Vapor Pressure vs Atmospheric Pressure

• Hereformore
In summary, the boiling point is the temperature at which the vapor pressure of a liquid is equal to the atmospheric pressure. This allows bubbles to form within the liquid, pushing the atmosphere back and allowing the liquid to boil. However, when the atmospheric pressure is higher than the vapor pressure, this does not prevent liquid from evaporating at the surface and diffusing into the air. This is due to the partial pressure of the vapor at the interface being higher than in the bulk of the air, allowing it to diffuse away.
Hereformore

Homework Statement

If the boiling point is the point at which vapor pressure > atmospheric pressure, so all of the water molecules can break free and fly into the atmosphere (i.e. overcoming the atmospheric pressure), then why is it that when atmospheric pressure > vapor pressure, the atmospheric pressure is pushing down on the water such that no vapor can escape too?

I guess the question I am asking is: Is vapor pressure independent of external pressure?2. Homework Equations

The Attempt at a Solution

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 I am 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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Also if I am comparing the vapor pressure of a solid vs that of a liquid (sublimation vs evaporation), then if I decrease the temperature of both solid and liquid at the same rate BELOW the melting point of the solid (but at different pressures), apparently the vapor pressure of the liquid will decrease faster than that of the solid, to the extent where Vapor Pressure of the Solid > Vapor PRessure of the Liquid.

What's going on here? I assume its a different case because we're in two differnet phases?

Last edited:
Being strict with the term, the boiling point (T) would verify Atmospheric Pressure = Vapor Pressure. So, at sea level (P = 1 atm), the vapor pressure of water has the value 1 atm (760 mm Hg, or whichever system should we use). And this is known for every component, where P-T graphs are plotted, showing that, for instance, pentane vapor pressure curve reaches those 760 mm Hg at 36.1 ºC (more volatility).
About your question on how is it that molecules can leave the surface via evaporation when Atmospheric Pressure > Vapor Pressure, I supposse it has to do more with statistical mechanics for gases, and it's probably not that simple as the molecules of atmosphere press down the molecules of the liquid (gas), and they wouldn't come up. Since solids are a more stable phase, models would show that the probability that a molecule of the solid comes out of it in a point below its fusion point (sublimation) is low, but non zero. In the end, we would be talking about Quantum Mechanics and its strange world...
To summarize: everything's quiet only at absolute zero.

Hereformore said:
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?

This is not quite correct. At the atmospheric boiling point, the vapor pressure is high enough to allow bubbles to form within the liquid by physically pushing the atmosphere back. However, if the vapor pressure is less than the atmospheric pressure, this does not prevent liquid from evaporating at the upper surface of the liquid, and then having molecules diffuse away into the overlying air. At the interface between the liquid and the air, the partial pressure of the liquid will be equal to the equilibrium vapor pressure of the liquid. The partial pressure of the vapor at the interface will be higher than in the bulk of the air, so the vapor can diffuse away into the air. This is what happens during ordinary evaporation.

wouldnt the atmospheric pressure push all of the water molecules down into liquid?
No. The vapor molecules diffuse into the atmosphere.

Chet

1. What is the difference between vapor pressure and atmospheric pressure?

Vapor pressure is the pressure exerted by the vapor molecules of a substance in a closed container, while atmospheric pressure is the pressure exerted by the weight of the air above a certain area.

2. How are vapor pressure and atmospheric pressure related?

Vapor pressure is directly related to atmospheric pressure. As atmospheric pressure increases, so does vapor pressure. This means that at higher altitudes, where atmospheric pressure is lower, liquids will have a lower vapor pressure and therefore evaporate more quickly.

3. How is vapor pressure affected by temperature?

Vapor pressure increases as temperature increases. This is because higher temperatures cause the molecules of a substance to have more energy and move more quickly, resulting in more molecules being able to escape into the vapor phase.

4. Why is vapor pressure important in everyday life?

Vapor pressure is important in everyday life because it affects the rate at which substances, such as water, evaporate. It also plays a role in determining the boiling point of a substance, and can impact the efficiency of certain processes, such as distillation.

5. How is vapor pressure measured?

Vapor pressure can be measured using a device called a vapor pressure thermometer, which measures the pressure exerted by the vapor of a substance in a closed container. It can also be calculated using the Antoine equation, which takes into account the temperature and properties of the substance.

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