## How does boiling exactly work?

After hours of reading and googling, I still do not really understand how boiling works.

I have some questions:

What I know is that boiling occurs when the vapor pressure is equal to the atmospheric pressure. But I thought that vapor pressure is the pressure exerted by the water vapor above a body of water. If the vapor is above the water surface, why are there bubbles in the body of water?

The explanation I have is that there are some pockets in the body of water where the state has changed from liquid to gas state. When the temperature is high enough, the vapor pressure exerted by these gas pockets is high enough that it can expand to form bubbles. Is that correct? If that's correct, then shouldn't the correct definition for boiling be "when the vapor pressure is equal to the atmospheric pressure PLUS the pressure from the body of water"?

Another complication is that two books stated specifically that the SATURATED vapor pressure is equal to the external pressure when boiling occurs. I don't understand why it has to be saturated vapor pressure, not just vapor pressure. Isn't it possible for the external pressure (including atmospheric pressure) to be lower than the saturated vapor pressure, so boiling can occur before the vapor pressure equals to the saturated vapor pressure?

The third question is related to open systems. I read on the web that the saturated vapor pressure of water cannot ever be higher than the external pressure (including atmospheric pressure) for an OPEN SYSTEM. My guess is that if the saturated vapor pressure exceeds the atmospheric pressure, bubbles will form and burst on the surface, “releasing” that pressure. Is that correct?

Thanks for reading, and if you find any mistakes in my thinking, please point them out; I'm always glad to learn.

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 In my experience, folks often use the term "vapor pressure" both in the sense of the partial pressure of vapor in the ambient atmosphere and in the sense of the partial pressure that would be present if the liquid in question were in equilibrium with the atmosphere. The latter usage matches the term "saturated vapor pressure" that you quote. Boiling occurs when the temperature of the liquid at a particular depth is such that the saturated vapor pressure is greater than the pressure at that depth. This means that bubbles, once formed, will tend to grow. Note that because of surface tension, tiny bubbles will tend to collapse of their own accord. A liquid which has no dissolved gasses and no nucleation sites can be heated well above its so-called "boiling temperature" before it actually boils. "Boiling temperature" is defined so that these factors do not enter in. As a practical matter, the variation of temperature with depth is greater than the variation of pressure with depth in a typical pot on a typical stove. So your pot will boil from the bottom in a so-called rolling boil. In a carefully engineered lab environment with a low rate of heating, a pot would evaporate from the top and bubbles would never form at depth, even though the pot would reach "boiling temperature".

 Quote by jbriggs444 ...in the sense of the partial pressure that would be present if the liquid in question were in equilibrium with the atmosphere...
Thanks for the reply! However, what do you mean by equilibrium? I thought you can never achieve equilibrium with the atmosphere, since in the end the atmosphere is so big you can never saturate it with vapor.

Another question is, would boiling become very difficult if it happens inside a very long vertical tube of liquid? Wouldn't the saturated vapor pressure need to be very great to counter the huge fluid pressure from the column of liquid?

 Quote by jbriggs444 A liquid which has no dissolved gasses and no nucleation sites can be heated well above its so-called "boiling temperature" before it actually boils.
I don't understand. The bubbles are made of water vapor, which is the gasesous form of water, so how do dissolved gases come into play?

## How does boiling exactly work?

 Quote by titaniumpen Thanks for the reply! However, what do you mean by equilibrium? I thought you can never achieve equilibrium with the atmosphere, since in the end the atmosphere is so big you can never saturate it with vapor.
Close the door to the lab if you have to. Or close the top of the aquarium.

 Another question is, would boiling become very difficult if it happens inside a very long vertical tube of liquid? Wouldn't the saturated vapor pressure need to be very great to counter the huge fluid pressure from the column of liquid?
Yes, the temperature required to boil water is higher when you are several miles down on the ocean floor. Note, however that there is a critical temperature beyond which water will turn to vapor regardless of pressure.

 I don't understand. The bubbles are made of water vapor, which is the gasesous form of water, so how do dissolved gases come into play?
Dissolved gasses will tend to come out of solution as you heat the water, thereby forming small bubbles. Those bubbles can act as nucleation sites, allowing steam to expand them further.

 Thanks a lot for the reply, I think I have the concepts clear now. It's a great feeling! Thanks! :D

 Tags boiling, heat, particles, thermodynamics, vapor pressure