MK - 4-11-2005 at 14:57
Hey, I could use a little help regarding how vapor pressure and boiling point change under different circumstances, (hence the title.) My textbook
does not explain the subject well, and so I could use your insights. Here are some questions I need to answer, as well as my answer to the questions
and my rationale. If you could tell me if my deductions are correct, and correct them if they arn't, I'd be very grateful-
* What would happen to the vapor pressure and boiling point of a liquid if the following changes took place?
1. External pressure decreased, temp constant
My Answer: Boiling point down since pressure and boiling point are directly proportional, no effect on vapor pressure (b/c no mention in the book was
made of outside pressure affecting vapor pressure.)
2. Temperature decreases by 20C, pressure constant
A: No effect on boiling point (for obvious reasons,) vapor pressure down because vapor pressure is directly proportional to heat
3. Liquid area is increased, temp and pressure constant
A: No effect on boiling point since size doesn't matter... no clue on the pressure one. Does more surface area mean more vapor pressure?
4. The liquid is moved to a higher elevation, temp constant
A: Easy. Boiling point goes down as well as pressure.
5. The amount of liquid is decreased, constant temp and pressure.
A: To be honest, I would put "no change" as my answer. Unless of course vapor pressure changes with quantity of liquid... does it?
Thanks for your help. I look forward to receving your answers, but I'm honestly even more interested in getting a good explanation of how
vapor pressure and boiling point work.
-MK
12AX7 - 4-11-2005 at 15:52
AFAIK, vapor pressure is a physical quantity of the substance, based on the equilibrium between molecules clinging together to form a liquid phase vs.
molecules seperating (by kinetic energy) to form a gaseous phase (usually dissolved in another gas as a partial pressure, but not necessary as pure
vapor can also exist alone at all temperatures). As a result, if a substance is left open such that the molecules escape, never to return, the
equilibrium is shifted and it will evaporate. In an enclosed space, the two forces counter and it establishes a partial pressure; the time it takes
to establish this will be an exponential curve depending on volume, temperature and various characteristics (partial pressure, etc.). So yes, I would
agree, #3 and #5 are almost trick questions.
Tim
Magpie - 4-11-2005 at 20:16
Vapor pressure vs temperature curves are not proportional but are generally exponential. Take a look at some curves in a handbook. Otherwise your
answers are generally in the right direction.
#3 - area has no effect on pressure
#5 - vp does not change with area
The Clausius - Clapeyron equation is an idealization of the vapor pressure temperature relation:
ln (p2/p1) = (lamda/R)(1/T1-1/T2)
where lamda = molal heat of vaporization
and R = ideal gas law constant
When the vapor pressure equals the external pressure boiling occurs.
[Edited on 5-11-2005 by Magpie]
[Edited on 5-11-2005 by Magpie]
MK - 6-11-2005 at 16:23
Thanks for your answers, verifies what I had supposed. It isn't a terribly complicated matter, but it is easy to overthink things and arrive at
the wrong conclusion. Thanks again guys
-MK