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Boiling Points of LiquidsHELP TEST TOMORROW!

  1. Nov 3, 2005 #1
    Ok, I have been going through my chemistry notes on molecular geometry...and I don't understand the following:

    In my notes it says that
    -"as pressure goes up, so does the boiling point"
    -"High vapour pressure means weak attractive force between the molecules and low vapour pressure means strong attractive force".
    -"as temperature goes up, so does vapour pressure"

    So my question is:

    If when the pressure of a system increases, why does the boiling point? When pressure increases, the attractive force is weaker...so why does the temperature have to be greater in order for the substance to boil? If the bonds are weaker less energy is needed to break them, which to me indicates the boiling point should decrease.

    I know I'm wrong, could somebody please help me?

  2. jcsd
  3. Nov 3, 2005 #2
    I not very good a this stuff so dun take my words for it but I heard from my teacher say when there is a higher pressure acting on the water more energy is requiered to make it boil. So you can easily "boil water on top of mountains since its lower pressure up there. Also thats because why people use pressure pots or what ever they call it to cook on high places such as mountains so the food will actually be "Cooked"
  4. Nov 3, 2005 #3
    ooh...is it because when the vapour pressure is high, there is more kinetic energy meaning the bonds are weaker. But in order for this to occur, the temperature has to be high, because if the temperature was low then there wouldn't be much kinetic energy in the first place. If there isn't much kinetic engergy, then the bonds are stronger and the molecules don't move around as much, meaning that the vapour pressure is low!

    I think I figured it out. Is this correct?:smile:
  5. Nov 4, 2005 #4


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    under particular conditions a liquid starts to boil when it has acquired enough energy to expand, so that the vaporization can occur within the liquid (cavity formation through nucleation, instead of soley at the surface). With increased surrouding pressure, a higher temperature will be required.

    another effect of applied pressure to a liquid at equilibrium (boiling point) is increased vapor pressure of that liquid. You can use the clapeyron equation to determine the new boiling point at different pressures.
  6. Nov 4, 2005 #5


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    Rocketboy, your concepts need a little reworking. Stuff like this (in fact, most of physical chemistry) is taught pretty poorly (IMO) in most schools. Unfortunately, physical chem lies at the basis of concepts in many other areas as well. Now that I'm done with the sermon...

    The boiling point is defined as the temperature at which the vapor pressure of the liquid equals the ambient pressure. If you increase the ambient pressure, you must generate a greater vapor pressure in order to have "boiling". This takes more energy, and hence, a higher temperature.

    Additional note : The higher pressure has very little effect on the actual intermolecular forces in the liquid phase (and this is small enough that it can be neglected for a good approximation).

    Proof of last statement :

    If there was a substantial effect on the intermolecular forces, you would expect to see a significant change in density (which is clearly a function of the molecular mass and intermolecular separation) that reflects the change in pressure. However, there is no noticeable change, and it is for this reason that water is refered to as an incompressible fluid. Strictly speaking though, there is a very weak dependence that goes like :

    [tex]\frac {\rho(P)}{\rho_0} \approx c (P - P_0) [/tex]

    The value of the compressibility c, for water is about 5 * 10-10 Pa-1, which is really tiny !

    Increasing the ambient pressure to 100 atm increases the boiling point by over 60% (to about 600K), but the density increases by only about 0.1%.

    Hence, this change in the boiling point is not a result of the chages in intermoleclar forces.
  7. Nov 4, 2005 #6


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    you're confusing the vapor pressure with that of an applied external pressure (technically not mutually exclusive, don't need to consider this for now). A general correlation that they teach you in general chemistry is that molecules with weak intermolecular attractions have a greater tendency to vaporize, which means a greater vapor pressure and a lower boiling point (again, a generalization). You should note that vapor pressure consists of the vapor of the liquid constituent.

    You should apply a different perspective when they refer to an increase in an external pressure, for instance atmospheric pressure. I think that this is the source of your confusion.
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