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Le Chatelier’s principle is an explanation of how systems in dynamic equilibrium respond to changing conditions. It states that if the conditions in a closed system change, the position of the equilibrium will shift to counteract the change.
- This article is about Le Chatelier's principle in chemistry.
- We will begin by looking at reversible reactions and dynamic equilibrium, before exploring what Le Chatelier’s principle is and how it can be used.
- We will then look at how external factors affect equilibria in accordance with Le Chatelier’s principle.
- Finally, we will investigate common reversible reactions used in industry.
Before we look at Le Chatelier’s principle in more depth, let’s quickly go over reversible reactions and equilibrium.
Equilibrium Le Chatelier's Principle
As you will have gathered, Le Chatelier’s principle depends on reversible reactions and equilibrium. We’ll start by looking at reversible reactions.
Reversible reactions
Reversible reactions are reactions that form products, which under different conditions can react together to form the original reactants again. (Check out Chemical Equilibrium for more information.)
There are a few terms concerning reversible reactions you should know about.
The reaction that goes from left to right, or from reactants to products, is called the forward reaction.
The reaction that goes from right to left, or from products to reactants, is called the backward reaction.
If there is more of the forward reaction than the backward reaction, we say that the forward reaction is favoured and that the equilibrium has shifted to the right.
If there is more of the backward reaction than the forward reaction, we say that the backward reaction is favoured and that the equilibrium has shifted to the left.
A reversible reaction. Anna Brewer, StudySmarter Original
Dynamic equilibrium
When you start a reversible reaction, the concentrations of reactants and products constantly change. But if you leave the reaction in a closed system, eventually the concentrations level off. When this happens, we say that the reaction has reached dynamic equilibrium.
Dynamic equilibrium is a state of a reversible reaction, in which the concentrations of products and reactants remain constant, and the rates of the forward and backward reactions are the same.
Chemical equilibria are examples of dynamic equilibria. This means that both the forward and backward reactions are constantly ongoing. But because they happen at the same rate, their effects cancel each other out; it looks like there is no overall reaction. In a dynamic equilibrium, although products and reactants are constantly being broken down and reformed again, the overall levels of each species don’t change.
Dynamic equilibrium graph. Anna Brewer, StudySmarter Original
What if we want more of the products than what the system is currently giving us? We can manipulate the equilibrium using Le Chatelier’s principle. We'll look at this next.
Le Chatelier’s principle definition
Let’s now move on to the main focus of this article: Le Chatelier’s principle. As we explored right at the start, Le Chatelier’s principle explains how equilibria respond to changing conditions. It states that disturbing an equilibrium, such as by changing its conditions or environment, shifts the position of the equilibrium to favour the side that reduces the disturbance and opposes the change.
Imagine you are in charge of letting people into a restaurant. The restaurant can only have a certain number of diners in it at any one time - it is limited by its number of seats. At busy times, this means that some people have to queue outside and wait to get in. However, to maximise profits, you ideally want as many people to be eating inside the restaurant as possible. At 2 o’clock, a big party ends, and lots of diners leave at once. The conditions have changed - suddenly there are far fewer people inside. In order to fill up the empty seats, you open the doors and let a lot of people inside at one go. You've reacted to the ‘disturbance’ caused by many people leaving at once by letting more people in, in order to counteract the change the disturbance caused.
Disturbing the conditions of an equilibrium causes the equilibrium to react in response. There are a few different ways in which we can disturb an equilibrium. These include:
- Changing the temperature.
- Changing the concentration of products or reactants.
- Changing the pressure.
Le Chatelier’s principle and temperature
We'll begin by looking at how temperature affects an equilibrium. Remember that Le Chatelier’s principle states that changing the conditions of an equilibrium causes the reaction to shift to oppose the change. Can you predict what will happen if you increase the temperature of a dynamic equilibrium?
The system will try to decrease the temperature to oppose the change. The only way it can do this is by favouring the endothermic reaction - the one that takes in heat as energy.
Here's an example. Look at this equation for the equilibrium involving nitrogen, hydrogen, and ammonia.
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ mol-1
The forward reaction is exothermic, while the backward reaction is endothermic. We can say the following:
- Increasing the temperature favours the backward endothermic reaction and causes the equilibrium to shift to the left. This will oppose the change by decreasing the temperature.
- Decreasing the temperature favours the forward exothermic reaction and causes the equilibrium to shift to the right. This will oppose the change by increasing the temperature.
Le Chatelier’s principle and concentration
What happens when you decrease the concentration of a substance? You essentially just have fewer molecules in the same volume. This is what happens when lots of diners leave the restaurant. In order to oppose the disturbance caused by the changing conditions, we need to let more people in, i.e., we need to increase the concentration of that substance. For example, if we decrease the concentration of the products, the equilibrium will shift to favour the forward reaction to bring the concentration of products back up again. If we decrease the concentration of the reactants, the equilibrium will shift to favour the backward reaction, in order to bring the concentration of the reactants back up again.
