[Burner]

In our school, we learnt to use Le Chatelier's Principle(I abbreviate it as LCP in the following) to predict the change in equilibrium positions of chemical systems AND its limitations. It states that if a equilibrium is subjected to a change, the equilibrium position will shift in a direction to minimise/oppose the change. We also learnt to use the reaction quotient(Q_c) to predict it.

However, my chemistry tutor told me to avoid using LCP as much as possible. He said that LCP has a lot of limitations such as oversimplifying the actual situation.

For example,

N₂)O₄(g) 2NO₂(g)  ΔH=+58 kJ mol^-1

The above reaction is carried out in a container with fixed volume.

If we heat up this equilibrium system, according to LCP, the equilibrium position will shift to the right and lowers the temperature by absorbing heat. However, the total pressure of the system increases also because of the increase in temperature. According to LCP, increase in temperature with shift the equilibrium position to the left to decrease the number of moles of gases, creating a self-contradictory prediction.

So, I would like to know:
- Do we still use Le Chatelier's Principle to predict changes in equilibrium positions? If yes, in what case do we use it or avoid it?
- Is Le Chatelier's Principle frequently challenged by chemists or be formally disproved or replaced?

P.S. I didn't learn equilibrium systems involving K_a, K_sp and etc.. Please use more general equilibrium systems examples if you would like to explain. Thanks.

[thetada]

Educators have raised issues over Le Chatelier's principle for decades. There's a thought provoking article linked below, the comments section of which is also well worth a read.

http://www.rsc.org/blogs/eic/2015/07/le-ch%C3%A2telier-principle-equilibrium

I teach A-levels in the UK, which I mention because our students are often asked to state the definition of the principle on the exam, so it's worth thinking about the grasp of the principle your own exam board requires.

[mjc123]

LCP is useful if used with care; it's important to understand the physical situation and what the change actually does. In particular, changes of pressure can give rise to confusion, for example in the following thread: http://www.chemicalforums.com/index.php?topic=83255.msg302201#msg302201. There, the total pressure was increased (at constant T and V) by adding an inert gas, with no effect on the position of equilibrium because the concentrations of reacting gases did not change. Likewise, when you heat up the system at constant volume, the concentrations don't change, so there is no driving force from that to shift the equilibrium. The equilibrium will shift because of the temperature rise in the endothermic direction.
Effectively, pressure in equilibria is a surrogate for concentration; P = (n/V)*RT. If we compare the equilibrium constants written in terms of concentration, K_c, and pressure, K_p;
K_c = [NO₂]²/[N₂O₄];        K_p = P_NO2²/P_N2O4
K_p = K_c*RT
Let us suppose, for simplicity, that ΔH° = 0 (referred to a fixed concentration standard state), so that K_c does not change with temperature. It will be seen that K_p increases with temperature, and this increase matches the way the pressures change with temperature at constant volume, so the position of equilibrium does not change.

[Burner]

If we compare the equilibrium constants written in terms of concentration, K_c, and pressure, K_p;
K_c = [NO₂]²/[N₂O₄];        K_p = P_NO2²/P_N2O4
K_p = K_c*RT

In fact I didn't learn about K_p either. Do you mean that we can express the 'equilibrium position' in different ways, such as by temperature and also pressure(edit)?

the total pressure was increased (at constant T and V) by adding an inert gas, with no effect on the position of equilibrium because the concentrations of reacting gases did not change. Likewise, when you heat up the system at constant volume, the concentrations don't change, so there is no driving force from that to shift the equilibrium. The equilibrium will shift because of the temperature rise in the endothermic direction.

I understand and agree with these deductions. However, it seems to go far beyond Le Chatelier's Principle: We can't simply say that the equilibrium position shifts to the left to decrease the number of moles of gases to 'oppose the change', but we need to consider what factors are being actually changed. So, am I right with that LCP oversimplifies the actual situation? Or there are some more 'ways' to use Le Chatelier's Principle?

Maybe LCP is suitable for qualitatively determine the direction of shift of equilibrium positions involving concentrations only but not with other physical factors such as pressure? By the way, the examples that my tutor used to oppose LCP all involves gases equilibrium systems.

[thetada]

Maybe LCP is suitable for qualitatively determine the direction of shift of equilibrium positions involving concentrations only but not with other physical factors such as pressure? By the way, the examples that my tutor used to oppose LCP all involves gases equilibrium systems.

LCP is a model with limitations like any other. It's a useful springboard from which to explore the intricacies of equilibria, which can, with caution as mjc123 points out, be applied to changes other than concentration. The difficulty with thermodynamics is that you get to a point where only maths will do and as a friend said to me recently, even then the explanations can feel unsatisfactory. Once you know all the relevant equations, you can derive one containing all the desired variables and consider changes within that framework. I suspect that attention paid to LCP is inversely proportionate to one's knowledge of those equations.