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I am a high school student and I am very confused in understanding the phase diagram and boiling. It seems like most people didn't understood the question well may be because I didn't explain the question well and concise, so I am editing this, lets take an example of water. We say at "atmospheric pressure" the boiling point of water is 373.15K and boiling occurs when vapor pressure is equal to the external atmospheric pressure. My questions are:

1)Why the temperature of water can't be increased more than 100 degrees?: By reading many answers and textbooks explanation ,It seems like its because as we increase the temperature of water ,at any temperature less than 100 degrees only its surface molecules gets enough energy to escape and if any other molecules in the bulk got enough energy somehow and started forming bubbles by separating from each other the bubbles will collapse because the external pressure is more than the bubble pressure but as the temperature of "some" region {because if all the regions have exact 100 degrees temperature why don't all water molecules escape?} reaches 100 degrees{at I atm pressure} molecules in the bulk there form bubbles they now got enough bubble pressure to "overcome" the external pressure because at this temperature more water molecules get separate increasing the bubble pressure and they rise and leave and as other molecules also reaches 100 degrees the process repeats. So in this explanation it seems the heat we are still supplying at 100 degrees is going in breaking bonds in the bulk and as they break those molecules form bubbles of steam which escapes. but I have following confusion with this explanation:

  1. How we can say that the pressure inside those bubbles is exactly equal to the "equilibrium vapor pressure" at 100 degrees? I mean they are two different things .Vapor pressure is the pressure of the water molecules which is above the liquid and those vapors come from just the surface molecules of the liquid which got converted to steam and on the other hand the bubbles are forming because of the separation of water molecules which got enough energy as they get at 100 degrees and breakes their H-bonding, so the region there got less dense and formed bubbles of steam which rises up. I mean bubbles can form at any temperature because at any temperature some molecules will have enough energy to how can we surely say the pressure inside those bubbles would be guided by the vapor pressure at that temperature?

2)If we see Raoult's law of vapor pressure ,it tells the vapor pressure at any temperature decreases even on adding "ideal non volatile solute" to the liquid because the solute will occupy some space at the surface which was earlier occupied by liquid molecules. But how will you explain the increase in boiling point with the explanation of boiling stated above? because an ideal solute doesn't changes the bonding structure of the liquid solvent ,so if bubbles are forming at 100 degrees because only at this temperature the molecules in the bulk gets enough bubble pressure so on adding ideal solute the bubble pressure should still remain the same because bonding is not changed by it so the same no. of water molecules should form bubbles making just enough bubble pressure to overcome the external atmospheric pressure.

Arun Bhardwaj
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  • Passing gigantic heat amount instantly into water of liquid form,- it will make whole water a vapor instantly, then if you'll repeat the process, vapor will be converted into plasma of mixed gas. You can't pass heat without affecting internal energy of absorbing system at hand. – Agnius Vasiliauskas Jul 13 '22 at 11:49
  • “I know Vapor pressure cannot rise more than 1atm, but even at 100 degrees its not 1atm for open container and would not be more than 1atm.” Both statements are incorrect. The vapor pressure passes smoothly through 1 atm at 100°C and continues to increase. But at this point, the water gas has the pressure to push liquid water out of the way, allowing large amounts to escape and expand quickly into the surroundings. – Chemomechanics Jul 13 '22 at 15:01
  • @Chemomechanics why the temperature of water cannot increase further than 100 degrees though? if it would increase we would just have more vapor pressure what's the problem with it? – Arun Bhardwaj Jul 30 '22 at 11:53
  • Broadly, since gaseous water is the stable phase above 100°C, any heating of the liquid water is absorbed immediately by the boiling process. The latent heat of vaporization is large that it sucks away any heating of the liquid water. It’s only after boiling is complete that further heating produces a temperature increase (of the gas). – Chemomechanics Jul 30 '22 at 19:08
  • Your edited question is not terribly easy to read, to be honest; there is a lot of text there and it is not easy to pick out what your actual question is. Can you edit it down to be clearer? – Michael Seifert Jul 31 '22 at 13:22
  • I first explained what I think boiling is and why temperature of water doesn't increase beyond 100 degrees at 1 atmospheric pressure from reading different texts and answers {so that people can correct me if its completely wrong} and based on that I put forward two things which seems wrong to me with that explanation . If I made it shorter I don't think people would get the picture I have in my mind. – Arun Bhardwaj Aug 01 '22 at 12:11
  • basically the questions are: 1) if my understanding of boiling is correct(for high school level and for my purposes) then it doesn't seem to follow raoult's law on a deeper level even on adding ideal solute 2)it doesn't clear why the equilibrium vapor pressure at any temperature has to be equal to inside bubble pressure at that temperature? – Arun Bhardwaj Aug 01 '22 at 12:15