You can also think about what would happen if you increased one of the concentrations - say, that of the reactants. The forward reaction would increase to try and ‘use up’ some of the extra reactant molecules. You can see this in the diagram below, where the arrow in bold shows the favoured reaction.
The effect of changing concentration on an equilibrium reaction. Anna Brewer, StudySmarter Original
Here’s the equation for the equilibrium involving ammonia again:
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ mol-1
We can say the following:
- Increasing the concentration of nitrogen and/or hydrogen will favour the forward reaction and shift the equilibrium to the right. This will oppose the change by using up some of the nitrogen and hydrogen, reducing their concentration.
- Decreasing the concentration of nitrogen and/or hydrogen will favour the backward reaction and shift the equilibrium to the left. This will oppose the change by forming more nitrogen and hydrogen, increasing their concentration.
Le Chatelier’s principle and pressure
Provided they are at the same temperature and in the same volume container, all gases have the same pressure per mole. Pressure is caused by the gas molecules randomly colliding with the sides of the container. Now imagine that we increase the pressure of the system. To oppose this change, the system will try to reduce the pressure by reducing the number of collisions that occur. The system can’t change the speed of the particles or the frequency of their collisions, but it can reduce the pressure by decreasing the number of gas molecules in the system.
Fewer molecules, fewer collisions - simple, right? Therefore, increasing the pressure favours the reaction that produces fewer moles of gas. On the other hand, decreasing the pressure favours the reaction that produces a greater number of moles of gas.
Take a look at this equation again:
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ mol-1
The forward reaction produces two moles of gas. The backward reaction produces four moles of gas. We can say the following:
- Increasing the pressure will favour the forward reaction and shift the equilibrium to the right. This is because the forward reaction produces fewer moles of gas than the backward reaction. This will oppose the change by reducing the pressure.
- Decreasing the pressure will favour the backward reaction and shift the equilibrium to the left. This is because the backward reaction produces a greater number of moles of gas than the forward reaction. This will oppose the change by increasing the pressure.
There are two things to note here. Firstly, changing the pressure only affects the equilibrium of gaseous species. You can ignore any moles of solids or liquids in the equation. Secondly, changing the pressure will have no effect on a gaseous equilibrium if both of the reactions produce the same number of moles of gas. There is no way of decreasing the number of moles of gas in the system - no matter which reaction is favoured, the equilibrium won’t change.
Le Chatelier’s principle and catalysts
Catalysts don’t affect the position of equilibrium. This is because they speed up both the forward reaction and backward reaction at the same rate. However, adding a catalyst is useful because it speeds up the time it takes for a system to reach dynamic equilibrium.
Catalysts are substances that increase the rate of reaction without being used up or changed in the process.
Le Chatelier's Principle - Key takeaways
- Le Chatelier's principle states that changing the conditions of an equilibrium reaction causes the equilibrium to shift to oppose the change.
- Increasing the temperature of an equilibrium reaction favours the endothermic reaction, whilst decreasing the temperature favours the exothermic reaction.
- Increasing the pressure of an equilibrium favours the reaction that produces the fewest moles of gas, whilst decreasing the pressure favours the reaction that produces the greater number of moles of gas.
- Increasing the concentration of the reactants favours the forward reaction, whilst increasing the concentration of the products favours the backwards reaction.
- Catalysts do not affect the position of equilibrium, as they speed up both the forward and backward reactions at the same rate. However, adding a catalyst is useful because it speeds up the time it takes for a system to reach dynamic equilibrium.
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Frequently Asked Questions about Le Chatelier's Principle
What is Le Chatelier’s principle?
Le Chatelier’s principle is an explanation of how systems in dynamic equilibrium respond to changing conditions. It states that if the conditions in a closed system change, the position of the equilibrium will shift to counteract the change.
What is an example of Le Chatelier’s principle?
An example of Le Chatelier’s principle is the Haber process, used to make ammonia. If we increase the pressure, this favours the forward reaction and increases the yield of ammonia.
What is Le Chatelier's principle and why is it important?
Le Chatelier’s principle is an explanation of how systems in dynamic equilibrium respond to changing conditions. It states that if the conditions in a closed system change, the position of the equilibrium will shift to counteract the change. It’s important because it allows us to manipulate the conditions of an equilibrium reaction in order to increase or decrease the yield.
How do you solve Le Chatelier's principle problems?
To solve problems involving Le Chatelier’s principle, you need to consider the effect of the change on the equilibrium. Ask yourself the following questions - which reaction is exothermic? Which reaction produces the greatest number of moles of gas? Remember that according to Le Chatelier’s principle, the position of the equilibrium will always shift to counteract the change in conditions. For example, if you increase the temperature of the system, the endothermic reaction will be favoured to take in some of the excess heat. If you increase the pressure, the reaction that produces the fewest moles of gas will be favoured. Analysing the equilibrium in this way should help you solve Le Chatelier’s principle problems.
What happens when you decrease the pressure according to Le Chatelier's principle?
Decreasing the pressure of a system at equilibrium favours the reaction that produces the fewest moles of gas. This is because one mole of any gas always takes up the same volume at a given temperature and pressure, so reducing the number of moles of gas reduces the overall pressure.
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