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I am being told that boiling can't take place in closed containers. Can someone please explain why this is so.

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  • $\begingroup$ Can you reproduce the reason you've been given for why boiling can't occur in a closed container? $\endgroup$ – J. Ari Jun 19 '17 at 1:23
  • $\begingroup$ Recall the fact that boiling point depends on external pressure. $\endgroup$ – Pritt says Reinstate Monica Jun 19 '17 at 2:27
  • $\begingroup$ If 'closed container' means 'closed system', boiling of course can take place. But I'm not sure whether both stands for the same thing or not... $\endgroup$ – PenPoint Jun 19 '17 at 2:40
  • $\begingroup$ To me closed container very much means closed system. Container as in a tupperware, or a pot with a lid. $\endgroup$ – Equinox Jun 19 '17 at 3:40
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    $\begingroup$ I am being told that beyond the ocean there is an island inhabited by two-headed people. Boiling surely can and does happen in closed containers; it is just that the pressure increases in the process, and so does the boiling point. $\endgroup$ – Ivan Neretin Jun 19 '17 at 5:57
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You can understand this well by thinking about why a liquid boils in the first place. Imagine you have a pot of water, and it's just sitting there in your kitchen. Intermolecular forces (specifically hydrogen bonding for H20 molecules) keep the water in liquid state. The air molecules zipping around the pot and above the water's surface represent a certain pressure. Let's say your kitchen has standard conditions, so 1 atm. Then, you put the pot on a stove, and heat the water up. This heat breaks the intermolecular bonds between the liquid molecules. Now, the only reason the molecules rise up and become water vapor is because the kinetic energy of the water molecules matches that of the surrounding air. Pressure moves from high to low, and as the water boils, certain molecules acquire the necessary speed (and therefore pressure) to enter the gaseous state. Keep in mind that heating the pot negligibly changed the kitchen pressure. The pressure in your kitchen should still be about 1 atm, since the vast air space should distribute the heat produced by the stove.

In a closed container, however, the story is different. Let's say you put a lid over your pot. Now, there is a relatively small amount of air trapped between the lid and the water's surface. Heating the pot heats the air as well, and pressure does not remain at 1 atm. That is because that small amount of air has to handle all that heat, whereas in the situation without the lid, a bigger volume (the kitchen) allowed for more heat dissipation. So now the small amount of air molecules are moving really fast. The heat is still breaking the intermolecular forces of the water, but this time, there is no (entropic) reason to move up into the gaseous state. The air molecules are moving just as fast, and therefore have just as much pressure, as water molecules that have broken bonds.

The key to your question is rooted in why the water molecule (without a lid) moves up out of the surface of the water and joins the air molecules, effectively becoming water vapor. Simply put, it is because the water molecule (because of the heat) has enough speed to zip out of the water. Pressure is intrinsically based on speed of molecules, and so you can think about this situation as: the water molecule(s) has a higher pressure than the air molecules, and thus they move up to join the air molecules, in order to balance out the pressure. A lid prevents this from happening because the water molecule(s) pressure will never be higher than the air molecules, since the air takes the same amount of heat as the water. (The preceding statement is true only in usual cases. There are caveats to this - see comments.)

Therefore, water does not (usually) boil in a closed container. A couple of things: There are many more explanations/videos online about this topic (boiling water, pressure cookers, more altitude = lower pressure, higher boiling point etc..)in case you want to strengthen your understanding further and explore the core of this question from different angles. Also, the intuition in this Khan academy video https://www.youtube.com/watch?v=hA5jddDYcyg is pretty good.

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    $\begingroup$ Mostly OK as an explanation except that water can and does boil in closed containers. But the pressure and temperature can get very high so the point where boiling happens is very different. And can be very dangerous. $\endgroup$ – matt_black Jun 19 '17 at 23:06
  • $\begingroup$ Apologies for leaving out the caveats to my explanation. You are correct, water CAN boil in a closed container, for example, if you add heat really quickly, or like you said, high temperatures. I was only trying to confine the problem to the context of the question, and I chose not to venture outside of what the asker was getting at. Thanks for the comment. $\endgroup$ – oceandevil24 Jun 28 '17 at 1:29
  • $\begingroup$ I would clarify that water molecules are always being exchanged between the gaseous and liquid state. Boiling is the phenomenon where there is a bulk movement of water molecules, in the characteristic form of bubbles, into the gaseous phase. $\endgroup$ – Tunk May 18 at 1:16
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boiling point is the temperature at which vapour pressure of the liquid becomes equal to the external pressure. But in case of closed vessel on boiling volume remains the same so pressure increases due to increase in temperature,which also cause the pressure of the air molecules in the container to rise, along with that of the liquid present so in that case vapour pressure of the liquid remains lower than that of the air molecules and therefore boiling does not occur.

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I am being told that boiling can't take place in closed containers. Can someone please explain why this is so.

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Source: https://upload.wikimedia.org/wikipedia/commons/thumb/9/9c/Simple_distillation_apparatus.svg/2000px-Simple_distillation_apparatus.svg.png

A liquid can boil in a closed container under certain conditions. For example, you can close off a distillation apparatus by putting a balloon on the only opening with the surrounding (position 9 in the picture). As long as you condense the vapors created by the boiling liquid using a cooler, you can boil off the entire liquid. (You should allow for some expansion, otherwise the glassware will be under high pressure - a dangerous condition.)

Also, if the container is large enough (like a house), you could boil water in it on a stove top with an open pot. You can also put a lid on the pot, and it will boil a little bit faster (because the lid keeps cold air from coming in). The boiling point might rise a little over the normal boiling point because the pressure in the house will increase a little bit as you boil off the water.

Heating liquid in a closed constant volume container

If the container is closed, all vapor generated by boiling will stay in the container. If it is at constant volume, the pressure will rise in it. That is what happens in a pressure cooker at the beginning of the cooking process. The pressure rises, and so the water in it will not boil at the normal boiling point. Instead, it will start boiling at a higher temperature. In a pressure cooker, there is a valve to release gas when the pressure reaches a certain value. Without the valve, the pressure would continue to rise until it stops boiling again, or the pressure cooker will explode.

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Boiling water in a closed container essentially requires you to add heat to a constant volume. As noted above, this increases pressure, and ways to get around this basically open the system.

However, heat pipes are small closed systems (really closed, and really small), and liquids are boiled in them to transfer heat from one end to the other. Pressure increase is nominal (i.e., within design specs). https://en.wikipedia.org/wiki/Heat_pipe

The key is that you can't keep adding something (neither material nor energy) to a closed system without exceeding some limitation. Heat pipes boil a liquid at one end and condense it at the other. So the system can remain closed because the contents don't increase: you put in heat, you take out heat.

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  • $\begingroup$ Very cool (or hot) application of phase transitions! $\endgroup$ – Karsten Theis Jul 9 at 14:02

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