5 Answers5

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FWIW, it is possible to heat water above 373.15 K without achieving boiling conditions. It's not easy, but with a proper container and proper care, you can get a super-heated liquid. Boiling, as you may have observed, initiates more easily of there are sharp-edged solids (like a rough pot interior) for bubbles to build up. Guess I should add that the phase diagrams indicate stable conditions, and super-heated or super-cooled situations are unstable.

Similarly, you can super-cool liquid water below 273.15 K if you are careful. Then drop a small solid, or shake the container, to get instant total conversion to ice.

Carl Witthoft
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  • can u tell me why after a certain temperature , whole ice melts into water? I mean its a reversible reaction so I think equilibrium should exists at all temperature just the equilibrium constant should shift, but instead the temperature remains constant till phase transition and equilibrium remain established{it just shifts, so it means here equilibrium constant is changing WITHOUT change in temperature} but when temperature starts rising quilibrium is lost ,,,so does it become irreversible? – Arun Bhardwaj Jul 15 '22 at 09:07
  • @ArunBhardwaj I suggest you get a college-level first year chemistry text (or equivalent online) to learn the basics of material states and transition functions – Carl Witthoft Jul 15 '22 at 10:44
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Note that a thermal inkjet printhead raises the temperature of the water-based ink in it to 270 degrees C at which point the ink in contact with the heater spontaneously explodes all at once into steam- which forms a bubble that kicks out a drop of ink.

You can superheat water in this way because it takes a certain amount of time for a vapor bubble to form in water at its normal boiling point of 100 C. If you dump in heat faster than the time it takes for the first vapor bubble to form and refrain from disturbing the hot water in any way, you can get the water superheated and thereby create a very energetic superheat vapor explosion in the water.

This is what sometimes makes a cup of water explode into steam when you are heating it in a microwave oven, just as you are removing the cup from the oven cavity. With practice and care, you can routinely superheat a 1" depth of water in a Snapple bottle and create a spectacular superheat explosion in the microwave which will blow all the water out of the bottle and create a big mess inside the oven.

As the superheat becomes more extreme, this initiatory time lag for the explosion tends to zero and you will then hit the thermodynamic limit of superheat for the water which is about 300 degrees C, compared to the "normal" boiling point of 100 C.

niels nielsen
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At standard atmospheric pressure and at 100C (373K) water boils - it undergoes a phase transition where from a liquid it becomes a vapor: it is still water in the sense that it is the same chemical substance, but it is not water anymore, in the sense that it is not a liquid - and this is what is meant when we say that it cannot be heated beyond 100C. Further heating means heating the vapor, but not the water.

