I'm learning about Le Châtelier's Principle and it makes sense to me, generally. But I'm having a hard time understanding the pressure aspect of it. Specifically, why it's not necessarily a change in pressure of a system that drives that system to a new equilibrium, but a change in pressure due to a change in volume.
From a mathematical standpoint it's extremely straightforward. For the reaction:
N2O4(g) ::equil:: 2NO2(g)
The value of Q is:
[itex]Q = \frac{(P_{\ce{NO2}}/P^{\circ})^2}{P_{\ce{N2O4}}/P^{\circ}}[/itex]
So if I compress the system and increase the partial pressures I'll drive the equilibrium towards more N2O4. BUT if I introduce a bunch of Ne gas to the system and maintain the volume (and increase the total pressure of my container) I will do nothing to the partial pressures and I won't drive the equilibrium in any direction.
What I'm having a hard time understanding is why. I've done some reading here (https://en.wikipedia.org/wiki/Le_Chatelier%27s_principle), here (http://www.chemguide.co.uk/physical/equilibria/lechatelier.html), and here (http://chemwiki.ucdavis.edu/Physical_Chemistry/Equilibria/Le_Chatelier's_Principle), but in each case they explain that this phenomenon occurs but not really why.
In my mental model I imagine the container being compressed and these molecules bumping around each other more and being "unhappy" which causes them to counteract that change by reducing the number of molecules and reducing the pressure. And so something similar should happen when Ne gas is introduced.
Since that doesn't happen, obviously my mental model is wrong/too simplistic. Why am I wrong and what is a better way to think about it?