ΔS being negative should be unfavourable (as it will raise the value of ΔG). Ah, now I see. It's favourable in terms of Gibbs' energy to increase entropy by expanding. Not sure why this won't hold true for real gases?
Right, unfavorable. Isothermal expansion (or diffusion) of an ideal gas is completely entropic. Therefore the "force" that causes it is an entropic force - the container is prohibiting the gas from maximizing its entropy, i.e., going to a state of low potential energy. This is what causes the pressure of the gas.http://en.wikipedia.org/wiki/Entropic_force
Actually you will notice that the Gibb Free Energy concept includes both an energetic and entropic component, and so it should not be surprising now that there are both energetic and entropic types of forces. Most people are familiar with energetic forces - such as the force which causes objects to fall (gravity) - but not so many are familiar with entropic forces. However molecular diffusion, gas pressure and even the tendency of a stretched rubber band to retract itself are all caused, at least in part, by entropic forces. (And it's not always just one or the other, obviously - the "force" which causes most chemical reactions are a combination of both energetic and entropic forces.
If you want to know more about this topic, you might find the following article interesting:http://johncarlosbaez.wordpress.com/2012/02/01/entropic-forces/
This concept also holds for a real gas - expansion is also primarily entropic. However it is not FULLY entropic. The internal energy of a nonideal gas DOES change during an isothermal expansion. This is because in a nonideal gas there are intermolecular "bondings" that become weaker (usually) as the volume increases. In this case there will be a force that resists entropic expansion, and the magnitude of that force is related to the strength of the intermolecular bonds. Because the pressure in this case is determined through a sum of both the entropic and energetic (enthalpic) forces, we can no longer say that the only force acting on a real, expanding gas is entropy. Even so, the fact that real gasses also expand spontaneously at most temperatures tells us that the outward entropic force exceeds the "inward" enthalpic force in most instances. As you cool the gas down, however, or compress it (reduce volume/increase pressure), the entropic force becomes weaker (or the enthalpic force becomes stronger), until such point as the enthalpic force dominates and the gas will no longer expand spontaneously. This state where the intermolecular forces are greater than the entropic driving force for expansion you may recognize otherwise as a condensed phase, or liquid.
The temperature at which this happens (for a given pressure) is one way to characterize the vaporization/condensation or sublimation/deposition point*. Now if you go the other way and warm the gas up, eventually the gas molecules become so far apart that the enthalpic force is practically zero (at least, in comparison to the entropic force), and thus you approach the ideal gas limit, where expansion is completely dominated by entropy.
Anyway, I think this is a cool alternative way to think about the processes of vaporization and gas expansion. Just some food for thought.
Also, I had another thought. If you have a system in which there are several gases (e.g. A, B, C and D) then does each gas take up a certain volume and a certain partial pressure of its own? By extension, can we say that, within any set of the gases (e.g. A, B and C, or A, C and D, or B, C and D, any set we choose, including just 1 gas or all 4 if we want) the gas laws will apply (e.g. if the gases are ideal, P(total of A, C and D)*V(total of A, C and D)=n(total of A, C and D)*R*T, which could then be true for every set of the gases in the container that we might choose)?
Assuming all the gasses behave ideally, this would be true. We have had this discussion before, and recently.http://www.chemicalforums.com/index.php?topic=67832.0
* I haven't actually tried this calculation to see how close you get to the real condensation point of a gas. It would be interesting to try...