[jsrozner]

So I've been puzzling over this for quite some time and have been unable to resolve it...


Assume that we have some reaction A + B -> AB and it's endothermic, so
heat + A + B -> AB.


By L'Chatelier's Principle, adding heat (increasing the temperature) should shift the equilibrium to the right.
But, if we look at the problem from a thermodynamic standpoint, I seem to get a different answer.
dG = dH-TdS
For the above equation, delta H is positive (endothermic) and delta S is negative (entropy decreases). If we increase the temperature in the equation G = H - TS, delta G should become more positive because S is negative. This would predict a shift to the left, rather than a shift to the right, contradicting the prediction based on L'Chatelier's Principle.


I get the same sort of contradiction when I consider        AB -> A + B + heat.


The other two possibilities A + B -> AB + heat and
heat + AB -> A + B don't result in the same problem.


There's probably some sort of convention or concept I'm ignoring or don't understand...or these two rules apply to different things (which I'm failing to distinguish between). Either way, hopefully someone can point out the flaw(s) in my logic.


Thanks

[Jorriss]

This is a good question, and I don't know much thermodynamics but here's my thinking.

Gibbs Free Energy is for when a reaction will occur spontaneously - without continually adding heat, so does the usual line of thinking apply when you aren't looking at systems in thermal equilibrium?