[Needaask]

What I understood from vapour pressure is that its the pressure exerted on the walls and surface of the liquid when the equilibrium of gas leaving and entering the liquid is the same.

However, would the vapour pressure be of a liquid in a sealed room with a vacuum above it? Or would the vapour pressure remain the same even if I have an atmosphere above it? I read this http://answers.yahoo.com/question/index?qid=20090309173048AAEAr1C and it suggested that at higher external pressure the vapour pressure remains the same. But if the external pressure is greater, wouldn't the water molecules be unable to escape so how does the vapour pressure still remain the same?

Also isn't the vapour pressure the pressure exerted on the walls and liquid surface, so why would my vapour pressure have to be equal to the atmospheric pressure for boiling to occur? I don't quite understand how vapour pressure is linked to the boiling point.

so I'm quite confused about this. Thanks for the help

[Corribus]

Vapor pressure is a quantity specific to the substance and relates to the (temperature dependent) equilibrium constant for vaporization.  It is the pressure of the vapor above a liquid when equilibrium has been reached between molecules escaping the surface of the liquid and molecules in vapor phase returning to the liquid.  It is something akin to a "hypothetical" value, because it assumes the system is closed.  In practice (open containers) equilibrium isn't really ever reached because the product (the vapor) is constantly removed from the system (diffusion into the atmosphere), but the vapor pressure still has a meaning with respect to boiling because bubbles cannot form until this pressure overcomes the atmospheric pressure.  Until this happens, vaporization is strictly a surface phenomenon.

Maybe this will help:

http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/vappre.html#c1

Boiling is actually a fairly complicated process at a microscopic level which involves a lot of fluid physics (cavitation, bubble formation, static fluid pressure, etc).  I'd be lying if I said I understood every aspect of it.

[Needaask]

Vapor pressure is a quantity specific to the substance and relates to the (temperature dependent) equilibrium constant for vaporization.  It is the pressure of the vapor above a liquid when equilibrium has been reached between molecules escaping the surface of the liquid and molecules in vapor phase returning to the liquid.  It is something akin to a "hypothetical" value, because it assumes the system is closed.  In practice (open containers) equilibrium isn't really ever reached because the product (the vapor) is constantly removed from the system (diffusion into the atmosphere), but the vapor pressure still has a meaning with respect to boiling because bubbles cannot form until this pressure overcomes the atmospheric pressure.  Until this happens, vaporization is strictly a surface phenomenon.

Maybe this will help:

http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/vappre.html#c1

Boiling is actually a fairly complicated process at a microscopic level which involves a lot of fluid physics (cavitation, bubble formation, static fluid pressure, etc).  I'd be lying if I said I understood every aspect of it.

Hi thanks for the reply.

So vapour pressure is the pressure exerted when equilibrium is met. But I still can't y understand how it affects the boiling point. I've managed to compile my questions

1. How does the higher vapour pressure explain a greater boiling point since vapour pressure is only when there is a closed system with a vacuum above it only?

2.The vapour pressure exerts on the liquid surface too, so how does it help to "push" the external pressure away? And so how does it enable boiling?

I'm still quite confused because I can't see the relationship between the pressure exerted on the liquid surface and the boiling point. It makes sense to just say when we are at a lower lying area the gas would have to do more work against the atmosphere because there is a higher external pressure. But now this "work against atmosphere" isn't the vapour pressure.

Thanks for the help

[delta609]

1.  Higher vapor pressure does not equal higher boiling point. It's the opposite.

2.  If the vapor pressure (pressure within the liquid you are trying to boil) is greater than the atmospheric pressure then there is nothing stopping molecules with sufficient kinetic energy from leaving the liquid, so they do.

[Needaask]

1.  Higher vapor pressure does not equal higher boiling point. It's the opposite.

2.  If the vapor pressure (pressure within the liquid you are trying to boil) is greater than the atmospheric pressure then there is nothing stopping molecules with sufficient kinetic energy from leaving the liquid, so they do.

Oops you're right for 1.

But for 2, I don't get the term vapour pressure here. How does the vapour pressure which exerts over the walls and surface of the liquid play a role? I'm thinking that the atmosphere will push the water so they can't vapourize so how can the vapour pressure even increase? Won't they remain as liquid except for the few molecules that have enough KE to escape.

[delta609]

You can look at vapor pressure like this. . .

Vapor pressure and the kinetic energy within a liquid go together. One doesn't exist without the other. The more kinetic energy within a particular liquid, the higher its vapor pressure due to the higher speed of the molecules.  Can you see where I'm going with this? When this kinetic energy (directly related to vapor pressure) is greater than the pressure exerted by the atmosphere above it on the surface, it can push through. 

If you don't already know, heat is kinetic energy.  That's why heating liquids allows for the kinetic energy (which increases vapor pressure) to increase and overcome the atmospheric pressure; hence liquids boil with added heat. 

Another way to look at it is picture a vacuum with zero atmospheric pressure.  With this abscence of atmospheric pressure pressing down on a substance, it will readily boil away due to the fact the vapor pressure within the substance in the vacuum has no other pressure to compete with. 

