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?

What sparked this question, is the following statement from a (school-issued) book of mine:

The dissolution of NaOH in water is exothermic. According to Le Chatelier's principle decreasing the temperature of the solution should favor solubility, however, in reality solubility of NaOH increases with temperature. This is therefore an exception.

However, if it weren't for the sharp users of Chemistry.SE I would've accepted my book's claims (since it corroborates seemingly well with common experience). O:)

I did find this article on the ACS page, but it's paywalled >_<

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    $\begingroup$ Most introductory textbooks make overly-broad statements that need to be corrected in later chapters or books - science is an onion where you have to peel back the layers and keep digging to get the the heart of the matter. But, Wikipedia's entry discusses this a bit and at least hints at why there are exceptions. $\endgroup$
    – Jon Custer
    Jan 12 '18 at 11:44
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    $\begingroup$ If Le Chatelier seems violated, then you're regarding an incomplete description of the problem. To actually violate Le Chatelier, you would need some seriously perverted thermodynamics i think. I'd be very happy to see an answer with some reasonable proof for this gut feeling of mine. ;-) $\endgroup$
    – Karl
    Jan 12 '18 at 19:13
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    $\begingroup$ @JonCuster I read that Wikipedia page twice, but couldn't find any reference to any exception of Le Chatelier's principle, not even any hints to it, as you said :/ $\endgroup$ Jan 17 '18 at 2:44
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    $\begingroup$ An exception is known in the ammonia (Haber) synthesis reaction when extra nitrogen gas is added at constant T, P, and moles of hydrogen & ammonia. See Uline & Corti, J. Chem Educ 2006, v83, p138 for a detailed discussion. Depending on nitrogen mole fraction the reaction moves to the left not right. $\endgroup$
    – porphyrin
    Jan 17 '18 at 14:34
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    $\begingroup$ @paracetamol, I agree that adding material to a system already at equilibrium is stretching the point a lot. I prefer a more conservative definition of Le Chatelier that is along the lines of 'If the system is in a stable equilibrium then any spontaneous change in its parameters will be counteracted to restore the equilibrium'. If this were not true then any fluctuation would be amplified and the equilibrium would not be stable. $\endgroup$
    – porphyrin
    Jan 17 '18 at 22:40

There are many exceptions to Le Chatelier's Principle (LCP) so it can certainly be violated. Recently, I have read many articles pertaining to this violation so I believe I should be able to provide an informed answer. Most of the violations of the principle can be seen in gaseous systems. I would not call them "violations", rather, just cases where LCP does not work so well

One of the most classic cases for such an exception, as presented by Cheung (2004), is the addition of nitrogen gas, at constant temperature and pressure, to an equilibrium mixture containing $\ce {N2}$, $\ce{H2}$ and $\ce {NH3}$. The mechanical approach invoking LCP would quickly conclude that there would be an increase in the amount of the gaseous product due to an increase in the amount of one of the reactant gases, favouring the forward reaction. However, this approach fails to consider the instantaneous increase in volume of the system, which would cause the partial pressure of $\ce {H2 (g)}$ to decrease. Now, if we consider the principle again, taking into account both changes, we would realise that the principle would tell us that the equilibrium would shift in both directions. Thus, we see that there is a possibility that LCP cannot give a definite prediction if the number of moles of gaseous products in the balanced equation is not equal to the number of moles of gaseous reactants.

Another example cited by Cheung (2004) is the dimerisation of $\ce {NO2}$. Consider an equilibrium mixture of $\ce {NO2}$ and $\ce {N2O4}$ in a sealed container. When we raise the temperature, we may expect that the position of equilibrium to shift to the left since the backward reaction to form $\ce {NO2}$ is endothermic. However, there is also an increase in the total pressure of the system when the mixture is heated. Again, by LCP, the pressure increase would likely shift the position of equilibrium to the right to the side with less moles of gaseous molecules. This is another case where there are indeterminate predictions given by LCP.


Cheung, D. (2004). The Scientific Inadequacies of Le Chatelier's Principle. Hong Kong Science Teachers' Journal, 22(1), 35-43. Retrieved July 27, 2018, from https://www3.fed.cuhk.edu.hk/chemistry/files/LCP.pdf.

Cheung, D. (2009). The Adverse Effects of Le Chatelier's Principle on Teacher Understanding of Chemical Equilibrium. Journal of Chemical Education, 86, 514-518. doi:10.1021/ed086p514

Quilez-Pardo, J., & Solaz-Portoles, J. J. (1995). Students' and teachers' misapplication of le chatelier's principle: Implications for the teaching of chemical equilibrium. Journal of Research in Science Teaching, 32(9), 939-957. doi:10.1002/tea.3660320906

  • $\begingroup$ The ammonia reaction is exactly what the paper @paracetamol linked talks about. $\endgroup$
    – ringo
    Jul 30 '18 at 5:06
  • $\begingroup$ Your examples prove the opposite of what you say: Once you look at your system properly, you find that LeChatelier is not violated at all! $\endgroup$
    – Karl
    Dec 29 '19 at 18:15
  • $\begingroup$ @Karl Yes, you are correct. I would not consider them "violations". These cases represent instances where the principle does not work so well (i.e. they indicate limitations of the principle). $\endgroup$ Dec 30 '19 at 2:37

I think Le Chatelier's principle is closely related to thermodynamics. Take the Clausius-Clapeyron equation for example. When water freezes into ice the volume increases. Therefore adding pressure helps ice melt into water at lower temperatures, which is how (I was told) skaters can skate on ice with less friction. Le Chatelier's principle is in the same spirit.

When applying Le Chatelier's principle to a chemical equilibrium, one should remember that the "counteract" statement only applies to intensive quantities. For extensive quantities, the shift of equilibrium actually helps with the applied change. For example, let's compare a normal ideal gas with no chemical equilibrium and the $\ce{NO2}$ gas with the equilibrium


If we add pressure to $\ce{NO2}$, the equilibrium moves forward to counteract the increase of pressure. But in doing so, the volume of the system becomes easier to compress (needing less pressure) than a gas that satisfies the equation $pV=nRT$ with no such reaction.

As for the dissolution of $\ce{NaOH}$ in water, I think there's more than just $$\ce{NaOH->Na+ + OH-}$$ going on. For example, the mixing entropy of $\ce{Na+}$ and $\ce{OH-}$ with the water molecules, the shift of the $\ce{H2O<=>H+ +OH-}$ equilibrium, and effects on the hydrogen bonds in water, etc. I think one (or maybe a computer) can still in principle correctly reproduce the trend of how solubility depends on temperature once all players of the game are identified. The water solution is a much more complicated system than a gas reactor. Thinking of heat as a product and raising temperature as providing more "heat" and applying Le Chatelier's principle this way is an oversimplification of what temperature can do.

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    $\begingroup$ Ice skating isn't due to pressure-induced melting. The surface of near-melting ice has a more disordered structure which makes it easy to deform. $\endgroup$
    – gsurfer04
    Jan 17 '18 at 17:54
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    $\begingroup$ @gsurfer04, thanks. I suspected this because skaters move very quickly and I doubt ice would have time to melt. But another experiment that cuts a big piece of ice with an iron wire attached to hanging weights should be genuine. $\endgroup$
    – Zhuoran He
    Jan 17 '18 at 19:03

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