The Atmosphere Is Not A Water Vapor Hotel
by Eric Hall
March 2, 2013
Most of us have probably experienced a warm, humid day in the summer. I imagine some of us in the northern hemisphere are right now wishing for one of those days; a day with the air thick where the sweat doesn't evaporate easily and walking into a building dried and cooled by air conditioning immediately gives rise to goosebumps on the skin. We might hear someone, perhaps even a meteorologist, describe the warm air as "holding more moisture." However, this is a disservice to the true science happening in the atmosphere.
Gases in the atmosphere are not dependent on one another for their content amounts in the atmosphere. At normal atmospheric pressures, the gases behave very well to what is known as the ideal gas law. This is a very good approximation when dealing with pressures around 1 atmosphere, because the inter-molecular interactions can largely be ignored. Derived from the ideal gas law is what is known as Dalton's Law or Dalton's law of partial pressures which states, "the total pressure exerted by the mixture of non-reactive gases is equal to the sum of the partial pressures of individual gases." This means the amount of water vapor has nothing to do with what the other gases in the atmosphere are doing, but instead it adds to the pressure of the atmosphere in that area.
Craig F. Bohren writes about this concept in his book Clouds In A Glass of Beer: Simple Experiments in Atmospheric Physics. He proposes a hypothetical scenario of a bottle partially filled with water. If the water in the air above the water is removed and the bottle sealed, the space above the water will not remain free of water molecules. As molecules escape the surface of the water, there will come a point where dynamic equilibrium is reached; which is when the evaporation rate equals the condensation rate. Professor Bohren continues:
When equilibrium is reached, the partial pressure of the water vapor (small compared with the total pressure, the sum of partial pressures contributed by each of the constituents of air) is called the saturation vapor pressure. But equilibrium vapor pressure would be a better term: "saturation" evokes, incorrectly, the image of a sponge. There is no end to the blather about the "holding power of air" and how air can "hold" more water vapor at high temperatures than at low temperatures; this implies that in air there is only so much space -- like rooms in a hotel -- between air molecules, and when filled with water molecules the air is saturated, just like the pores of a sponge. But air doesn't "hold" water vapor -- it coexists with it. Indeed, the presence of air (oxygen, nitrogen, etc.) in the space above the liquid is largely immaterial: the pressure of a vapor in equilibrium with its liquid would be nearly the same with or without air. It is worth noting here that everything has a vapor pressure; that of mercury at room temperature, for example, is about one-thousandth that of the atmosphere at sea level. That of most solids, especially at room temperature, is very much lower -- but is not zero. Your skin has a vapor pressure. Fortunately, it is rather low or you wouldn't be here to read this -- you would have evaporated away long ago.Another good thought experiment comes from Steven M. Babin at the University of Maryland. He writes:
Imagine a closed container containing a beaker of pure water and a beaker of ocean water. Place the two solutions side by side so that they are at the same atmospheric temperature and pressure. The air above these two solutions is at the same temperature and pressure. If air "holds" water vapor, then the two solutions should have the same saturation vapor pressure. However, the saturation vapor pressure above the saline solution is less than that above the pure water. In the saline solution, the salt ions replace some of the water molecules so that fewer water molecules are available for evaporation. Therefore, the presence of the salt reduces the rate of evaporation from the saline solution compared to the solution of pure water. This then is the reason why the saturation vapor pressure above the saline solution is less than that above pure water. Note that the presence of initially identical air above the solutions could not account for this difference.In essence, the amount of water vapor has to do with the properties of the liquid from which the vapor is being evaporated, not the properties of the atmosphere. The atmosphere is largely empty space, thus any new vapors are largely unaffected.
The Bad Clouds FAQ adds another great explanation debunking the claim that the atmosphere acts like a sponge for water vapor. In answering a question about cloud formation and the idea that "cold air cannot hold as much water vapor," the author, Alistair B. Fraser, writes:
...watch the formation of cumulus clouds. These are the puffy white clouds which form on a summer's day over the Sun-warmed ground. The clouds form at the top of rising columns of air. As the air rises to a region of lower pressure, its density drops (the molecules get farther apart) and yet that is where the cloud forms. If you were to believe the silly explanation of the water being squeezed out because the molecules were getting closer together, then you should also expect that the clouds would have formed, not in the rising air, but in the sinking air, because it is there that the air density is increasing.The Bad Cloud FAQ uses a great analogy which is probably the most important lesson in this lesson. As they state, just because one gets the right answer reducing the fraction 16/64 by cancelling the 6's does not mean you will get the correct answer in all cases. In the case of water vapor, the idea of the holding capacity of the air does largely describe the behavior, but is not correct and will not describe some cases. We often hear the term relative humidity, and at the surface it rarely exceeds 100%. But higher in the atmosphere, the air can often be "super-saturated," which simply means the rate of condensation is greater than the rate of evaporation.
Amazingly, this bad science can spread and be very difficult to scrub out before its spread continues. For example, USA Today's meteorologist continues to spread this falsehood under their explanation of dew point. Although he uses quotes, his statement is:
When the air can no longer "hold" all of the water vapor in it, the vapor begins condensing into ordinary, liquid water. If the air is in the sky, the vapor condenses into cloud drops. If the air is right above the ground the vapor condenses to make fog (which is nothing but a cloud that's on the ground.) If the air is touching things such as grass or a car's windshield, it will condense to make dew.Although it simplifies the concept, it is not correct. I see no reason to simplify it when the real science is actually more fascinating, and is good science. I hear meteorologists use the term "holding water" on occasion as well, and it always makes me cringe a little. I have even caught this explanation in a few science textbooks, which makes keeping the science correct an even more difficult task.
While I understand the need to reduce explanations of science down to their essence when explaining a concept to someone who is not an expert in a field, it should never be at the expense of the science itself. Whether it means spending a few extra minutes explaining a concept, finding new models or examples, or just some hard work getting up to speed to understand a concept, we should never sacrifice the truth for the sake of simplification. While it may give the correct answer some of the time, it can lead to misunderstandings when faced with other concepts, and can actually damage the reputation of science. I for one will always strive to be concise, but correct as well. I don't ever want my explanations to be just "good enough," or just enough to "hold water."
by Eric Hall
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