Hope this helps.

[Needaask]

You can look at vapor pressure like this. . .

Vapor pressure and the kinetic energy within a liquid go together. One doesn't exist without the other. The more kinetic energy within a particular liquid, the higher its vapor pressure due to the higher speed of the molecules.  Can you see where I'm going with this? When this kinetic energy (directly related to vapor pressure) is greater than the pressure exerted by the atmosphere above it on the surface, it can push through. 

If you don't already know, heat is kinetic energy.  That's why heating liquids allows for the kinetic energy (which increases vapor pressure) to increase and overcome the atmospheric pressure; hence liquids boil with added heat. 

Another way to look at it is picture a vacuum with zero atmospheric pressure.  With this abscence of atmospheric pressure pressing down on a substance, it will readily boil away due to the fact the vapor pressure within the substance in the vacuum has no other pressure to compete with. 

Hope this helps.

Hi thanks

I get the explanation using kinetic energy. But why can I look at it this way? Because this is what I'm thinking: when I have a liquid the atmosphere will push the liquid surface so very few molecules have enough energy to leave. However when I increase the temperature the kinetic energy of the molecules increase so they can do work against the attractive forces as well as the atmosphere.

So from this explanation I don't see where vapour pressure comes in at all. Why would there still be a pressure acting on the liquid surface by the vapour of the liquid itself? Since only very few molecules can vaporize unlike in a sealed container with no atmosphere?

Thanks

[Borek]

Because this is what I'm thinking: when I have a liquid the atmosphere will push the liquid surface so very few molecules have enough energy to leave.

It doesn't work this way. Presence of other gases doesn't change the vapor pressure over a substance, just like it doesn't change the energy of the substance molecules (their energy is a function of their temperature). As they have the same energy, they can leave the surface the same way, regardless of what is above.

[Needaask]

Because this is what I'm thinking: when I have a liquid the atmosphere will push the liquid surface so very few molecules have enough energy to leave.

It doesn't work this way. Presence of other gases doesn't change the vapor pressure over a substance, just like it doesn't change the energy of the substance molecules (their energy is a function of their temperature). As they have the same energy, they can leave the surface the same way, regardless of what is above.

Oh so they will still have the same vapour pressure despite having an atmosphere above it. But so what is the significance of this vapour pressure having to be equal to the external pressure of the atmosphere (or any gas above it)?

Because what I understood from explanations here and from Khan Academy is that the vapour pressure pushes away the atmosphere?

But since vapour pressure is just the pressure exerted on the walls and liquid surface I can't seem to relate them together for this.

Thanks for the help everyone hope to get a clearer insight from you guys too

[Borek]

When the vapor pressure equals atmospheric pressure it is possible for the evaporation to start inside the liquid, not only on the surface. That's how the boiling starts (see the definition of boiling).

[Needaask]

When the vapor pressure equals atmospheric pressure it is possible for the evaporation to start inside the liquid, not only on the surface. That's how the boiling starts (see the definition of boiling).

Hi Borek. I googled boiling point and the relationship is "that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid"

But I don't quite understand how vapour pressure is related here. Isn't vapour pressure just the pressure exerted at the surrounding such as the walls and liquid surface when an equilibrium between how much water leaves and enters? So I don't understand why there's a relationship between the two of them.

Thanks

[Borek]

Isn't vapour pressure just the pressure exerted at the surrounding such as the walls and liquid surface

You don't need surroundings for the pressure to exist. Presence of the gas is enough.

when an equilibrium between how much water leaves and enters?

And there is no need for equilibrium for the pressure to exist. Equilibrium that you are mentioning refers to the saturated vapor.

[Needaask]

Isn't vapour pressure just the pressure exerted at the surrounding such as the walls and liquid surface

You don't need surroundings for the pressure to exist. Presence of the gas is enough.

when an equilibrium between how much water leaves and enters?

And there is no need for equilibrium for the pressure to exist. Equilibrium that you are mentioning refers to the saturated vapor.

Hi thanks again for the help

Why isn't the surroundings required, because isn't the pressure exerted on walls?

Oh I didn't know that, thanks for pointing that out. But still in this case why does the vapour pressure have to be equal to the external pressure for boiling to take place? I still can't see the link for it..

Thanks for the help

[Borek]

why does the vapour pressure have to be equal to the external pressure for boiling to take place?

When the vapor pressure is below the atmospheric, vapor is unable to overcome the pressure and create a bubble inside of the solution. For that it must at least equal the atmospheric pressure.

[Needaask]

why does the vapour pressure have to be equal to the external pressure for boiling to take place?

When the vapor pressure is below the atmospheric, vapor is unable to overcome the pressure and create a bubble inside of the solution. For that it must at least equal the atmospheric pressure.

What do you mean by overcome? Because as you mentioned the presences of other gases doesn't change the vapour pressure. So why would the vapour pressure be the one doing work against the atmosphere?

Thanks