules must move away from each other to accommodate or dissolve the sodium and chloride ions. Now these water molecules with salt ions between them are in a more disordered state than they were without the salt ions. This means that less of them will have to escape into the air above the surface to reach that same level of entropy as with the plain water. It means that the water with the salt ions will have a lower vapor pressure than the plain water, but both systems will have the same level of entropy (disorder).

When you boil both liquids, you are putting heat (energy) into the system. The plain water will reach boiling (a certain high vapor pressure) at 100°C while the salt-water will reach boiling (the same exact high vapor pressure) at a higher temperature.

The state of the salt ions entropy also adds to this. In short, when the salt ions are dissolved in the water, they are in a disordered state that they like, adding their entropy to the system. Now when the water molecules evaporate and leave them behind, the salt ions must come back together in an ordered state of salt crystals, thus losing entropy. This loss of entropy also requires more energy (a higher temperature) to vaporize the water molecules and crystallize the salt molecules.

Hope this helps.

Joel Jadus


The "simple" answer is that the more concentrated a solution is the fewer molecules of the solvent there are at the surface of the solution so, as a result, the lower is the vapor pressure of the solvent. As a result, as you point out, it is the number of particles of solute (and not its chemical identity) that determines the decrease in vapor pressure of the solvent. In any case the solvent/solute/vapor are at equilibrium. The "real" answer is much more complicated and does depend upon the interactions of solute and solvent. The technical term for the simple case is "colligative properties" which is derived from the word "collective". This is meant to imply that it is the number of particles and not their chemical identity that is important. In the "real world" the interactions are much more complicated and does depend upon the details of the interactions between solvent and dissolved substance (solute).

Vince Calder



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