Roger V.
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  • I simply mean that even when water is boiling , it doesn't converted to vapor as a whole even at 100 degrees and much amount is still present in the liquid state in the container, So why can't the temperature of this liquid doesn't increase more than 100 degrees on supplying heat? why is it that only at this temperature the heat supply would result in phase transition and not increasing temperature of liquid state? – Arun Bhardwaj Jul 13 '22 at 11:54
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    See Latent heat - once water has reached the boiling point, the heat provided to it goes into breaking the bonds between molecules, rather than increasing their average kinetic energy, i.e., the temperature. Once these bonds are broken, the temperature will start increasing again, but now it is not water anymore. – Roger V. Jul 13 '22 at 11:59
  • Perhaps it could add to understanding - the water does not become a vapor immediately when it reaches 100C. One needs to continue heating it, so that water at 100C converts into vapor at 100C. – Roger V. Jul 13 '22 at 12:02
  • I am asking it on a more basic level, "why only after 100 degrees the further heat supplied is used in breaking bonds" not before and after, why? just because at 100 degrees the saturation vapor pressure is 1 atm what is it even significant? – Arun Bhardwaj Jul 13 '22 at 12:03
  • Saturated vapor pressure is another relevant term here. The ease at which bonds can be broken or restored depends on pressure - at lower pressures (e.f., at a mountaintop) water boils at lower temperature. – Roger V. Jul 13 '22 at 12:10
  • "at high pressure boiling occurs at high temperature" doesn't necessarily means "only at that temperature when its SVP becomes equal to atmospheric pressure" .Can you prove it mathematically by any means? No right? We just know it as facts and trying to act as if its logical and obvious – Arun Bhardwaj Jul 13 '22 at 12:15
  • SVP is the atmospheric pressure here - we are dealing with a constant pressure process, i.e., we assume that the pressure is equalized very quickly and always remains at 1Pa. We could have a situation where this is not the case, so that the vapor pressure just above boiling water is higher, so that the SVP would be higher and the water would be heated to a higher temperature... but then we would not be dealing with constant pressure conditions, and there would be no sense to talk about atmospheric pressure. – Roger V. Jul 13 '22 at 12:22
  • SVP is not the atmospheric pressure here, how can it be? the definition of SVP is" the vapor pressure in quilibrium with the liquid phase" so clearly it only exists at equilibriums and In open container for all temperatures there is "no equilibrium" and V.P at any temperature would be lesser than the SVP of that temperature, so what's so special in 100°C ? why the relation that" at those temperature whose SVP=atmospheric pressure " is required for boiling? please think from the starting and consider what's happening at all temperatures rather than trying to justify what books directly says – Arun Bhardwaj Jul 13 '22 at 12:38
  • The vapor pressure in an open vessel is the same as the atmospheric pressure - otherwise we will have pressure gradient that will result in the wind equalizing pressure. – Roger V. Jul 13 '22 at 12:43
  • the total pressure above the liquid is always 1atm and I have written this in the question too but that is made of both watervapors and the air ......its true for all temperatures and it has nothing to do with my question so please revisit it...........In one line my question is "why the relation that"at those temperature whose SVP=atmospheric pressure is required for boiling"? like here the SVP of water at 100 degrees is 1atm so we say boiling occurs and for temperatures less than this ,SVP<1atm we say boiling will not occur . Why boiling has anything to do with this relation? – Arun Bhardwaj Jul 13 '22 at 12:47
  • @ArunBhardwaj I merely tried to help you. If you are not content with my answer, no need to get upset. Hopefully someone else explains it better than me. – Roger V. Jul 13 '22 at 12:53
  • I apologize if I have used any wrong language by mistake, I got in some rage mode because this is creating too much confusion in my mind and people treat it as "obvious" .I hope you will understand me – Arun Bhardwaj Jul 13 '22 at 12:57
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Boiling means that bubbles of pure water vapor form below the surface of the liquid; there is no air below the liquid surface, but the pressure at these locations is essentially 1 atm because of contact with the air above (very low hydrostatic contribution to the pressure). So the bubbles of water vapor are in equilibrium with the liquid water at 1 atm. The bubbles then rise and leave through the surface.

