Equilibrium and Evaporation /h3>

name         Jodi
status       educator
grade        9-12
location     NY

Question -   We were talking about equilibrium and how the rate
of reactions need to be the same.  How does water evaporate at
room temperature, for example from an open bottle of spring water?
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The molecules of water at the surface of the water like being around other
molecules, but they are jostling around a lot, and can sometimes 'jump' into
the air. This is basic mechanism for the evaporation you see. Likewise, any
water vapor in the air can hit the surface, and 'stick'. You see this
process when 'sweat' forms on the outside of a cold drink on a humid day. In
reality, both processes are always occurring, but one typically occurs
faster than the other (evaporation when air is dry or liquid is hot, or
condensation when air is damp or liquid is cold).

The ability or willingness of molecules in the liquid to jump into the air
is called its 'vapor pressure'. It is called the vapor pressure because if
you add up all the molecules 'jumping' into the air, they exert a pressure
(just like any other gas). A liquid like acetone (nail polish remover) has a
very high vapor pressure, in other words more of its molecules like to jump
into the gas phase. A liquid like olive oil has a very low vapor pressure
because its molecules do not like to jump into the air.

If the air is dry enough, more molecules will jump from your water bottle
into the air than will stick from the air into the water. Over time, the
water will continue to lose molecules to the air, and eventually all will be
gone.

As an interesting side note, when you cool a liquid, its vapor pressure
tends to go down. When you heat the liquid, the vapor pressure tends to
rise. When the vapor pressure of the liquid equals atmospheric pressure, the
liquid boils. The definition of boiling point is when the vapor pressure
equals atmospheric pressure. This is why water boils at a cooler temperature
at high altitude -- at high altitude, atmospheric pressure is lower, and so
a lower temperature is required to raise the water's vapor pressure to match
it.

Hope this helps,

Burr Zimmerman
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Hi, good question.

Since the container is not closed, the
water vapor can escape and the
system will not come to equilibrium.
In a dry room, the water will continue to
evaporate at a more or less constant
rate until all of the water is gone.

Le Chatelier's principle says that if
you continuously remove the product of a
reaction, you will continue to shift towards
more and more product formation - so this example is
in accord with this principle.

If the container were closed, eventually the rate of
evaporation would equal the rate of condensation and
an equilibrium would be established.

best, Dr. Topper
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At room temperature, some fraction of water molecules will have 
enough energy to break free from the liquid and become gas 
(evaporate). Some fraction of water molecules in gas phase 
will also be captured in liquid (condense). Evaporation and 
condensation are occurring simultaneously on the water surface. 
Assuming that the air in the room is not already saturated with 
water vapor (for that given temperature), the rate of evaporation 
will exceed the rate of condensation, and you get net evaporation.  
In a sealed system (e.g. a closed water bottle), once allowed enough 
time, there enough water vapor in the air that evaporation and 
condensation rates are equalized, and equilibrium is reached.

Don Yee
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Jodi,
 
Look up a graph of "Maxwell's Distribution of Molecular Speed". 
Essentially, this shows that the temperature of any system with a 
statistical number of particles is the result of the random motion 
of the particles in the system, and that the range of possible 
speeds of these particles is broad and dynamic.
 
This graph suggests two important points pertinent to your question: 
(1) temperature is the average of all the range of particle motions, and 
(2) at any given temperature, there will be a finite (and constant) 
number of particles that have enough speed to escape the liquid phase.

Let's say (for the sake of this discussion that the graph shown in 
this site:

http://en.wikipedia.org/wiki/Image:MaxwellBoltzmann.gif

Is *not* representing different gases at the same temperature, but 
rather one particular liquid at different temperatures. You can then 
imagine that as the temperature of the liquid increases, the 
probability density skews to the right (to higher speeds). Now, let 
us focus on the blue and yellow lines. Let us say that the yellow 
line represents the liquid in thermal equilibrium with the room 
(is now at room temperature). Further, let us say that 1000 m/s 
is the required speed in order for a liquid particle to escape 
the liquid phase and enter the gas phase. If this is so, then if 
the temperature is the yellow line, then there will be a small 
amount of particles that will have that requisite velocity. Let 
us say that some of these particles do escape - if so, then the 
average temperature will drop. If the average temperature of the 
liquid drops (due to the escape of the fast moving particles) and 
the liquid is now described by the blue line (with a lower maximum 
since the number of particles decreased), then this liquid is now 
able to have heat enter it since it will be of a lower temperature 
than the environment. This means that in absorbing heat from the 
environment, rising in temperature (equilibrating once again with 
the room), the blue line becomes the yellow line again. And again, 
there will be particles that are able to escape the liquid.

I am sure you can now see how a glass of water left on a kitchen 
counter will evaporate and dry out - despite the fact that the 
water is nowhere near its boiling point.

Greg (Roberto Gregorius)
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Every pure substance can coexist with another phase of the same 
composition. A "phase" of a pure substance is the arrangement of 
molecules comprising the substance. The common phases are: Gas, 
Liquid, and Solid. There can be exceptions, but to keep it simple 
a substance has only one gas phase, one liquid phase, and one solid 
phase. The relative amount of each phase depends on the substance, 
the temperature, and the pressure -- here we are ignoring mixtures.
In particular water at a temperature of 25 C. has an equilibrium 
vapor pressure of about 25 mm of Hg. That makes a good approximation 
to remember since the "magic" number is "25". If the liquid is in an 
open container some of the molecules in the gas phase can diffuse 
(wander away) or be moved by air movement near the water's surface. 
So eventually all the water will evaporate from an open container.
Your discussion of how the forward and reverse rates of a process or 
reaction is the same has a hidden constraint that frequently is not 
clearly pointed  out -- the system (gas + liquid) is in a CLOSED 
container. It does not apply to an open container, which is able 
to "exchange" both heat and molecules with its surroundings.

Vince Calder
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