Why in open container and atmospheric pressure the temperature of water can't be increased more than 373K?

Question

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.

Answer 1

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.

Answer 2

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.

Answer 3

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.

Answer 4

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.

Answer 5

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.