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Effect of air pressure on boiling point of water

  1. Sep 7, 2012 #1
    I know that the boiling point of water is lower when the atmospheric pressure is less (such as at the top of a mountain), but why is that?

    Once a water molecule has the energy to escape from the binding forces that keep it with the other water moelcules in the liquid, I don't understand why the movement of the surrounding air molecules (i.e pressure) alters the chances of escape of that molecule into the surrounding air?
  2. jcsd
  3. Sep 7, 2012 #2
    The pressure acts as an external force on the water. This inhibits the ability of water molecules to break apart as they are constantly being pushed together by the surrounding air pressure.
  4. Sep 7, 2012 #3
    That in itself cannot be the explanation - for the gaps between the molecules in the air are so much larger than the gaps between the molecules in the water.

    If the water molecule has the energy to escape from its neighbouring water molecules, and does so, why does it not then just get jumbled up with the surrounding air molecules - and stray away from the water?
  5. Sep 7, 2012 #4


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    You need to keep in mind that evaporation happens at any temperature. Not just at boiling point. There will always be particles with enough energy to overcome the binding forces and depart the surface. (Ionic fluids notwithstanding.) However, if there are any particles of water in the air, there is also a chance that a particle will "stick" to the surface. So at any given temperature you have simultaneous processes of evaporation and condensation. The rate of evaporation depends on temperature. Rate of condensation depends on concentration of water particles in the air. Number of particles per volume at given temperature gives you pressure. Partial pressure, to be specific, since you are only interested in concentration of water molecules. Partial pressure at which evaporation and condensation are balanced out is the vapor pressure of the liquid at that temperature.

    Now, picture a small bubble of water vapor inside the fluid. If the pressure in that bubble is above vapor pressure, the vapor in the bubble will condense. If the pressure in that bubble is bellow vapor pressure, more fluid will evaporate into the bubble. The pressure in the bubble, however, must be balanced out by atmospheric pressure. So if the vapor pressure of the liquid at given temperature is bellow atmospheric pressure, any vapor bubble inside the fluid will collapse. If the vapor pressure exceeds atmospheric, vapor bubbles in the fluid will increase in size. It's that growth of vapor bubbles that we call boiling. Since vapor pressure of the fluid varies with temperature, boiling temperature naturally depends on atmospheric pressure.
  6. Sep 7, 2012 #5
    Thanks K^2. Yes the observation is that boiling involves bubbles of the vapourised water - but that is probably a consequence of heating the water from below.

    Take the case of water at say 75 centigrade in a pan - and rather than heat it, reduce the air pressure. Once the air pressure gets to a low enough pressure to allow the water to 'boil', what will happen? Will bubbles appear?

    I am still wanting to understand what happens at the molecular level? A bubble of water vapour will have billions of molecules in it. I don't understand how pressure can affect what the water molecules are doing at the interface with the air molecules?
  7. Sep 7, 2012 #6

    Boiling is not about what "water molecules are doing at the interface", as is the vaporization below boiling point. Boiling "happens" in the volume of the liquid.
    K^2 have explained it very well.
  8. Sep 10, 2012 #7
    Thanks, but I still don't understand how the pressure of the air alters the boiling point?

    So keeping with the example of the water in a pan at 75C, and the air pressure is reduced, at a certain reduction in the pressure, bubbles of water vapour will form in the body of the water and rise to the surface, i.e. the water boils.

    Before this low pressure was reached, bubbles of water vapour did not form in the body of the water because the pressure on the surface of the liquid by the air molecules was doing what?

    The air molecules have such large gaps between themselves, as compared to the distance between the water molecules, that their pressure on the water must be like trying to hold back a wave on the beach by throwing pebbles at it!

    Yes, some of the water molecules will be pushed back into the surface (and thus create a pressure on the water vapour bubble inside the body of the water), but most of the water molecules will have large moments when they are not being hit by the air molecules, why don't these just leave the body of the water - as they in effect have no pressure on them?

    I think the explanation will involve understanding how the water molecules at the surface 'hold' on to each other - but that on its own still seems to leave something missing from the explanation?

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  9. Sep 10, 2012 #8


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    They do. Water does evaporate off the surface without boiling. Frozen ice at 1 atmosphere of pressure evaporates into dry air. Boiling and evaporation are two separate processes.
  10. Sep 10, 2012 #9
    So at the point of boiling, when the pressure has dropped by just a fraction from the moment before, a small bubble of water vapour forms in the body of water, say a cubic millimtre in size. This happens because the pressure on the whole body of liquid reaches the point to allow (some of) the water to boil.

    And in that bubble of water vapour will be say a quadrillion water molecules. So in what way were those quadrillion water molecules different from the quadrillion water molecules next to the bubble, which did not expand to form a bubble of water vapour?

    I still don't see how the pressure on the water molecules alters the boiling point of the water?
  11. Sep 10, 2012 #10


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    They happened to be a quadrillion water molecules with a bit higher-than-average kinetic energy that just happened to be at the water/bubble interface.

    If there were a small bubble filled with a water vapor at atmospheric pressure and just under the boiling point, 999 trillion molecules of water would evaporate into the bubble and 1000 trillion would condense out over some length of time.

    The bubble would shrink.

    If there were an identically sized bubble filled with water vapor with the ambient pressure reduced by 0.2 percent and temperature left unchanged then 999 trillion molecules of water would evaporate into the buttle but only 998 trillion would condense out.

    The bubble would grow.

    Growing bubbles = boiling.
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