Lowering Vapor Pressure
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?
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.
Here is a good site with a good explanation:
Todd Clark, Office of Science
US Department of Energy
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
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.
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).
Click here to return to the Chemistry Archives
Update: June 2012