# Why does boiling occur when vapor pressure equals ambient pressure?

Here is my understanding of the boiling process: pressure of the atmosphere is pressing down on the liquid and this pressure is propagated throughout the liquid. That is, pressure everywhere inside the liquid is equal to the ambient pressure (I'm ignoring the weight of the liquid columns for simplicity). In order for bubbles to form inside a liquid, the gas molecules must push outwards with a greater pressure than the existing ambient pressure to "make room" for the bubble. The pressure that the bubbles are able to exert on its liquid surroundings is the vapor pressure and hence we need to be at a temperature where vapor pressure at least equals the ambient pressure.

But my source of confusion is that ambient pressure is not the only pressure pressing down on the liquid. There is also the vapor pressure of the liquid in addition to the ambient pressure. Then according to my understanding of how boiling works, this means that the pressure inside the bubble (which equals the vapor pressure of the liquid) must be greater than the ambient pressure plus the vapor pressure of the liquid which makes no sense. So I'm assuming I have an incorrect understanding of how boiling works and was wondering if someone could help me correct it.

• Commented Sep 2, 2021 at 8:36
• @Poutnik Thanks for the reference. Is the answer in this post basically saying that the vapor pressure is higher at the bottom of a container than at the top of the surface (due to higher pressure at the bottom), allowing the bubbles to form at the bottom? Commented Sep 2, 2021 at 9:13
• Which answer in which post ? (the link was a hint for you that you are supposed to do basic search before posting a question) // Saturated vapor pressure depends on T, not p (1). But when p is higher, higher T is needed to p_vap = p, therefore the boiling point slightly depends on the depth. As convection is involved, there may occur streams of temporarily overheated metastable liquid.// (1) In your Q context, otherwise things are more complicated in tiny details. Commented Sep 2, 2021 at 10:02
• The ambient pressure is just one, there is no such a thing you describe. Or you have a pressure cooker, and indeed boiling stops. Commented Sep 2, 2021 at 10:18
• @KarstenTheis, sorry, that was a bad pun on supercritical -- above critical temperature and pressure, there is no boiling. Commented Sep 3, 2021 at 16:16

But my source of confusion is that ambient pressure is not the only pressure pressing down on the liquid.

Pressure does not press down. When something is under pressure, it exerts it in all directions. Maybe it would help imagining this problem in the absence of gravity (you would have to put the sample in a stretchy balloon to get some pressure while being able to change the volume).

So the balloon would exert some pressure on the liquid. If the liquid is hot enough that the vapor pressure is higher than the pressure, and if you can get some tiny bubbles going, they will expand as more and more liquid goes into the vapor phase, increasing the volume of the bubbles while maintaining sufficient pressure not to collapse.

• Thanks for the reply. So to summarize: the reason why the boiling occurs when vapor pressure equals ambient pressure is bc bubbles would otherwise collapse if the vapor pressure inside the bubble was less than the ambient pressure it feels from the liquid. Is that correct? Commented Sep 3, 2021 at 10:40
• @rofimfao Yes. Getting it to make tiny bubbles is a question of kinetics and nuclei forming, more complicated... when you listen to a pot about to boil you can witness some of that
– Karsten
Commented Sep 3, 2021 at 18:55
• @rofimfao Here a link to the question exploring the sounds: chemistry.stackexchange.com/q/125368/72973
– Karsten
Commented Sep 3, 2021 at 19:34