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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).
    http://www.elp.com/articles/print/volume-81/issue-1/power-pointers/primer-on-supercritical-steam.html [Broken]
    Last edited by a moderator: May 6, 2017
  13. Jan 6, 2017 #12
    I don't think i can give you a perfect explanation, but will try to make a sense.

    As everyone says, Temp(t) is average kinetic energy of the molecules ( or atoms ). Let's assume t be the value of average kinetic energy for the molecules in the water that is being heated up. As you heat water you increase the kinetic energy of molecules and the average kinetic energy of the molecules increases ( t value increases )

    Now to understand why water cannot hold average kinetic energy beyond a value (i.e Temp value beyond a point or 100 C at sea level), we need to understand Vapour Pressure. This is a simply the pressure that is exerted by the water molecules in gaseous state just above the liquid surface ready to evaporate. The vapour pressure of the water ( force exerted by gaseous state water molecules)along with atmospheric pressure put the liquid molecules in vessel ( container or pot whatever) form escaping into the atmosphere.

    Because you are increasing the kinetic energy of the molecules by supplying heat, the molecules move faster and randomly get out of liquid. This is where vapour(along with atmospheric) pressure comes in and conflicts if the pressure applied by the vapour on the surface is more, then the molecule is forced back into the liquid but if the pressure if weaker than the kinetic energy it escapes into atmosphere, what we call as boiling (vapourisation).

    Vapour pressure of water is 1 atm at 100 C that is equal to the atmospheric pressure. So when the average kinetic energy( t ) of the molecules reach the value of 100 (as we assumed at first) they have equal kinetic energy as that of vapour pressure. Further addition of kinetic energy ( temp ) would give the molecules just that final kick needed to overcome that pressure to vaporise. So they just boil at that point. That clearly explains why we cannot rise the temp of liquid beyond boiling point at a pressure. As you asked why we cannot input more energy into system, the energy you put is simply used by molecules to overcome that pressure get vapourised. Why would any one be staying back at college after getting their degree?? to increase the strength of college? , similarly why would molecules stay back in the pot when they've got energy to leave it.

    To make it simple it's just like a tug of war between the energy of liquid molecules (kinetic energy) and vapour pressure at 100 C, if liquid molecules get enough energy they win over vapour and leave the pot as soon as they cross 100 C. If they haven't they just stay and get energy till they reach 100 C.

    BTW 100 C doesn't define water boiling point, where water boils is defined as 100 C hence we have such perfect valued boiling point.

    Hope it helped. ;)
    Last edited: Jan 6, 2017
  14. Jan 7, 2017 #13


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    You see the highlighted sentences I think you have a clue why the student might not feel his question is answered yet. He asks why does water boil at 100°, we answer him it doesn't always which I expect he knows, and other answers I expect he knows. He probably doesn't feel that these have answered the question why does it boil at all? There is vapour and vapour pressure at all temperatures. But we don't in common experience see a gradual equilibrium gradually changing from more to less water, less to more vapour. It seems to me that much of the explanation given sounds very like explanation of a temperature dependent chemical equilibrium. But there you always have left a little bit however small of the unfavoured substance in the equilibrium. You don't have small drops of liquid water which are just very tiny above 100°.

    In the quoted analogy the students don't just drift away one by one until at the end of the year there arenone left. They nearly all leave at once - and moreover implicitly a causal explanation has been given for that. So doesn't somebody need to add some explanation involving a highly cooperative process in which molecules faraway are helping to keep a molecule at the surface from leaving and vice versa in such a way that almost one leaves, we all leave?
  15. Jan 8, 2017 #14
    Thank you for your perspective, I just tried to make sense why that happens (boils at 100 C)so that any body who reads it can understand without any preliminary knowledge and I've clearly stated that at the beginning.

    The lines highlighted by you are there to add humor if you can understand ;) , though they literally are not false. As I said if I had acquired the knowledge and a degree so as to consolidate the prior , I wouldn't be staying at the college wasting my time. Yes student wont drift away one by one but, if someone has got what they need will they be staying for everyone to get it? this holds the same for boiling, if a molecules has got enough energy it just leaves because its inter molecular bonding has been broken. This is what I intended to say through that line. Like you said that students wouldn't go until the end of the year, would be agreeing that every student took exactly one year to get that knowledge ?? Dude I just casually correlated AN INDIVIDUAL MOLECULE with A STUDENT that's it, if that misled you, I wouldn't hesitate to ask for an apologize.

    There is vapour and vapour pressure at all temperatures. yes there is, but is not equal to atmospheric pressure until the temp is 100 C. That obviously makes sense that water has got enough pressure to turn into vapour despite the pressure of atmosphere to keep it in the pot. There are various values for vapour pressure of water at various temperatures, but every value is less than atmospheric pressure and that is the reason why water doesn't boil. Until the water gets enough pressure against atm it would just stay gaining energy. At 100 C because the atm no longer is able to keep them together they just boil to from vapour.

    "But we don't in common experience see a gradual equilibrium gradually changing from more to less water, less to more vapour. It seems to me that much of the explanation given sounds very like explanation of a temperature dependent chemical equilibrium. But there you always have left a little bit however small of the unfavoured substance in the equilibrium." Sorry I didn't get that, I would be delighted if you can elaborate it for me. I wanted to be simple so I hadn't written any equilibrium stuff.

    To my knowledge I had never said that.

    The topic was about boiling water not exceeding 100 C and I think my answer gives a reason.

    What you asked is primarily why boiling occurs and what he asked for was " why does that happen only at a particular point? " Hope you got the difference.

    Answer to your question is relatively simple to understand, boil water in closed container, if there is not space for the vapour to exist in the container theoretically the water would not boil no matter how high the temperature is ( provided the container can withstand the vapour pressure at those temperatures). What is stopping/keeping water molecules at the surface from leaving? It's undeniable that the container walls (atoms/molecules of container) are responsible for holding the molecules in liquid phase. Similarly atmosphere (atoms/molecules present in the atmosphere ) is responsible for maintaining a liquid phase until is can hold ( i.e the pressure it can apply = 1 atm) , if it cannot hold beyond the pressure of 1 atm what would keep the water in the pot so that you can give more heat energy.

    Just try to visualize the concept, I hope you wont take long to get clarified.
  16. Jan 8, 2017 #15
    This is basically not correct. Water vapor can escape from the pot at temperatures substantially below 100 C (even if the total pressure is 1 atm) simply by evaporating. You have come perilously close to perpetrating misinformation (probably, unintentionally). But, if you persist along these lines, you will receive warning points and possible banning.
  17. Jan 8, 2017 #16
    This thread has been inactive from over two years, and the responses provided in the posts from 2 years ago adequately addressed the OP's question. Therefore, I am hereby closing this thread.
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