Chet Miller
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  • but how does it explain that liquid water temperature shouldn't rise above 100 degrees? – Arun Bhardwaj Jul 13 '22 at 12:01
  • 100 C is the equilibrium vapor pressure of water at 1 atm. Since the pressure can't get higher than 1 atm, the temperature can't get above 100 C. – Chet Miller Jul 13 '22 at 12:04
  • say water is at 60 degrees in open container, still vapors are forming and diffusing into the surrounding right?here also there is no equilibrium and the vapor pressure above liquid is lesser than saturation vapor pressure at 60 degrees same thing is happening at 100 degrees{V.P LESS THAN 1ATM}and even if the temperature would increase more than 100 degrees the vapors would form and diffuse out keeping vapor pressure above the liquid lesser than 1atm. so it doesn't tell why it can't increase more than 100 degrees. You are confusing two concepts "equilibrium of vapors and water" and "boiling" . – Arun Bhardwaj Jul 13 '22 at 12:10
  • if you could simply tell me that why "SVP{saturation vapor pressure}=Atmospheric pressure" this relation is necessary for boiling then I would think you have answered my question. – Arun Bhardwaj Jul 13 '22 at 12:17
  • In order for a bubble to grow under the water surface, it has to push the water surrounding the bubble (which is at 1 atm) back. So the water vapor pressure within the bubble has to be 1 atm. (to balance the pressure of the surrounding water). – Chet Miller Jul 13 '22 at 14:15
  • that's the minimum pressure needed, the maximum can be anything so it doesn't explain why temperature can't be more than 100degrees ,,,,if it would be more than 100 degrees the SVP would be higher than atmospheric pressure what's the problem in it?? – Arun Bhardwaj Jul 15 '22 at 09:01
  • also how would you explain the rise of boiling point in ideal solutions?? in ideal solution the same force of attraction exists between each pair so it wouldn't affect the vapor formation INSIDE bulk liquid and it shoud still occur at 100 degrees?? I know in books VP gets affected because there are fewer water molecules over surface but inside also the concentration is changing but the bonding of water molecules is still same so for THEM to become vapors the temperature should still be 100 degrees right but it doesn't? – Arun Bhardwaj Jul 15 '22 at 10:53
  • We are talking about boiling of pure substances here, right? – Chet Miller Jul 15 '22 at 22:48
  • For an ideal solution, do you think that the vapor pressure curves of the two components are identical? – Chet Miller Jul 16 '22 at 11:56
  • I have written that vapor pressure{which actually is outside the liquid} will schange on adding solute,,,,,I am saying that how the solutes affect the bubble formation?the bonding is not changes so "INSIDE" bubbles should still have 1atm pressure at 100 degrees only?if you think boiling is characterized by bubble formation which actually happens inside the liquid. – Arun Bhardwaj Jul 19 '22 at 06:05
  • Even in multicomponent boiling, if the external pressure is controlled to 1 atm, the total pressure inside the bubbles is 1 atm. Of course, the temperature doesn't have to be 100 C. – Chet Miller Jul 19 '22 at 10:50
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First, water molecules are sticky. Think of tape. They stick together when in contact. But if you pull hard enough, they come apart.

Second, water molecules are vibrating madly and slamming into each other violently. According to Molecular Speed Distribution, typical speeds are $1500$ mph in $100^o$ C water.

When water boils, it is because the molecules vibrate or collide so hard they fly apart.

There is a range of speeds at any given temperature. The faster molecules will come apart and fly away, leaving slower ones behind. This is why water stays at $100^o$ as it boils. Another factor is that it takes a certain amount of energy to break the bond. This slows the separated molecules.

Even below boiling, some of the molecules are vibrating fast enough to come apart. So evaporation can occur without boiling.

Air pressure plays a role. If there are lots of air molecules around, evaporating water molecules will run into them. They might be bounced back into the water. They might lose energy to the air molecules. The net effect is to suppress boiling and slow evaporation.

mmesser314
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