|Aug25-11, 03:29 PM||#1|
Boiling point and vapor pressure
I am trying to understand the concept of boiling from a very fundamental perspective. Most textbooks say that: as a liquid is heated, it's vapor pressure increases. When the vapor pressure reaches the surrounding pressure(or atmospheric pressure at that point), then boiling occurs.
I agree with the above point.
My question is: when the vapor pressure reaches the surrounding pressure, what physically happens that allows vapor to form all through the liquid for it to boil.
|Aug25-11, 03:40 PM||#2|
Pressure is what prevents a volume of water from spontaneously flashing into steam - be it air pressure above or water pressure around it. If a volume of water gets hot enough that it "wants" to be steam at a certain pressure and the pressure around it is lower than that, it will spontaneously flash to steam.
|Aug25-11, 04:00 PM||#3|
Thank you for your reply.
So if the vapor pressure reaches the surrounding atmospheric pressure, how does that lead to the pressure around the volume of water being lower than the pressure with which the steam wants to form?
By "wants", do you mean that the molecules break away from surrounding molecules by absorbing the latent heat?
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