Does pressure affect equilibrium vapor pressure (or RH)?

In summary, at atmospheric pressure, water will boil at 100C. However, when external pressure is reduced to the point of boiling, water will boil at a lower temperature.
  • #1
Elquery
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TL;DR Summary
Does atmospheric pressure affect the equilibrium vapor pressure of water vapor? What about partial pressure, and what about relative humidity?
I'm wondering if I'm on the right track and if anyone is willing to steer me on if not:

Equilibrium vapor pressure (EVP—also referred to as saturation vapor pressure) is dependent only on temperature. Outside pressure has no bearing.
Now, of course, with lower external pressure (atmospheric), water will boil at lower temperatures. This led me to initially believe that EVP is affected by pressure. At the extreme, a vacuum would cause total and near instantaneous phase transition to vapor. Therefore it must be that the EVP has been severely lowered.

Maybe not. The EVP of water hasn't changed, it's simply that the outside pressure has been reduced to—or below—that pressure.

Thought experiment: Imagine a sealed container with water in it. Then imagine pulling a perfect vacuum in that container. If the container is large enough, all water will evaporate (boil) and fill the space with water vapor at some pressure below the EVP (for that temperature). If the container is small enough, however, evaporation can cause the space to reach the EVP and this vapor pressure will cause equilibrium to be reached while some liquid water remains. So it would essentially boil until the vacuum was filled by enough vapor to reach EVP, at which point any remaining liquid water would stay as such.

Hopefully I'm sort of on track.
One final question assuming the above is true: Does pressure affect relative humidity?
Relative humidity is the partial vapor pressure / EVP. And if it's true that pressure does not affect EVP, then the only way it would is if it affects partial pressure. The factors that affect partial pressure elude me. It seems like its just a function of a dynamic state that hasn't reached equilibrium.

At the extreme, it would seem that if you lower the pressure enough to cause boiling, you would have increased the likelihood of a fully saturated state. If this extreme applies on a scale prior to boiling point, then lowering pressure would be likely to increase partial pressures and hence RH at a given temperature. Not sure about this...
 
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  • #2
Elquery said:
...water will evaporate (boil)
Evaporation and boiling are not the same thing. Please state your understanding of what each those terms mean. Recognizing the difference will probably answer your question...
Does pressure affect relative humidity?
No.
At the extreme, it would seem that if you lower the pressure enough to cause boiling, you would have increased the likelihood of a fully saturated state.
Do any of these processes depend on "increased likelihood"? What does that even mean?
 
  • #3
Elquery said:
What about partial pressure, and what about relative humidity?

They are strictly related, so it is in fact a single question, not two.
 
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  • #5
Elquery said:
At the extreme, it would seem that if you lower the pressure enough to cause boiling, you would have increased the likelihood of a fully saturated state.

In a sealed container you're going to end up with saturation no matter the temperature and regardless of the presence of air. The term "saturated air" is sort of confusing in that the air itself is irrelevant with respect to the amount of water vapor present.
 
  • #6
russ_watters said:
Evaporation and boiling are not the same thing. Please state your understanding of what each those terms mean. Recognizing the difference will probably answer your question...

Let's see. I think evaporation is the process of water molecules leaving the liquid state and entering the gas state. At equilibrium, the same amount will leave that re-enter. It can be viewed statistically. It occurs at the boundary.

Boiling point is the temperature at which the 'external pressure' (not sure if that's the best language to use) on the system equals the EVP. So at atmospheric 'external pressure', the temperature is around 100C. When this happens, water molecules entering the vapor state can do so from within the liquid mass (bubbles can form), thereby eliminating the restriction that it occur at the boundary.

I guess the only fundamental difference I see is that one is restricted to boundary, the other become a free for all. But in both cases, it has to do with available energy for molecules to escape the forces of attraction (cohesion) of the liquid state. It simply takes less energy when external pressure is reduced (at least when reduced to the point of boiling).

Do any of these processes depend on "increased likelihood"? What does that even mean?

Good question. This may be the source of my confusion. I suppose I view a lowering of pressure (at least when to the point that boiling occurs) as reducing the amount of time needed for saturation to occur. My reasoning being that the reaction is no longer restricted to the boundaries. So presumably the transition will occur much more rapidly. Is this wrongheaded?

I can't work out whether this would apply with pressure reductions prior to this threshold of boiling point. Since it's still restricted to boundary interactions, it seems perhaps not.
 
Last edited:
  • #7
JT Smith said:
In a sealed container you're going to end up with saturation no matter the temperature and regardless of the presence of air. The term "saturated air" is sort of confusing in that the air itself is irrelevant with respect to the amount of water vapor present.

Yes, I think Russ helped me work out that I actually mean increased the rate that EVP (saturation) will be reached. The EVP will remain the same regardless of air (dependent only on temperature), but the rate at which EVP will be reached would seem to be increased if a liquid is 'boiling.' Is this still off?

Taking the chance that I'm right, my next question becomes: is this only true at the critical point of boiling, or is there a scaled effect with pressure changes prior to boiling.
 
  • #8
JT Smith said:
In a sealed container you're going to end up with saturation no matter the temperature and regardless of the presence of air.

Only if there is enough water.
 
  • #9
Elquery said:
Taking the chance that I'm right, my next question becomes: is this only true at the critical point of boiling, or is there a scaled effect with pressure changes prior to boiling.

It's generally faster to boil because vaporization isn't constrained to the surface. So now you want to know if the lower the pressure the faster the boil? I wouldn't think so but one can often dream up special cases. For example, if you have a very smooth vessel and minimal contaminants then for a given temperature boiling won't happen when the pressure is just below the boiling point but will at some lower pressure. That's kind of a picky case though. Similarly, you'll never reach saturation if there isn't enough water or enough heat to drive vaporization far enough.
 

Related to Does pressure affect equilibrium vapor pressure (or RH)?

1. Does increasing pressure increase or decrease equilibrium vapor pressure?

Increasing pressure typically increases equilibrium vapor pressure. This is because higher pressure forces more molecules into the gas phase, increasing the number of gas molecules present and therefore increasing the vapor pressure.

2. How does pressure affect relative humidity (RH)?

Pressure does not directly affect relative humidity. However, higher pressure can increase the amount of water vapor that air can hold, so if the amount of water vapor stays the same, the relative humidity will decrease. Conversely, lower pressure can decrease the amount of water vapor that air can hold, leading to an increase in relative humidity.

3. Is there a relationship between pressure and equilibrium vapor pressure?

Yes, there is a direct relationship between pressure and equilibrium vapor pressure. As pressure increases, so does equilibrium vapor pressure. This is because higher pressure means more gas molecules present, which leads to a higher vapor pressure.

4. Can pressure affect the boiling point of a liquid?

Yes, pressure can affect the boiling point of a liquid. When pressure is increased, the boiling point of a liquid also increases because it takes more energy for the liquid molecules to overcome the increased pressure and enter the gas phase. On the other hand, decreasing pressure can lower the boiling point of a liquid.

5. How does pressure impact the phase equilibrium of a substance?

Pressure plays a crucial role in the phase equilibrium of a substance. It can determine whether a substance is in a solid, liquid, or gas phase. For example, at a certain temperature, increasing pressure can cause a substance to transition from a gas to a liquid, while decreasing pressure can cause it to transition from a liquid to a gas. Pressure also affects the equilibrium between the different phases, as seen in the relationship between pressure and equilibrium vapor pressure.

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