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Chemistry Forums for Students => Physical Chemistry Forum => Topic started by: yxy112358 on March 03, 2014, 04:38:49 AM

Title: Equilibrium shift by adding inert gas - constant pressure / constant volume
Post by: yxy112358 on March 03, 2014, 04:38:49 AM

When I learn the chemical equilibrium, I have a question that has confused me for long time. When I add inert gas into an equilibrium system,

(1) if it is under constant temperature and constant volume, I know that the equilibrium does not shift, since the molarity of each of the original part does not change.

(2) if it is under constant temperature and constant pressure, I am told that the equilibrium will shift towards the direction with more molecules. But why is that? Could anyone give me an explanation with deduction from equilibrium constant or Dalton's law of partial pressures or any other related concepts?

Thank you all.  :)

Title: Re: Equilibrium shift by adding inert gas - constant pressure / constant volume
Post by: Borek on March 03, 2014, 04:51:22 AM

Why don't you try to derive it by yourself? Assume you know Kp or Kc, assume some initial V and P, see what happens when you add inert gas - what happens to partial pressures (or concentrations - depending on whether you start with the Kp or Kc). System has to react to reach equilibrium - you should be able to predict in which direction the equilibrium will move.