# Vapor Pressure vs Boiling Point Confusion

Discussion in 'MCAT Study Question Q&A' started by manohman, Sep 27, 2014.

1. ### manohman 2+ Year Member

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Jul 27, 2014
So Vapor pressure is the result of the phenomenon when the surface molecules of a liquid gain enough kinetic energy to leave the liquid and become a gas. The force per unit area of these molecules is the vapor pressure.

So it makes sense as T goes up, Vapor Pressure should increase since on average more molecules have the energy to break free from their intermolecular bonds.

From an intermolecular bond perspective this makes sense. But form a pressure perspective im still a bit confused. If the boiling point is the point at which the vapor pressure > the atmospheric pressure so all the molecules can keep leaving the surface, then how is it that molecules can leave the surface via evaporation when pAtmosphere > pVapor Pressure? wouldnt the atmospheric pressure push all of the water molecules down into liquid? Just as all of the water molecules push against the atmospheric gas molecules to free themselves when pvapor pressure > patmosphere?

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3. ### 403710Guest 5+ Year Member

718
169
Jun 12, 2011
Brooklyn, NY
I think I understand what you're asking. Thing is, systems that we know are in equilibrium are actually in a "dynamic equilibrium." Meaning there is a constant flow in both ways, but the net change is unnoticeable. That should explain why a cup of water will appear to be about as full a day later, compared to a cup of boiling water which would quickly evaporate.

However, a cup of water nowhere near its boiling point will still eventually completely evaporate in the cases we are familiar with because the kinetic energy of individual molecules within that glass vary greatly. There will always be a very small percentage of molecules that gain the sufficient amount of kinetic energy to break free of the intermolecular bonds. Eventually it will be noticeable as more and more molecules break free. It's similar to the average speed curve of gas molecules, from which the rms is derived. There will be molecules with any kinetic energy, at any temperature.

Evaporation in the case of glasses of water that we are familiar with are spontaneous processes, it's just that the addition of energy to the greatly catalyzes the process.

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4. ### manohman 2+ Year Member

150
5
Jul 27, 2014
Ah I see thank you that makes so much more sense.

So the differnece is that at a liquid's boiling point, all of the molecules have enough energy to escape into the atmosphere (is this why boiling water "rolls"/moves around?

I recall that the bubbels themselves are due to the water coming out of solution since solubility of gases decrease as temperature increases but the liquid itself moving around, is that due to all of the particles now having enough KE to escape? And the Vapor Pressure Atmospheric pressure concept is a way to describe this? Since atmospheric pressure factors into how much energy a particle will need to escape? (more external pressure = need more energy to escape)?

5. ### 403710Guest 5+ Year Member

718
169
Jun 12, 2011
Brooklyn, NY
Np. Happy to help.

Can you clarify a little? I may be a bit tired and maybe didn't read it enough times, but I'm not sure I understand what you're asking.

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6. ### benjaminl1nusProbationary Status 5+ Year Member

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May 12, 2011
I think I understand what you're asking. You're asking why the water is leaving solution and yet there is still water moving around. In a way, yes, it is due to all the particles not having enough KE to escape. The more direct explanation I can think about is surface tension. Imagine 10 people standing in a straight line, all holding hands. a person runs toward that line, and is resisted from moving forward by the people's linked hands. the result is that the line is no longer straight, and curves as a result of the stress of the person running. That's kind of what is going on. Does that help?