Le Chatelier’s principle is one of the most important concepts with regard to chemical equilibrium. It allows us to predict the behavior of a system in equilibrium under various conditions. In this tutorial, we will cover the definition of Le Chatelier’s principle, as well as how it applies to changing concentrations, temperature, and pressure.
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- Gas pressure
- Dalton’s law of partial pressures
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- Common Ion Effect
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Overview: Dynamic Equilibrium
Dynamic equilibrium, or chemical equilibrium, refers to the state a chemical reaction is in when the forward and reverse reactions are at equal rates. The concentrations of products and reactants both remain constant. This does not mean that the concentrations of reactants and products necessarily must be equal, just that neither concentration is changing, since the rates of formation are equal.
Chemical equilibrium and dynamic equilibrium are the same in most cases, although technically you could have a case of dynamic equilibrium where there are no chemical changes.
What is Le Chatelier’s principle?
Le Chatelier’s principle states: if a system in a state of dynamic equilibrium is disturbed by a change to its conditions, then the position of equilibrium will shift to counteract the change. Let’s explore how this applies, specifically, to different cases.
How does Le Chatelier’s principle apply to specific cases of dynamic equilibrium?
For the next few explanations, let’s use the example of the following chemical equilibrium:
N2 (g) + 3H2 (g) ⇋ 2NH3 (g)
This famous reaction for the synthesis of ammonia is the Haber process.
1. Changes in concentration
If you increase the concentration of one of the reactants, then the position of equilibrium will shift towards the products to counteract the added reactant. For example, if more N2 is added to the system, it will be reacted with H2. This will, in turn, produce more ammonia.
Similarly, if you decrease the concentration of one of the reactants, then the position of equilibrium would shift to the left. In this case, if N2 was removed, more of the NH3 would decompose into the reactants.
This same concept applies to situations where the product is removed from the system—the position of equilibrium shifts to the right to counteract the change, producing more of the product.
2. Changes in pressure
Changes in pressure only apply to reactions involving gases—this includes reactions where not all of the reactants are gases.
If the pressure is increased, then the position of equilibrium will shift to decrease the pressure. It will shift towards the side with less moles of gas. This is because the side with less moles of gas will have less gas molecules to collide with the sides of the container, creating pressure. Essentially, this shift produces fewer molecules of gas in order to reduce the pressure. In this example, then, equilibrium would shift to the right. The left (reactants) side of the equation has 4 moles of gas, and the right (products) side has 2 moles.
Similarly, if the pressure is decreased, then the position of equilibrium will shift to increase the pressure. It will shift towards the side with more moles of gas. Again, this is to produce more molecules of gas, which create a greater pressure when they collide with the walls of the container. Hence in this example, equilibrium would shift to the left.
Note: if there are an equal number of moles of gas in both sides, then increasing/decreasing the pressure will have no effect on the position of equilibrium. Similarly, the addition of an inert gas has no effect on the position of equilibrium. Though it increases the total pressure, it does not react with the other gases, so it does not increase their partial pressures.
Remember that increasing or decreasing the volume of a system’s container impacts the pressure. Transferring the reaction to a smaller volume increases the pressure, and a larger volume decreases the pressure.
3. Changes in temperature
The forward reaction in this case is exothermic, meaning that the reverse reaction is endothermic. This is important to know when considering the changes in temperature.
If the temperature is increased, then the position of equilibrium will shift to reduce the temperature, meaning it will need to absorb the extra heat. Thus, it will shift in the direction of the endothermic reaction (as an endothermic reaction absorbs heat). In this case, as we said, the back reaction is endothermic, thus the position of equilibrium will shift to the left.
If the temperature is decreased, then the position of equilibrium will shift to increase the temperature, meaning it will need to release extra heat. Thus, it will shift in the direction of the exothermic reaction (as an exothermic reaction releases heat). In this case, as we said, the forward reaction is exothermic, thus the position of equilibrium will shift to the right.
Note about catalysts & Le Chatelier’s principle
The addition of a catalyst, actually, has no impact on the position of equilibrium. This is because catalysts speed up the rates of both the forward and reverse reactions. This means that the two rates remain equal, as they both change by the same amount.
Example of Le Chatelier’s principle: Cobalt Chloride complexes
Cobalt (II) chloride is often used to demonstrate Le Chatelier’s principle. A solution of cobalt chloride is pink, but if enough concentrated hydrochloric acid is added, a blue complex is formed.
[Co(H2O)6]2+(aq) (pink) + 4Cl–(aq) → [CoCl4]2–(aq) (blue) + 6H2O(l) + heat
If you add more reactants, the solution turns more blue. If you add more products, the chemical equilibrium shifts to the left, and the solution turns pink. It is easy to get an intermediate stage that is violet. So adding more HCl, adding table salt, or cooling the solution, will make it more blue. Adding water, or heating the solution, will turn it more pink.