Distillation Theory Part II - Vapor Pressure and Boiling

To explain distillation, it is helpful to consider a bowl of pure liquid.  For this post, we will use a bowl of water at room temperature.  While the liquid may appear to be calm and peaceful, pure bedlam is occurring beneath the surface.

Temperature is simply a measurement of the average speed of the molecules in a sample.  In the example of our room temperature bowl of water, the average speed of the water molecules is 587 meters per second or 1,313 miles per hour.  Now, the water molecules don’t travel long distances, because there are literally billions of billions of water molecules in the bowl.  These molecules are all slamming into each other at this average speed and bouncing off of each other.

Some molecules are moving faster, and some are moving slower.  But the surface of the liquid is where this becomes interesting.  A molecule traveling fast enough can escape the liquid phase and join the gas phase.  Also, a molecule can be at the surface and be struck by a fast moving molecule and catapulted into the gas phase.  This is called evaporation.  The fastest molecules are the ones that escape, thereby lowering the average speed of the liquid (known as temperature).  This is why evaporation is a cooling effect.

What is interesting is an equilibrium will from above a liquid of a fixed temperature.  Every bowl of water at 20°C will have the same proportion of water/air in the vapor phase directly above the water.  The reason for this is, the water molecules that join the gas phase will cool and condense back down into the liquid phase at a fixed rate, while more water molecules will leave the liquid phase and go up into the gas phase.  For 20°C water, the pressure of the water in the gas phase can be measured, and has been found to be 17.54 mmHg.  Atmospheric pressure is about 760 mmHg, therefore the water makes up 2.31% of the gas phase.  The water simply pushes the air aside.

As the water is heated, it takes up more and more of the gas phase above the water.  At 90°C, the vapor pressure of water is 525.8 mmHg.  When you are heating a pot of water to boil, you can see and feel the water amount in the vapor phase increasing and pushing the air aside.

Then, at 100°C, the water starts to boil.  This is the point where the vapor pressure of the water is equal to the atmospheric pressure.  The water can now freely migrate from liquid phase to gas phase.  How quickly all the water is boiled is determined by how much heat is being applied to the vessel.  This is atmospheric distillation/boiling.  In the next post I will explain vacuum distillation.

Post 2 in the theory of distillation series.