Contradiction between Henry's Law and Le Chatlier's Principle

johnny_b_good

Le Chatlier's Principle is used to determine the direction of a reaciton based upon a stress put on the system. In addition, Henry's law states that the solubility of a gas is related to the partial pressure of that gas. Therefore I present a seemingly contradictory setting:

Example for Carbonation of Water:

CO2 (g) ⇔ CO2 (aq) ---- (1)
H2O (l) + CO2 (aq) ⇔ H2CO3 (aq) ------ (2)

If I apply an external pressure on the system with an oxygen take (pure O2 (g) -- I know, very dangerous, but just a theoretical situation), then Le Chatlier's principle would argue that reaction number (1) would shift towards the right to attain the lowest possiblity energy setting. However, if we use Henry's law, then we would say that the partial pressure of CO2 (g) remains constant ... since we are adding pure O2 (g) ... and therefore, the solubility of the gas would remain constant. That is Le Chat's says concentration of CO2 (aq) increases, but Henry's says that CO2 (aq) remains constant.

Am I missing something relatively large here? thanks for the help

Nessdude14

Henry's law requires a constant temperature, and by increasing the pressure to the system, you'll cause a change in temperature, so Henry's law is invalid in this case. The reaction will behave just as Le Chatelier's principle suggests.

A gas law that would be applicable to this scenario is the ideal gas law, PV=nRT. You could determine the change of gas molecules expected from a pressure change if you also know the volume and temperature change.

johnny_b_good

Ok thanks for the help. I forgot that it dealt with constant temperature.

johnny_b_good

Well, on that note. Wouldn't increasing the partial pressure always generate a net temperature increase (no matter how small)? And thus invalidate Henry's Law?