Derivation of the Van der Waals equation

As a specific example of the application of perturbation theory, we consider the Van der Waals equation of state. Let be given by a pair potential:

with

This potential is known as the hard sphere potential. In the low-density limit, the radial distribution function can be shown to be given correctly by or

is taken to be some arbitrary attractive potential, whose specific form is not particularly important. Then, the full potential might look like:

  
Figure 1:

Now, the first term in is

where

is a number that depends on and the specific form of .

Since the potential is a hard sphere potential, can be determined analytically. If were 0, then would describe an ideal gas and

However, because two particles may not approach each other closer than a distance between their centers, there is some excluded volume:

  
Figure 2:

If we consider two hard spheres at closest contact and draw the smallest imaginary sphere that contains both particles, then we find this latter sphere has a radius :

  
Figure 3:

Hence the excluded volume for these two particles is

and hence the excluded volume per particle is just half of this:

Therefore Nb is the total excluded volume, and we find that, in the low density limit, the partition function is given approximately by

Thus, the free energy is

If we now use this free energy to compute the pressure from

we find that

This is the well know Van der Waals equation of state. In the very low density limit, we may assume that , hence

Thus,

from which we can approximate the second virial coefficient:

A plot of the isotherms of the Van der Waals equation of state is shown below:

  
Figure 4:

The red and blue isotherms appear similar to those of an ideal gas, i.e., there is a monotonic decrease of pressure with increasing volume. The black isotherm exhibits an unusual feature not present in any of the ideal-gas isotherms - a small region where the curve is essentially horizontal (flat) with no curvature. At this point, there is no change in pressure as the volume changes. Below this isotherm, the Van der Waals starts to exhibit unphysical behavior. The green isotherm has a region where the pressure decreases with decreasing volume, behavior that is not expected on physical grounds. What is observed experimentally, in fact, is that at a certain pressure, there is a dramatic, discontinuous change in the volume. This dramatic jump in volume signifies that a phase transition has occurred, in this case, a change from a gaseous to a liquid state. The dotted line shows this jump in the volume. Thus, the small flat neighborhood along the black isotherm becomes larger on isotherms below this one. The black isotherm just represents a boundary between those isotherms along which no such phase transition occurs and those that exhibit phase transitions in the form of discontinuous changes in the volume. For this reason, the black isotherm is called the critical isotherm, and the point at which the isotherm is flat and has zero curvature is called a critical point.

A critical point is a point at which

Using these two conditions, we can solve for the critical volume and critical temperature :

and the critical pressure is therefore

Using these values for the critical pressure, temperature and volume, we can show that the isothermal compressibility, given by

diverges as the critical point is approached. To see this, note that

Thus,

It is observed that at a critical point, diverges, generally, as . To determine the heat capacity, note that

so that

Then, since

it follows that

The heat capacity is observed to diverge as . Exponents such as and are known as critial exponents.

Finally, one other exponent we can easily determine is related to how the pressure depends on density near the critical point. The exponent is called , and it is observed that

What does our theory predict for ? To determine we expand the equation of state about the critical density and temperature:

The second and third terms vanish by the conditions of the critical point. The third derivative term can be worked out straightforwardly and does not vanish. Rather

Thus, we see that, by the above expansion, .

The behavior of these quantities near the critical temperature determine three critical exponents. To summarize the results, the Van der Waals theory predicts that

The determination of critical exponents such as these is an active area in statistical mechanical research. The reason for this is that such exponents can be grouped into universality classes - groups of systems with exactly the same sets of critical exponents. The existence of universality classes means that very different kinds of systems exhibit essentially the same behavior at a critical point, a fact that makes the characterization of phase transitions via critical exponents quite general. The values obtained above for , and are known as the mean-field exponents and shows that the Van der Waals theory is really a mean field theory. These exponents do not agree terribly well with experimental values ( , ). However, the simplicty of mean-field theory and its ability to give, at least, qualitative results, makes it, nevertheless, useful. To illustrate universality classes, it can be shown that, within mean field theory, the Van der Waals gas/liquid and a magnetic system composed of spins at particular lattice sites, which composes the so called Ising model, have exactly the same mean field theory exponents, despite the completely different nature of these two systems.

Another problem with the Van der Waals theory that is readily apparent is the fact that it predicts for certain values of the density . Such behavior is unstable. Possible ways of improving the approximations used are the following:

1.

Improve the approximation to .

2.

Choose a better zeroth order potential .

3.

Go to a higher order in perturbation theory.

Barker and Henderson have shown that going to second order in perturbation theory (see discussion in McQuarrie's book, chapter 14), yields well converged results for a square well fluid, compared to ``exact'' results from a molecular dynamics simulation.