Department of Energy Argonne National Laboratory Office of Science NEWTON's Homepage NEWTON's Homepage
NEWTON, Ask A Scientist!
NEWTON Home Page NEWTON Teachers Visit Our Archives Ask A Question How To Ask A Question Question of the Week Our Expert Scientists Volunteer at NEWTON! Frequently Asked Questions Referencing NEWTON About NEWTON About Ask A Scientist Education At Argonne Equilibrium and Evaporation
Name: Jodi
Status: educator
Grade: 9-12
Location: NY
Country: N/A
Date: April 2007

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?



Replies:
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


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


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


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)


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



Click here to return to the General Topics Archives

NEWTON is an electronic community for Science, Math, and Computer Science K-12 Educators, sponsored and operated by Argonne National Laboratory's Educational Programs, Andrew Skipor, Ph.D., Head of Educational Programs.

For assistance with NEWTON contact a System Operator (help@newton.dep.anl.gov), or at Argonne's Educational Programs

NEWTON AND ASK A SCIENTIST
Educational Programs
Building 360
9700 S. Cass Ave.
Argonne, Illinois
60439-4845, USA
Update: June 2012
Weclome To Newton

Argonne National Laboratory