Ask A Scientist Chemistry Archive

Question:
  Simply put, my question is why would the boiling point of a liquid
(take water, for example) be elevated when a nonvolatile solute is
added to it?  (Another way to put this is: "I don't understand boiling
point elevation" :)  )
 
  When I asked this question of my teacher, he explained that molecules
with enough kinetic energy to evaporate must be on the surface of the 
liquid to actually evaporate (become gaseous), and that the particles of 
the nonvolatile solute would take some of the room there is on the surface
of the liquid, and hence there would be less VOLATILE particles on the
surface (and so less can evaporate).  But what I do not understand is 
the first part of that explanation:  Why would molecules with enough
kinetic energy to become gaseous have to be on the surface to evaporate?
Are the spaces between molecules generally immense relative to the size
of the molecules?  E.g., in Rutherford's Scattering Experiment (if I
remember correctly), when alpha particles were shot into a thin sheet o
of gold, 99% of the alpha particles went straight through, so couldn't g
highly energetic gaseous molecules essentially make it through a liquid 
without a great deal of bouncing around (or am I just underestimating
the thinness of the gold sheet in the experiment), and evaporate without
being on the surface?
 jonathan ben-ami

Answer 1:
Check out the Chemistry Archives section, discussion #91.  It talks about
when salt is added to water why the freezing point goes down and the
boiling point goes up.  You may be looking for more than there actually
is.
-Joe Schultz

Answer 2:  Also remember that alpha particles being bare nuclei are 
MUCH smaller than molecules, so I don't think you can draw any conclusions 
about evaporation from the Rutherford scattering experiment
 jade hawk

Answer 3:
Also, alpha particles are quite a bit smaller than gold atoms
and they are fired through the gold foil coherently (i.e.,
all in the same direction) with a fairly high average
kinetic energy. Then, one has the surface tension of the
water-air interface to overcome (solvent molecules are
"organized" differently at surfaces than they are at the
interior of the solution)...but fundamentally, one has
to realize that the solven molecules are essentially
undergoing random thermal fluctuations (Brownian motion)
as they move around and collide with one another, and therefore
the odds of a molecule follwoing a straight-line trajectory
from the bottom of the beaker up to the surface are very tiny.
 
Man! We're getting some great questions!!
 
-topper
 
Answer 4:
   Your teacher is incorrect.  Vapor can form anywhere in a boiling liquid.
If you watch a boiling pot of water you will see bubbles form throughout
the water and on the bottom and sides of the pot. Also, if you pour a layer
of oil on top of the water it will still boil in the same way, in
particular at the same temperature, without having any exposed surface at
all.  The reason dissolved solutes increase boiling point is that the
solute must come out of solution in order for the water to boil.  This
costs entropy (the entropy of solution).  Boiling is entropically driven,
hence the reduction in the net entropy gain of boiling results in a higher
temperature needed for the reaction to go.  To put it without jargon: for a
little packet of water with dissolved salt to turn to steam the salt atoms
must, in the course of their random zooming about, ALL simultaneously leave
the packet.  This is not a likely event.  It becomes more likely as the
temperature (i.e. the average speed of zooming about) becomes higher,
though, and at a certain temperature above the ordinary boiling point it
becomes sufficiently likely to allow boiling in spite of the handicap.  You
can also see that the effect will naturally increase with the concentration
of dissolved solutes (i.e. the number of salt atoms per packet that
must simultaneously leave).  Incidentally, at the low energies involved here
molecules are quite hard and space-filling to each other.  You can regard
them as little hard balls, which are as closely packed as they can be in the
liquid state. That is why, for example, it is so very hard to compress or
expand a liquid --- there's no space.  This is the basis for hydraulic
equipment, e.g. the brakes in your car.
 christopher grayce