Le Chatelier's principle: Are there any exceptions?

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ChloePCooper
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The way Le Chatelier's principle is presented in most introductory chemistry books (high-school) is as though it's an indisputable law of the physical world (in the sense that we're never shown an exception, not that its universality is explicitly stated).

This is supposed to differ from the more frequently encountered half-assed "laws" like Ohm's Law or Hooke's Law (which are really just generalizations drawn from a few, everyday cases... not real "laws" as such).

Now I'm curious. Is there really no exception to Le Chatelier's principle (for a closed, isolated system)?

Wording this differently: Is there any situation (closed, isolated system) where conditions conducive to Le Chatelier's principle are in place, yet the expected Le Chatelier's "response" (the tendency to oppose change) is not observed upon effecting the said change (in parameters such as concentration, temperature, volume, etc.)?

If there are any exceptions that one comes across in everyday life, I'd prefer to hear those (but since that's extremely unlikely, I suppose just about any exception will do).

Are there any theoretical grounds for exceptions to Le Chatelier's principle?

If Le Chatelier's principle cannot be violated, then why is it so?
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Kian Stevens
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This is a very good question!
As far as I'm aware, there are no exceptions to it, as Le Chatelier's principle is just an observation based on thermodynamics and kinetics
Obviously, thermodynamics is never 'wrong' just like other explanations for the universe aren't 'wrong', it isn't just a generalisation based on what is seen... This is why the principle is, as far as I'm aware, never violated, as it's not theoretical and is instead essentially thermodynamics etc. in words

The most primitive way to think about this, which crops up every second of your day, is via Newton's third law... As we know, all actions have an equal and opposite reaction, i.e. whatever you push will push back on you to achieve an equilibrium state... You're pushing down on the floor, and the floor will push up against you, that sort of thing
This is the same thing here, as doing one thing to a system in dynamic equilibrium will cause the system to do the opposite thing, again to achieve an equilibrium state

Let's generalise a system and some situations which may arise, as to why this is the case:
  • Changing the concentration of reactants: increasing the concentration of one side will shift the equilibrium so that the concentration decreases again, meaning equilibrium is restored. For example, increasing the concentration of the reactants will shift the equilibrium to the right, so that the concentration of products increases (and simultaneously the concentration of reactants decreases) so that equilibrium is restored
  • Changing the temperature of the system: increasing the temperature of the system will shift the equilibrium so that the temperature will decrease again (and vice versa), meaning equilibrium is restored. For example, if the forward reaction is endothermic and you increase the temperature, the equilibrium will shift to favour the products as the endothermic reaction will absorb any added energy and the temperature will decrease again... Or, if the forward reaction is exothermic and you decrease the temperature, the equilibrium will shift to favour the products as the exothermic reaction will release more energy and the temperature will increase again... This can be said about all possible situations
  • Changing the pressure of the system: this is essentially the same as changing the volume, as P \propto \frac{1}{V}. Just like all the times before, increasing the pressure of the system will cause the equilibrium to shift to favour the side which causes the pressure to decrease again (and vice versa), meaning equilibrium is restored

So as you can see for a generalised system, these processes will occur simply to maintain an equilibrium, so will happen regardless of the system
As mentioned, it will happen all due to thermodynamics and kinetics
Last edited by Kian Stevens; 1 year ago
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