Boiling point and vapour pressure

[americast]

Hi all,
I am facing difficulty in understanding the reason behind the fact that a liquid boils when the atmospheric pressure is equal to its vapour pressure.

We know that at the vapour pressure, the air and the liquid remain in equilibrium, so there's no net evaporation or condensation. So, why does this not hold true at the boiling point? According to me, at the boiling temperature, there should be no vaporization at all.

Gramercy...

 

[russ_watters]

There is no net evaporation if the pressure of the vapor equals the saturated vapor pressure at the liquid's temperature. Boiling occurs when the vapor pressure equals atmospheric pressure.

That's 14.7 psi atmospheric pressure or boiling saturated vapor pressure for water, but at room temperature saturated vapor is only 0.4 psi (0.027 bar).

 

[americast]

Okay. So, at the boiling temperature, the liquid is in equilibrium with its vapours, not the entire air, right?

Can you tell me why does a liquid boil only when the outside pressure is equal to its vapour pressure at that temperature? What is the reason behind this phenomenon?

 

[Chestermiller]

Okay. So, at the boiling temperature, the liquid is in equilibrium with its vapours, not the entire air, right?

Can you tell me why does a liquid boil only when the outside pressure is equal to its vapour pressure at that temperature? What is the reason behind this phenomenon?

Under these conditions, bubbles can form under the liquid's surface because the gas within them has a high enough pressure to physically push the surrounding liquid back. If the vapor pressure were less than the total pressure, if a bubble were try to form, it would collapse.

Chet

 

[russ_watters]

Okay. So, at the boiling temperature, the liquid is in equilibrium with its vapours, not the entire air, right?

No, with boiling it is versus the entire air. And it isn't really an equilibrium either: that's why boiling can be so violent.

Can you tell me why does a liquid boil only when the outside pressure is equal to its vapour pressure at that temperature? What is the reason behind this phenomenon?

Again, you have it backwards. But the "why" of boiling is simple: since the saturated vapor pressure inside the water is at/above atmospheric pressure, bubbles spontaneously form inside the liquid. Normal evaporation, on the other hand, only occurs on the surface because the air pressure is high enough to prevent bubbles in the liquid.

[Same idea, slightly different wording from Chet]

 

[americast]

Thanx... I think I have got it!

No, with boiling it is versus the entire air. And it isn't really an equilibrium either: that's why boiling can be so violent.

Say, while boiling, the air above the liquid comprises 1% vapour (of that particular liquid) and rest ordinary air (other gases). In this situation, the liquid is in equilibrium with that 1% of its vapours only while it is not in equilibrium with the rest of the air. And thus boiling takes place, am I right?

Gramercy again...

 

[russ_watters]

Say, while boiling, the air above the liquid comprises 1% vapour (of that particular liquid) and rest ordinary air (other gases). In this situation, the liquid is in equilibrium with that 1% of its vapours only while it is not in equilibrium with the rest of the air. And thus boiling takes place, am I right?

Sorry, I can't make any sense of that, but it looks wrong. Boiling liquid really isn't in an equilibrium with anything. Not with the air and not with the water vapor in the air.

 

[americast]

Boiling liquid really isn't in an equilibrium with anything. Not with the air and not with the water vapor in the air.

But, why not?
While boiling, the entire air (including the small fraction of vapour) above the liquid is at the vapour pressure of that liquid (at that particular temperature). So, at least, the vapours of the liquid should be in equilibrium with the liquid. That is how vapour pressure is defined...Vapour pressure is the pressure exerted by the vapour on the liquid when the two are in an equilibrium.

 

[Chestermiller]

But, why not?
While boiling, the entire air (including the small fraction of vapour) above the liquid is at the vapour pressure of that liquid (at that particular temperature). So, at least, the vapours of the liquid should be in equilibrium with the liquid. That is how vapour pressure is defined...Vapour pressure is the pressure exerted by the vapour on the liquid when the two are in an equilibrium.

The partial pressure of the water vapor is equal to the equilibrium vapor pressure (a) within the rising bubbles in the liquid and (b) at the surface of the water (i.e., the interface with the air in the room). However, further out into the room, the partial pressure of the water vapor is less than the equilibrium vapor pressure, even at the ambient temperature of the room. So, there is a substantial gradient of water vapor concentration in the room air, with very high values at the liquid surface, and much lower values away from the surface. The gradient in concentration is responsible for diffusive transport of water vapor away from the interface, in conjunction with the convective motion of the air within the room. As long as the water is boiling (and, if the room air is not circulated much through the house), the room humidity will gradually increase.

Chet

 

[russ_watters]

But, why not?
While boiling, the entire air (including the small fraction of vapour) above the liquid is at the vapour pressure of that liquid (at that particular temperature).

In a non-boiling equilibrium situation, "equilbrium" refers to the fact that no net liquid is leaving or entering the liquid. That situation does not exist in boiling. If you want to say that there is some kind of pressure equilibrium, that situation always exists, so there isn't any special situation that would be useful to call an "equilibrium". Either way, it is not referring to the same thing.

So, at least, the vapours of the liquid should be in equilibrium with the liquid. [when boiling]

Again, in what sense? At a zero mass flow? No. At the same (vapor) pressure? No. At the same temperature? No.

That is how vapour pressure is defined...

But it isn't how boiling is defined.

Vapour pressure is the pressure exerted by the vapour on the liquid when the two are in an equilibrium.

Yes.

Maybe I see what the issue is. You think because the definition of boiling includes the words "vapor pressure" that an equilibrium needs to exist because the definition of vapor pressure includes the concept of equilibrium. This isn't the case. Vapor pressure isn't only applicable when there is an actual equilibrium, it is also a theoretical value that can be applied in situations when there isn't an equilibrium. For the boiling point, you can say that if the actual pressure of the vapor were equal to atmospheric pressure, there would be an equilibrium, but obviously it isn't: the atmosphere is mostly nitrogen and oxygen, not water vapor.

Indeed, the difference between the actual pressure of the vapor and the theoretical vapor pressure value of the water at its temperature is information that can be used to figure out how fast it will evaporate (a non-equilibrium situation).

Perhaps the definition of "vapor pressure" is confusing and implies any partial pressure of vapor. But really, it's more about what is happening to the liquid than the vapor. Try it this way for the theoretical definition: it is the pressure of vapor required to stop the liquid from evaporating.

 

[americast]

For the boiling point, you can say that if the actual pressure of the vapor were equal to atmospheric pressure, there would be an equilibrium, but obviously it isn't: the atmosphere is mostly nitrogen and oxygen, not water vapor.

Exactly. The very small fraction of water vapour present above the boiling water should have been in equilibrium with the liquid. Now I understand that this does not happen as convection takes those vapours away from the surface of the boiling liquid.

Thanx a lot!