Applications of Le Chatelier’s Principle in Chemical Equilibrium

Le Chatelier’s Principles

Le Chatelier’s principle states that when dynamic equilibrium is disturbed by a change in conditions, the position of the equilibrium shifts to cancel the change and restore the equilibrium. When a chemical reaction is in equilibrium and the pressure, temperature, or concentration of the product or reactant changes, the equilibrium shifts in the opposite direction to accommodate the change. This page describes the changes in the equilibrium position due to these changes and briefly explains why the catalyst does not affect the equilibrium position. Le Chatelier’s principle is an observation of the chemical equilibrium of the reaction. This indicates that changes in system temperature, pressure, volume, or concentration will bring predictable and opposite changes to the system, reaching a new equilibrium state. You can actually use Le Chatelier’s principle to understand the reaction conditions that promote product formation. This idea was originally discovered and developed by Henry Louis Le Chatelier and Karl Ferdinand Brown.

Changes in concentration: Increasing the concentration of one of the reactants shifts the equilibrium position towards the product, cancelling out the added reactants. For example, when N2 is added to the system, it will be converted to H2. This will produce more ammonia. Similarly, lowering the concentration of one of the reactants shifts the equilibrium position to the left. In this case, when N2 is removed, more NH3 is decomposed into a reactant. 

The same concept applies to situations where a product is removed from the system. The equilibrium position shifts to the right to counteract changes and produce more products.

Changes in pressure: Pressure changes apply only to reactions involving gas. This includes reactions where not all reactants are gas. Increasing pressure shifts the equilibrium position and reduces pressure. Shift to the side with less gas mol. This is because on the side with a small number of moles of gas, there are few gas molecules that collide with the side surface of the container and generate pressure. Basically, this shift produces less gas molecules to reduce the pressure. Therefore, in this example, the equilibrium shifts to the right. The left side of the equation (reactant) has 4 moles of gas and the right side (product) has 2 moles. As with the, lowering the pressure shifts the equilibrium position and raises the pressure. Switch to the side with the most gas moles. Again, this helps create more gas molecules and creates more pressure when they collide with the walls of the container. Therefore, in this example, the equilibrium shifts to the left. 

Addition of an Inert Gas: What is the equilibrium position of the reaction when an inert gas such as krypton or argon is added to the reaction vessel? Answer: There is nothing. Remember to always shift the system so that the product-to-reactant ratio remains the same Kp or Kc. Since the inert gas does not react with either the reactant or the product, it does not affect the product/reactant ratio and thus the equilibrium.

Changes in temperature: In this case, the forward reaction is exothermic, which means that the reverse reaction is endothermic. This is important to know when looking at temperature changes. As the temperature rises, the equilibrium position shifts and the temperature drops. That is, it needs to absorb excess heat. Therefore, it shifts towards an endothermic reaction (because the endothermic reaction absorbs heat). In this case, the equilibrium position shifts to the left because the reverse reaction is an endothermic reaction as described above. When the temperature drops, the equilibrium position shifts and the temperature rises. This means that additional heat needs to be released. Therefore, it shifts towards an exothermic reaction (because the exothermic reaction releases heat). In this case, as mentioned above, the equilibrium position shifts to the right because the forward reaction is exothermic.

Effect of a catalyst: The catalyst speeds up the reaction without being consumed in the reaction. The use of a catalyst does not affect the position and composition of the equilibrium of the reaction, as both the forward and reverse reactions are accelerated by the same factor. 

For example, consider the harbour process for synthesizing ammonia (NH3). 

N2 + 3H2 ⇌2NH3 

In the above reaction, in the presence of iron (Fe) and molybdenum (Mo), it acts as a catalyst. They accelerate all reactions, but do not affect equilibrium.

Conclusion

When external stress is applied to a system in dynamic equilibrium, the system shifts the equilibrium position to nullify the effects of stress. 

You can stress the chemical system by changing the concentration, pressure, or temperature. Therefore, it can be expressed as: 

When the chemical system is disturbed in dynamic equilibrium by changing the concentration of the reactant or product. Or by changing the partial pressure of the gaseous reactant or gaseous product. Or in the case of temperature, the equilibrium position is changed in that direction and a new equilibrium state is established. H. Either a forward reaction or a reverse reaction is preferred.