Understand the fundamental reason for the vast difference in properties between gases and the condensed phases (liquids and solids).
We found that gas properties can be understood in terms of a simple ideal gas equation of state or a somewhat more complicated van der Waals equation of state. Remember the causes of deviations from ideality are size and intermolecular forces. However, we did not have to delve very deeply into the types of intermolecular forces and their influences on the gas properties because the gas molecules are, on the average, very far apart. One of the assumptions of the kinetic molecular theory was that the distances between the centers of the gas molecules are much greater than the molecular diameters. In the condensed phases this is not a valid assumption because the average distance between the molecular centers is much smaller in the condensed phases than in the gas phase. The average distance between molecular centers in the gas phase is much greater than in the liquid phase, which is greater than in the solid phase (dgas>>dliquid>dsolid). The answer to the question "So what?" is that because the molecular centers are much closer together in the condensed phases, we must more thoroughly consider the intermolecular forces.
1.) Sketch a diagram for the three phases that clearly demonstrates the different distances between the centers of molecules in the three phases so that a layperson can visualize the three phases on a particle level.
2.) Sketch a diagram that shows an exponential decrease in intermolecular forces with increasing distance between centers of molecules and show approximately on this diagram where the three phases would be.
Web Author: Dr. Leon L. Combs
Copyright ©2000 by Dr. Leon L. Combs - ALL RIGHTS RESERVED