Water Vapor in the Atmosphere

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Main Question or Discussion Point

So I understand generally how humidity in atmospherics work but there's this one thing I can't reconcile with the fundamentals.

First, my understanding of the basics. Water boils at 212 F @ 14.7 psia (1 ATM). At 213F @ 14.7 psia that water is now 1F superheated vapor. So we're all on the same page the boiling point here is also referred to the "bubble point" (i.e. the liquid to saturated liquid-vapor transition line). The dew point as discussed in normal atmospheric terminology is the point at which the water vapor (superheated vapor) starts to condense out (i.e. the vapor to saturated liquid-vapor transition line).

Ok. The thing I can't reconcile is under "normal" atmospheric conditions, how is there superheated water vapor in the air? My first intuition was that as altitude increases the pressure decreases and therefore so does the bubble point, but doing some rough calculations to convert standard atmosphere to a pressure at a given altitude where clouds start to form (~6500 ft), it doesn't seem like the pressure by itself drops enough to lead to evaporation. Additionally, as the altitude increases the temperate decreases so it seems to me the decreasing pressure and temperature cancels each other out, at least to some extent. Lastly, evaporation has to happen at ground level around bodies of water to begin with.

So, how is evaporation even occurring under normal atmospheric conditions (i.e. around 14.7 psia and way, way below 212F).

Below I'm including a ph-diagram of water for reference.

Any help would be greatly appreciated!!

steamtemplines.png
 

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  • #2
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Other general terminology for clarity:

Boiling point = bubble point
Condensation point = dew point = Drybulb temp @ 100% RH = Wetbulb temp @ 100% RH
 
  • #3
russ_watters
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First, my understanding of the basics. Water boils at 212 F @ 14.7 psia (1 ATM). At 213F @ 14.7 psia that water is now 1F superheated vapor....

Ok. The thing I can't reconcile is under "normal" atmospheric conditions, how is there superheated water vapor in the air?
There isn't. As you note, it would have to be 213F for there to be superheated vapor in the air.
The dew point as discussed in normal atmospheric terminology is the point at which the water vapor (superheated vapor) starts to condense out (i.e. the vapor to saturated liquid-vapor transition line).
I don't think you'll find the term "superheated" in the definition of dew point.
So, how is evaporation even occurring under normal atmospheric conditions (i.e. around 14.7 psia and way, way below 212F).
It happens because the vapor pressure is below the saturation pressure, so water diffuses into the air to try to raise the vapor pressure. The actual mechanism is that some individual molecules gain enough energy to break their bond with the water, while the average is much lower. At saturation you have water molecules constantly jumping into and out of each state.
 
  • #4
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There isn't. As you note, it would have to be 213F for there to be superheated vapor in the air.

I don't think you'll find the term "superheated" in the definition of dew point.

It happens because the vapor pressure is below the saturation pressure, so water diffuses into the air to try to raise the vapor pressure. The actual mechanism is that some individual molecules gain enough energy to break their bond with the water, while the average is much lower. At saturation you have water molecules constantly jumping into and out of each state.
To make sure I understand you correctly. The ph-diagram I'm referencing is really just the average and not representative of the whole? Although on average water boils at 212F @14.7 psia, in some circumstances water enters saturation at a much lower temperature condition (e.g. 75F @ 14.7 psia)? If that is correct what outside factors play a part in influencing that behavior? If you haven't inferred from my use of superheat, I'm in the HVAC industry. Would these same principles apply to all fluids (specifically common refrigerants) or is this something unique to water?

Lastly, if in atmospheric terms dew point is not referring to the transition of superheated vapor to saturated liquid-vapor, what is atmospheric dew point referring to specifically?
 
  • #5
russ_watters
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The ph-diagram I'm referencing is really just the average and not representative of the whole?
Yes, though I hope this fact doesn't lead you astray....
Although on average water boils at 212F @14.7 psia, in some circumstances water enters saturation at a much lower temperature condition (e.g. 75F @ 14.7 psia)?
No. Saturation is a bulk property, not a description of an individual molecule. There are lots of normal circumstances where water can be saturated at below 212F -- every temperature has an associated saturation pressure.
If that is correct what outside factors play a part in influencing that behavior?
What, the statistical part? That I'm not sure, but it shouldn't have much relevance to your main question which appears to be a simple definition error on what it means to be "superheated". Again, by definition superheated vapor is above the boiling point at the specified pressure. On a 70F day when the dew point is 60F, that's not superheated vapor, it's just vapor.
If you haven't inferred from my use of superheat, I'm in the HVAC industry.
No....I'm in the HVAC industry, but I'm not sure why that would matter. Superheated vapor (steam) isn't something ever encountered with water in the HVAC industry. Not even with a steam boiler (if functioning properly). It only occurs with refrigerants.
Would these same principles apply to all fluids (specifically common refrigerants) or is this something unique to water?
Yes, the definitions are general definitions.
Lastly, if in atmospheric terms dew point is not referring to the transition of superheated vapor to saturated liquid-vapor, what is atmospheric dew point referring to specifically?
Just plain vapor. I suppose you could say "sub-saturated" vapor, but I don't think that's a typically used term.
 
