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Why doesn't boiling water exceed 100 deg C?

  1. Aug 15, 2014 #1
    I know that water can be superheated, but my question is more fundamental. What causes water to boil at 100 deg C? I know that there are several factors (atmospheric pressure, nucleation sites, hydrogen bonding, surface tension, etc) but that's not what I'm getting at. I'm interested in the more fundamental reason as to why water boils at the temperature that it does (why a boiling solution doesn't exceed its boiling point). I'm not sure if I'm making sense as some people will probably tell me that I just answered my own question, but idk, why can't one "input" more energy into the system and cause water to boil at a higher temperature than 100 deg C?
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  3. Aug 15, 2014 #2


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    Pure water boils at 100 C only when the ambient pressure is 1 atm. If you were to take your pot of water up to the top of a tall mountain and start to boil it, your pot of water would boil at a temperature significantly less than 100 C, because the ambient pressure is no longer 1 atm. For example, at an elevation of 3000 m above sea level, the boiling point of water is reduced to about 90 C.


    Conversely, if one heats water in a closed environment where the pressure is greater than one atmosphere, boiling will not occur until the temperature exceeds 100 C.

    The boiling point of water can be changed also by dissolving another substance in it or mixing it with another substance. That is the principle behind adding anti-freeze to the water in a car radiator: not only is the boiling point of the mixture raised above 100 C, but the freezing point of the mixture is reduced below 0 C. On a hot day, the water in the radiator stays liquid above 100 C because the cooling system also is pressurized to a certain extent. If the radiator cap is removed suddenly, the pressure inside the cooling system drops at once to ambient, and the hot coolant can begin to vaporize.

    This page explains a little more behind the physics of the boiling point:

  4. Aug 15, 2014 #3


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    I think all of that is true and none of it answers the question, which I think can rationally be extended to mean (and I read it as meaning) "why does pure water boil at exactly 100 at sea level? Why can't you raise the temperature to 105 degrees?"

    I don't know why. I know what HAPPENS if you try to raise the temperature ... you just boil it faster.

    If you could manage to hit the water with a massive dose of microwaves throughout its mass all at once such that all localized areas boiled, then you'd likely get superheated steam at least momentarily until the pressure had time to spread the steam out but I'm not sure that counts as the temperature of the water.
  5. Aug 15, 2014 #4
    Superheated liquid-water in microwave on YouTube ... http://youtu.be/1_OXM4mr_i0?t=23s
  6. Aug 15, 2014 #5
    I know that you can superheat the liquid, but that's not what I'm getting at. Like phinds rephrased, why does water boil at 100 deg C at STP? (I know that intermolecular forces etc come into play, but that's not what I'm getting at.) Before anyone brings up the latent heat of vaporization etc that's also not what I'm getting at. Another way to ask is this: I have an open pot of boiling water and want to bring the boiling water temperature to 250 degrees C. Why can't I?
  7. Aug 16, 2014 #6


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    Because 100 deg C is the temperature that water molecules have enough energy to overcome their intermolecular bonds and break away from each other and become a gas.

    It's a result of how strongly the molecules are attracted to each other, if the intermolecular bonds were (magically) made stronger then the boiling point would rise; a higher temperature (average kinetic energy) would be required to break the bonds and free the molecules into a gas.

    (Disclaimer; A Mechanical engineering students intuitive understanding)
  8. Aug 16, 2014 #7
    Allegedly that is possible ...
  9. Aug 16, 2014 #8


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  10. Aug 16, 2014 #9
    I havnt read the entire thread, but liquids boil when their vapour pressure is higher than the enviroment, you read some about it.. Its general thermodynamics. However, what that mean is if you raise the pressure on a liquid, the vapour pressure will have to be higher in order to exceed the external pressure...
  11. Aug 16, 2014 #10
  12. Aug 16, 2014 #11


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    Boiling implies a phase change from liquid water to steam (water vapor). At 1 atm, water boils (changes phase) at 100°C. The boiling point can be raised by increasing pressure (as in a pressure cooker) or decreased by lowering the pressure as evidenced by SteamKing.

    The phase change occurs when the molecules in the liquid receive sufficient energy to overcome the interatomic forces that maintain the material in liquid form.

    Often we experience saturated steam/liquid, as opposed to dry steam (which is superheated).

    In industry we use 'pressure vessels' heat steam at high pressures. Sufficient energy must be input to superheat the steam, such that the temperature increases above the 'saturation' temperature.

    We have superheated or dry steam, but also supercritical steam ( > 22 MPa).
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