Reversible reactions and dynamic equilibrium are important concepts in chemistry. Le Chatelier's Principle explains how systems in dynamic equilibrium respond to changing conditions. It states that when the conditions of a closed system change, the equilibrium will shift to counteract the change. This article will explain Le Chatelier's Principle in detail, including how external factors affect equilibria in accordance with the principle.
Ammonia, NH3, is a key ingredient in many fertilisers, synthetic fibres, and plastics. Under normal atmospheric conditions, the yield of ammonia is low. However, using Le Chatelier’s Principle, we can increase the yield of ammonia dramatically by changing the conditions of the reaction.
Le Chatelier's Principle can be used to predict the direction of a chemical reaction in response to changes in temperature, concentration, volume, or pressure. It can also be used to explain the effect of temperature on equilibria, as the equilibrium constant is dependent on temperature.
Common reversible reactions used in industry can also be explained using Le Chatelier’s Principle. For example, if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium will shift to counteract the change and reestablish an equilibrium.
Le Chatelier's Principle is an important concept in chemistry and can be used to understand how systems in dynamic equilibrium respond to changing conditions. It can also be used to predict the direction of a chemical reaction in response to a change in conditions, as well as to explain the effect of temperature on equilibria.
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 are special reactions that can go both ways, meaning that the products can react together to form the original reactants again. (If you want to learn more, check out Chemical Equilibrium.)
There are some key terms to understand when it comes to reversible reactions.
The reaction that goes from reactants to products is called the forward reaction. Conversely, the reaction that goes from products to reactants is called the backward reaction. If there is more of the forward reaction, we say that the forward reaction is favored and the equilibrium has shifted to the right. On the other hand, if there is more of the backward reaction, we say that the backward reaction is favored and the equilibrium has shifted to the left.
During a reversible reaction, the concentrations of reactants and products are constantly changing. However, if you leave the reaction to proceed in a closed system, the concentrations will eventually level off. At this point, the reaction has reached dynamic equilibrium.
Dynamic equilibrium is a special state of a reversible reaction where the concentrations of products and reactants remain constant, and the rates of the forward and backward reactions are equal.
Chemical equilibria are examples of dynamic equilibria. This means that both the forward and backward reactions are constantly occurring, but they happen at the same rate, effectively canceling each other out. As a result, it appears as if there is no overall reaction taking place. In a dynamic equilibrium, the products and reactants are continuously being broken down and reformed, but the overall levels of each substance remain constant.
Le Chatelier's principle is a concept that explains how equilibria respond to changes in their environment or conditions. It states that if an equilibrium is disturbed, such as by modifying its environment or conditions, the position of the equilibrium will shift in favour of the direction that opposes the change and reduces the disturbance.
For instance, imagine you are managing the seating arrangements for a restaurant. The restaurant has a limited number of seats, and at busy times, there may be a queue of people waiting outside. However, you want to maximize profits by filling up as many seats as possible. At 2 o'clock, a large party finishes, and many people leave the restaurant at once. Suddenly, there are many empty seats in the restaurant. To fill them up, you open the doors and let more people in at once. You've responded to the "disturbance" caused by many people leaving at once by allowing more people in to counteract the change the disturbance caused.
There are several ways to disturb an equilibrium, including changing the temperature, altering the concentration of reactants or products, and modifying the pressure. Le Chatelier's principle states that the equilibrium will shift in favour of the direction that opposes the change and reduces the disturbance caused by the change.
According to Le Chatelier's principle, if the temperature of a dynamic equilibrium is increased, the system will try to oppose the change by favouring the endothermic reaction, which absorbs heat as energy. In contrast, if the temperature is decreased, the system will favour the exothermic reaction, which releases heat.
For example, consider 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. If the temperature is increased, the equilibrium will shift to the left and favour the endothermic backward reaction, which will absorb heat and decrease the temperature. If the temperature is decreased, the equilibrium will shift to the right and favour the exothermic forward reaction, which will release heat and increase the temperature.
Therefore, the direction of the equilibrium shift depends on the nature of the reaction and the direction in which the temperature is changed.
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.
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.
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.
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 can be 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.
It is important to note that while a catalyst does not affect the position of equilibrium, it can still have an impact on the equilibrium reaction. For example, if the forward reaction is slow and the backward reaction is fast, a catalyst that speeds up the forward reaction will allow the equilibrium to be reached more quickly. This means that the position of equilibrium will be achieved faster, even though the catalyst did not shift the equilibrium position.
In summary, Le Chatelier's principle is a useful tool for predicting how an equilibrium reaction will respond to changes in conditions such as temperature, pressure, and concentration. Catalysts do not affect the position of equilibrium, but they can speed up the time it takes for the system to reach dynamic equilibrium. Understanding these principles is crucial for controlling chemical reactions and optimizing their efficiency.
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.
