[Skiznibbler]

I read that vapor pressure is the amount of pressure in the air above a liquid when the vapour is in dynamic equilibrium with its liquid counterpart. For example when theres a vapor pressure the vapors are condensing at the same speed that the liquid is evaporating so the vapor stays the same. Is this what vapor pressure is?

I know that increasing the temperature will increase the vapor pressure. Does a higher vapor pressure just mean theres a higher concentration of vapor present therefore a larger amount of vapors condensing and liquid evaporating simultaneously?

Also when the air is said to be saturated or at dew point does that mean no more vapors can be dissolved in the air or does it just mean that the liquid and vapors have reached equilibrium and the vapor pressure will remain the same unless the temperature is increased.

One other quick question. When a liquid reaches it's boiling point does that mean it cannot take any more heat and will remain the same temperature? For example when water starts boiling does the boiling water remain at 100C or is it only the vapors that are guarenteed to be 100C?

[Yggdrasil]

Yes, vapor pressure represents a dynamic equilibrium.  In fact, you can think of vapor pressure as the equilibrium constant for a chemical reaction.  For example, for water you have the "reaction":

H₂O _(l) <--> H₂O _(g)

Where the K_p for the reaction is the vapor pressure.

A higher vapor pressure does mean a higher concentration of vapor present at equilibrium.  Increasing the temperature increases the rate of evaporation.  To counteract this increased rate of evaporation, a higher rate of condensation is needed.  Since the rate of condensation depends on the concentration of vapor, the system achieves a higher rate of condensation by increasing the concentration of vapor present at equilibrium.

When something is saturated with vapor, it means that the partial pressure of the gas is equal to its vapor pressure.

Like any first-order phase transition, vaporization occurs at a fixed temperature.  Any heat that is put into the system is used to break the intermolecular bonds that hold the molecules together in the liquid so that it a molecule can break free into the gas phase.  The energy needed to break these bonds is known as latent heat or the enthalpy of vaporization.

[Skiznibbler]

Thanks alot. That explains it. I have another quick question though.

How does the vapor pressure interplay with the atmospheric pressure? Lets say I heat some water to 50C. The water starts evaporating quicker and more steam enters the air. I know that for the vapor pressure to be equal to atmospheric pressure I have to heat the water to 100C. Obviously when I raise the temperature from 50C to 100C the whole room doesn't just fill up with steam so theres exactly the same amount of H2O molecules as there are N and O molecules in the air? Have I got this mixed up?

I know that pressure is measured by the amount of molecules hitting a surface so the atmospheric pressure exerted on the water is basically nitrogen and oxygen atoms hitting the surface of the water. How can the same amount of vaporized H2O molecules be hitting the surface of the water unless the room has filled with H2O molecules? What I mean is the vapor doesn't just hover above the water when it evaporates instead it travels around the room. How can water vapor exert the same amount of pressure on the liquids surface as the air just by heating the liquid to boiling point? Does the water vapor literally take the place of the air just above the liquid and exert its own pressure in place of the regular airs pressure?

[Yggdrasil]

Vapor pressure refers to only the partial pressure of the gaseous form of your liquid.  In a closed system at equilibrium (e.g. if you seal a beaker containing water), the partial pressure of water vapor above the liquid will be the vapor pressure of water regardless of the partial pressures of other gases in the air (e.g. N₂, O₂).

The boiling point of a liquid, however, does depend on atmospheric pressure.  Why is this?  Well, as I mentioned before, the vapor pressure is essentially the K_p for the reaction H₂O _(l) <--> H₂O _(g).  If this is the case, why does vaporization take place only at the surface of water?  For other chemical reactions, the reaction will take place all throughout the solution and not only at the surface.

The answer has to do with pressure.  When the vapor pressure of a gas exceeds the atmospheric pressure, the vaporization of water does occur all throughout the solution (not entirely true, more on this later), creating a phenomenon known as boiling.  Boiling happens only when vapor pressure exceeds atmospheric pressure because vaporization in the middle of the solution requires the formation of bubbles that have an internal pressure large enough to counteract the external pressures that try to squeeze the bubble back together (roughly atmospheric pressure).  Since the interior of the bubble is entirely vapor, the highest pressure inside the bubble is the vapor pressure of the liquid.  So, when the vapor pressure is smaller than atmospheric temperature, any bubbles are immediately crushed back into solution by atmospheric pressure and evaporation can occur only at the surface of the liquid.  However, once the vapor pressure of the liquid exceeds atmospheric pressure, bubbles can start forming and the liquid boils.

This explains why the boiling point of a liquid is dependent on atmospheric pressure.  The boiling point of the liquid does not change because a change in atmospheric temperature affects the vapor pressure; rather, it is the boiling point itself that depends on atmospheric temperature.  Vapor pressure, as I mentioned before, is entirely independent of atmospheric pressure.

(as an aside, there is some pretty cool physics/chemistry involved in the formation of bubbles.  As it turns out, the limiting step of bubble formation is whether the internal pressure of the bubbles can overcome the surface tension of the water [i.e. the intermolecular bonds that hold water molecules together].  Interestingly, surface tension forces are largest for infinitesimally small bubbles and decrease as the size of the bubble increases.  Most of the time, bubbles need to be nucleated; bubbles need to form along a surface where a scratch can help disrupt the surface tension of water and allow a larger bubble to form.  One consequence of this is the ability to create superheated solutions; heating water in a very smooth container will allow water to be above its boiling point because the vapor molecules will not yet have enough energy to overcome the surface tension of water and there are no sites for nucleation.  Disrupting the surface tension of water, for example, by stirring the solution, nucleates bubbles, often rapidly and violently, causing an explosion of vapor.  This is why you need to add boiling chips to any solution that you heat in laboratory glassware.  Similar phenomena involving surface tension and nucleation are important in crystallization [and help lead to the fractal patterns you see in snowflakes])