Talk:Le Chatelier's principle

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WHY???

WHY is Le Chatelier's principle true? What are the physical reasons for it? Is it a 2nd law of thermodynamics thing?

—The preceding unsigned comment was added by 67.99.238.198 (talk) 01:54, 12 January 2007 (UTC).

It's related to kinetics, if my understanding is correct. (Perhaps a professional who has more experience could check this, to make sure that it's accurate?) There is actually no "drive" or "force" that pushes an equilibrium in a certain direction, it's just that when one reactant has an increase in concentration, it collides with its fellow reactants more often, producing a higher rate of reaction in the forward direction than the product side is reacting in the reverse direction. Eventually, the increase of products means that their collisions increase too, causing the reverse-reaction rate to increase as well. Once the new forward and reverse rates have become equal, i.e., Q = K, the equilibrium has been reached again. I hope this was helpful. 70.149.178.146 03:37, 9 March 2007 (UTC)

My knowledge of this is not as good as it should be, but it is in my opinion that when you look at it in terms delta G, you realize that since delta G is neither positive nor negative, the reaction will no longer become spontaneous in either direction. If the reaction is not spontaneous in either direction, the reaction will stop moving in either direction. So maybe that will help.Q E11even 04:19, 14 March 2007 (UTC)

Is that spelling correct?

I thought it was spelt "Le Chatelier", without the circumflex. I have a copy of the original paper by Henri Le Chatelier (Comptes Rendus de l'Académie Française 99, 786-9 (1884)) and there's no accent there. Neither is there one in an article I read on him in Chemistry in Britain, nor in "Modern Physical Chemistry" by Liptrot et al, nor in the classic "Physical Chemistry" texts of Moore or Atkins, nor in Encyclopædia Britannica. In fact I don't think I've seen the accent anywhere other than on Wikipedia.--CSM 18:30, 17 April 2006 (UTC)

I agree. Even Wikipedia's own Henri Louis Le Chatelier Article does not use the circumflex. I suggest a title change. --anomymous user 2:01, 24 April 2006 (EST)

I just want to add something... im a 15 year old studying this for my GCSE's but it still amazes me how some people can be so pedantic! The person who added the origional article (whoever he or she may be) deserves respect for thier work... not people critising about the spelling of a forign word.

Hello... its me again. Just wanted to prove a point.. ORIGIONAL, CRITISISING and FOREIGN.

:I think you just proved his point.134.88.164.209 01:55, 6 March 2007 (UTC)

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Chemistry (8e, Raymond Chang, McGraw-Hill) spells it Le Châtelier.

Dmbrown00 04:50, 27 January 2007 (UTC)

Chemistry the Central Science (9e, Theodore Brown, H. LeMay Jr., Bruce Bursten, and Julia Burdge, Pearson Education Inc.) also spells it Le Châtelier. Fluffybun 17:11, 4 March 2007 (UTC)

Hey is that the standard text book or the AP one? I've the AP :) Arc88

Effect of Adding an Inert Gas

Could someone please verify whether the edit on 3 January 2008 concerning the effect of adding an inert gas was accurate?

Of the first 20 results obtained when searching for '"le Chatelier's principle" "inert gas"' on Google, 12 state that adding an inert gas has no effect on equilibrium when volume is held constant and only 5 state that equilibrium is affected, of which 4 quote Wikipedia and the fifth is Yahoo Answers.

I thought that equilibrium was only affected by the partial pressure of each gaseous substance, which is unaffected by adding an inert gas, rather than the total pressure of the system? Vqors (talk) 09:20, 14 February 2008 (UTC)

You are correct. I tracked down a legit source and made the changes. I also changed the "changes in pressure" section, since that, too, implied a change in total pressure would shift the equilibrium. This is one of those misconceptions that I can't stand. Brittlandk (talk) 08:23, 23 February 2008 (UTC)

Let us come to a consensus on this issue before I roll back the change. I am in agreement with the above that when an inert gas is added at fixed volume, the equilibrium is unaffected (as stated in my original post) because the concentrations of the other gases are still the same. However, adding an inert gas at constant pressure and allowing the volume to increase (ex. inside a balloon), would in fact result in a decrease in the concentrations of all other gases involved, thus resulting in a shift.

For example, let's consider CO + 2 H2 <--> CH3OH in a balloon. If we add enough helium to double the volume of our balloon (still at same pressure), the concentrations of CO, H2 and methane would all be halved.

Yes, this agrees with my original source (and others).

Since the decreases in concentrations would affect the left side more than on the right, a shift would occur to the left to "fill the void". Please verify your sources and advise. Brent Woods 22:44, 1 March 2008 (UTC)

In summary, I agree that adding an inert gas at constant volume (thus increasing total pressure) will not result in a shift because the partial pressures are unchanged. However, if the addition of the inert gas causes a volume increase, the partial pressures of all gases involved are decreased, thus causing a shift towards the side with more moles. Brent Woods 15:56, 2 March 2008 (UTC)

Yes, I see what you are saying. It is important to make the distinction whether the change in total pressure is allowed to change the volume of the system. Thank you for clarifying this.Brittlandk (talk) 04:38, 12 March 2008 (UTC)

...so effectively it depends whether it is an open or closed system. Open system, outside a balloon, would not affect the pp of each reactant/product, so make no difference, but a closed system would affect pp of a reactant/product and so would change equilibria. yeah??...--82.9.21.247 (talk) 19:22, 15 April 2008 (UTC)