  • #6
davenn
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So I understand generally how humidity in atmospherics work but there's this one thing I can't reconcile with the fundamentals.
Obviously not by the rest of your post :wink:

You need to read up on some the basics …. start here ……

https://en.wikipedia.org/wiki/Water_vapor

and then on to here …….

https://en.wikipedia.org/wiki/Dew_point

Then ask any specific questions relating to anything there that your are still unsure of.
Quote specific parts and those here at PF can guide you to getting a better handle on the processes :smile:


Dave
 
  • #7
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Evaporation of water is very different from boiling. For evaporation to take place, what matters is the equilibrium vapor pressure of water at the temperature of the liquid (not the total atmospheric pressure). The air is made up of mostly nitrogen and oxygen at a partial pressure (at ground level) of about 100000 Pa. The partial pressure of water vapor in the air is only about 1% of this, or on the order of 1000 Pa. In order for water to evaporate, the equilibrium vapor pressure of water at the liquid temperature must exceed the partial pressure of the water vapor in the bulk of the air. On the other hand, for boiling to occur (a phenomenon that takes place beneath the surface of the liquid water), the equilibrium vapor pressure of water must exceed the entire pressure of the surrounding air, so that bubble can form beneath the surface. So generally, evaporation can take place at a much lower temperature than boiling.
 
  • #8
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@Chestermiller:
I'm sitting here trying to decide if cavitation is better described as 'local boiling,' or 'opportunistic evaporation.' :)
 
  • #9
russ_watters
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@Chestermiller:
I'm sitting here trying to decide if cavitation is better described as 'local boiling,' or 'opportunistic evaporation.' :)
Cavitation is [local] boiling.
 
  • #10
jim hardy
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Ok. The thing I can't reconcile is under "normal" atmospheric conditions, how is there superheated water vapor in the air?
@Jed K

try to relate these kinds of questions to your everyday experiences.

there's a real simple mental trick that'll (i hope) make this one intuitive.
It rests on the definition of "partial pressure". Think back to that.

In my tenth grade chemistry class we were taught that for a given volume of a mix of gases,
the total pressure is the sum of the pressures that each component of the mixture would have if it were the only one present.

So imagine a one foot cube filled with air and water vapor.
If the temperature of the mix is greater than its dewpoint ,
that means the water vapor is above the saturation temperature for its partial pressure. So it's superheated.

If you cool the mix , when it's cooled down to the dewpoint the water vapor is now at its saturation temperature for its partial pressure and further cooling will make dewdrops.
You see that every foggy morning.

If you compress the mix, you raise the pressures of both the air and the water vapor.
At some point the water vapor's partial pressure reaches saturation pressure for whatever is its temperature
so it will begin to condense.
That's why you have to(or at least should) drain the water from your air compressor tank daily.

Any help ?

old jim
 
  • #11
jrmichler
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The relationship between water vapor and air is summarized in a psychrometric chart. Google the term for many good examples, discussions, and references.
 
  • #12
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So, how is evaporation even occurring under normal atmospheric conditions (i.e. around 14.7 psia and way, way below 212F).


View attachment 229686
Vaporization is not restricted to boiling. Anytime that a surface of liquid water is exposed to the free atmosphere, gross vaporization will occur. It is spontaneous and it is continuous. The rate of vaporization depends upon surface temperature, salinity, and a host of other factors, but numbers on the order of 1025 water molecules per square meter per second leaving the water surface to become vapor molecules are common.

Since, at atmospheric pressure, some 99.99% of the space overlying a water body is empty space and the remaining portion is composed of speeding air molecules, atmospheric pressure does not act as much of a "lid". The role of the atmospheric pressure is to assist the surface tension system of bonded water molecules in limiting the escape of water molecules from the water body. It does this by bludgeoning the outside of this surface tension network with a continuous bombardment of air molecules. This bombardment transfers a continuous series of overlapping impulses to the outer surface of the surface network. Each molecule of the bonded surface network experiences some four hundred million such impulses per second. And yet, in that same second, some four million water molecules manage to escape around (and through) that poor surface molecule!
 
  • #13
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Vaporization is not restricted to boiling. Anytime that a surface of liquid water is exposed to the free atmosphere, gross vaporization will occur. It is spontaneous and it is continuous. The rate of vaporization depends upon surface temperature, salinity, and a host of other factors, but numbers on the order of 1025 water molecules per square meter per second leaving the water surface to become vapor molecules are common.

Since, at atmospheric pressure, some 99.99% of the space overlying a water body is empty space and the remaining portion is composed of speeding air molecules, atmospheric pressure does not act as much of a "lid". The role of the atmospheric pressure is to assist the surface tension system of bonded water molecules in limiting the escape of water molecules from the water body. It does this by bludgeoning the outside of this surface tension network with a continuous bombardment of air molecules. This bombardment transfers a continuous series of overlapping impulses to the outer surface of the surface network. Each molecule of the bonded surface network experiences some four hundred million such impulses per second. And yet, in that same second, some four million water molecules manage to escape around (and through) that poor surface molecule!
Please provide a peer reviewed reference for this analysis.
 

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