Moving Beyond Le Châtelier

test tubes used in equilibrium activity

It is time for us as chemistry teachers to move beyond the Le Châtelier Principle as justification for why disturbances to equilibrium systems cause particular “shifts”. I came across the idea of ditching Le Châtelier in an article by Eric Scerri for the Royal Society of Chemistry in which he summarizes the shortcomings, and outright failures, of the principle. He expands on this idea in a post on his blog. I suggest you read both of these articles in full, as I will only summarize Scerri’s key points. I will then focus on how I approached equilibrium in my class this year to give my students a more rigorous understanding of the concept that can be more broadly applied.

The Problem with Le Châtelier

The Le Châtelier Principle is generally stated along the lines of: “If a system in a state of equilibrium is disturbed, the position of equilibrium will shift in order to counteract the change”. Other terms used in lieu of “counteract” are “oppose”, “relieve”, or “reduce”.

My immediate problem with this statement has always been that it is teleological. That is, the explanation for why the phenomenon occurs is the existence of an intrinsic nature for it to occur: the shift in equilibrium happens in order to restore equilibrium. This often lends itself to personification, “the system wants to restore equilibrium” or “needs to restore equilibrium” and other such appealing, yet fallacious statements. A priori arguments justifying the change based on the result should be replaced with a more rigorous a posteriori logic based on the value of the reaction quotient, Q.

Scerri points out that the Le Châtelier Principle accurately predicts shifts in only one case: changes of concentration. In the cases of pressure and temperature changes students can use sound judgement and come to incorrect conclusions. Consider a system at equilibrium is contained in a closed vessel. The volume of the vessel is reduced. A student could soundly reason that the reaction would counteract the change in volume by shifting to produce more gas so that the volume increases. This counteracts the change, but of course is incorrect. The system actually shifts to produce fewer gas particles.

I agree with Scerri that a principle that fails two out of three times should be abandoned. What follows is how I went about this in my AP Chemistry class this year. In this post, I will discuss concentration and pressure disturbances to equilibrium. I will address temperature disturbances in an upcoming post.

My In-Class Sequence

I do not cover equilibrium in my first year chemistry course, so this is my AP students’ first exposure to the topic. This comes immediately after we have studied chemical kinetics. The first activity I do is a variation on a common activity with bingo chips in which students simulate different equilibrium systems.1 From this activity, students uncover the meanings of dynamic equilibrium, the reaction quotient, the equilibrium constant, and the effects of adding or removing reacting species from an equilibrium mixture.

Following this activity, we summarize our findings in a class discussion and briefly review notes on equilibrium and the relationship between Q and K. Then, students complete the new AACT Simulation as a homework assignment to build their fluency with Q and K.

The next class I present the students with two equilibrium systems with discussion questions. 

     1. If additional Cl- ions were added to the system, would the rate of the forward or reverse reaction increase?

     2. If CuCl42- is added to the system, would the rate of the forward or reverse reaction increase?

To both of these questions, I am looking for a kinetic argument along the lines of, “If additional Cl- ions are added, there are more likely to be successful collisions between reactant molecules that increase the rate of the forward reaction to produce more products.”

Following the discussion, I demonstrate this for my students by adding 12 M HCl dropwise to a test tube containing 0.5 M CuCl2(aq). The solution turns a bright yellow/green color, confirmation of their prediction.

Then I ask them the “challenge question”: if the system is diluted by the addition of water, will this affect the equilibrium position? If so, will the rate of the forward reaction or the reverse reaction increase?

My class is small, only six students, so they discussed the question together. With larger classes, I would break them in to groups and let them use whiteboards. The conversation that followed was a rich discussion. Some students initially thought no change would occur because water does not appear in the equilibrium expression, but eventually the conversation turned to Q and K. In doing their analysis, they independently assumed 1 M concentrations at equilibrium and did the following analysis:

They correctly concluded that the value of Q would increase regardless of the actual value of K, so the reverse reaction rate would increase and more reactants would be produced.

To see if they are correct, we use a wash bottle to add water to the green/yellow solution from the first demonstration and they observe that the solution turns back to blue/green, evidence that the reverse reaction has increased in rate to produce more Cu2+(aq) ions.

Figure 1 - 0.5 M CuCl2 (left); After adding 12M HCl (center); Center tube diluted with DI water (Right)

Then we turned our attention to System 2.

I ask them what would happen if various species were added or removed from the equilibrium system, and we discussed the effect on Q. Then I asked them the challenge question for this example, “If the volume of the container were halved, would this affect the position of equilibrium. If so, would the rate of the forward or reverse reaction increase?” Having the experience with the first system, they decided to assume 1 atm equilibrium partial pressures and realized that halving the volume would double each of the individual partial pressures of each gas:

Q is now greater than K. Therefore, the reaction will proceed to produce more SO2Cl2(g).

The students tried to solve a problem based on knowledge they already had about the reaction quotient, and came to the correct conclusion. Using the Le Châtelier Principle, using their previous knowledge could lead them to an incorrect conclusion, but using the reaction quotient, they will be correct.

Conclusion

Treating disturbances in equilibrium in this manner allows students to logically reason through any situation and even address difficult situations, like dilution, with ease and accuracy. Any existing chemistry curricular materials can be easily modified to remove the teleological Le Châtelier Principle and replace it with more rigorous Q vs. K arguments that connect everything back to the simple idea of the ratio of reactants and products in a system. Students will now be able to produce logical justifications instead of relying on the crutch response of “because Le Châtelier!”

For AP teachers, this approach would have made approaching the notorious 2016 Free Response Question #6 much easier. Part (b) of this question asks about the effect of dilution on an aqueous system, the mean score was 0.45/4, and the modal score was zero. I hear many AP teachers say that Q vs. K arguments are now “in vogue”, as if with resentment, but these arguments represent more logically sound chemistry. Isn’t that what we all aspire to teach?


1 My activity is based on “Penny-Ante Equilibrium” from Flinn ChemTopics Volume 15: Equilibrium. You can find it below under Supporting Information.

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Comments 1

Tom Kuntzleman's picture
Tom Kuntzleman | Wed, 12/06/2017 - 10:53

Hi Kaleb. I enjoyed this article. It gave me a lot of new things to think about regarding the principle of Le Châtelier (POL), especially your point that “My immediate problem with [POL] has always been that it is teleological.” I’ve not heard that argument before. I have been familiar with Eric Scerri’s work on the POL for quite some time, mostly because I happen to be of opposite opinion: I like the POL. (In fact, I’m not convinced that the POL leads to erroneous results when it is correctly applied and pertinent assumptions are made clear). We can talk about the validity of the POL if you’d like, but I’d actually prefer to move away from a discussion of whether POL is correct or not to something a bit deeper (for now I will begrudgingly grant for sake of argument that POL can lead to erroneous predictions).

 

The point I’d like to claim is that it is sometimes okay to sacrifice rigor – and even truth – to enhance student understanding. In my opinion, if a principle, model, or rule is correct more than 90% of the time (or under certain conditions), then I’d say it is okay to use it with students. I would also say that such “partially correct” principles, models, and rules often create a scaffold onto which further student learning can be built. A few examples of these “partially correct” concepts include the ideal gas laws, the Bohr Model (this model is wrong 118 out of 118 times), and Newtonian physics. I happen to think that the POL –whether it is 100% correct or not – when appropriately presented is a good qualitative principle that helps students predict outcomes when a system has been bumped out of an equilibrium state.