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Le Chatelier’s Principle

Le Chatelier’s principle

When a system at equilibrium is subject to a change, the system will shift to partially counteract the change.

For example, for the following reaction:

In a 1.00 \(L\) reaction vessel at a fixed temperature and pressure, if equal mole amounts of \(A\) and \(B\) are added, they will begin to react to form \(C\). However, the reaction vessel will never contain only \(C\). This is because once \(C\) is formed, it begins to react in the reverse reaction, forming \(A\) and \(B\).

\[\text{A(g)} + \text{B(g)} \rightleftharpoons 2\text{C(g)} \quad \quad \Delta\text{H} = -\text{ve}\]

Once the system has reached equilibrium, both the forward and reverse reactions are taking place at an equal rate. If any changes are made, this results in the system shifting to partially counteract this change.

Worked example

\[\text{A(g)} + \text{B(g)} \rightleftharpoons 2\text{C(g)} \quad \quad \Delta\text{H} = -\text{ve}\] For example, if more reactant \(\text{A}\) is added, the system can respond to this increase in concentration of \(\text{A}\), by moving in the forward direction, partially counteracting the increase by decreasing the concentration of \(\text{A}\). In doing so, more of product \(\text{C}\) will be produced. This means the yield of \(\text{C}\) is increased. This is a way that chemists can adjust a reaction mixture to increase the yield of a desired product.

Use this page to revise the following concepts within Le Chatelier’s Principle:

Predicting the effect of changes to equilibrium systems
Factors that change the position of equilibrium
Concentration-time graphs

Predicting the effect of changes to equilibrium systems

As mentioned above, the system will shift away from the side of a reaction that has had a reactant or product added to it (in either liquid or gas equilibrium systems).
For the reaction
\[\text{A(g)} + \text{B(g)} \rightleftharpoons 2 \text{C(g)}\]
increasing the concentration of \(A\) or \(B\) will shift the reaction forwards, while adding \(C\) will shift the reaction in reverse. Conversely, the opposite is true if we take away species. If we remove \(C\), the system will shift in the forward direction, to generate more \(C\) to partially counteract the decrease. If we take away \(A\) or \(B\), the system will shift in the reverse direction, decreasing the yield of \(C\).
For a gaseous system, changes in pressure results in the equilibrium shifting in a direction opposite to the change. The pressure of a gas depends on the number of particles per volume, so the equilibrium will shift to increase or decrease the number of particles to counteract a change in pressure. For example, increasing the pressure would shift the direction to the side with the least number of particles, as this will decrease overall pressure, partially opposing the original increase.
For the reaction
\[\text{A(g)} + \text{B(g)} \rightleftharpoons \text{C(g)}\]
increasing the pressure will shift the reaction forwards to the side with only one particle, while decreasing the pressure will shift the reaction in reverse.
For an aqueous solution, we cannot change the pressure as liquids are incompressible. However, we can add water, and thus dilute the solution. Dilution causes the reaction to shift to the side with more particles, to oppose the decrease in concentration. Similarly, removing water, making the solution more concentrated, can cause a shift to the side with less particles.
For the reaction
\[\text{A(aq)} + \text{B(aq)} \rightleftharpoons \text{C(aq)}\]
adding water will dilute this system, and favour the reverse reaction, to the side with more total particles, to partially offset this change.
When we look at reactions that are exothermic or endothermic, we can consider energy to act as a reactant or product. For example, an exothermic reaction produces heat energy, while an endothermic reaction consumes it.
In this way, increasing the temperature of a reaction will cause an endothermic reaction to shift forwards, and partially counteract the increased energy. Conversely, it will cause an exothermic reaction to shift in the reverse direction, and partially counteract the increased energy.
For the reaction
\[\text{R(g)} + \text{S(g)} \rightleftharpoons \text{Q(g)} \quad \quad \Delta\text{H} = -\text{ve} \]
increasing temperature would favour the reverse reaction, to partially oppose this energy change. A decrease in temperature would favour the forward reaction.
If the reaction were endothermic, increasing temperature would favour the forward reaction, and decreasing temperature would favour the backward reaction.

Factors that change the position of equilibrium

Temperature is the only factor that also changes the value of the equilibrium constant, \(\text{K}\). For an exothermic reaction, increasing temperature causes a decrease in \(\text{K}\) (as the reverse reaction is favoured, and products decrease). For an endothermic reaction, increasing temperature causes an increase in \(\text{K}\) (as the forward reaction is favoured, and products increase).

A catalyst has no effect on yield, as it only allows rate to increase, and does not affect the position of equilibrium.

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Concentration-time graphs

Concentration-time graphs show equilibrium systems as they are established. An example is shown below.

Once we understand what happens to the concentrations of products and reactants at equilibrium, we can easily see how changes to these can lead to equilibrium systems responding by moving in the forward or reverse directions, according to Le Chatelier’s principle.

As mentioned above, the system will shift away from the side of a reaction that has had a reactant or product added to it. Since \(\text{CO}\) is a reactant, the system will shift forward. This will decrease the concentration of both reactants and increase the concentration of products. However, it will not decrease the \(\text{CO}\) past its initial value, only partially offset the change.
For a gaseous system, changing the pressure results in the system shifting in a direction opposite to the change. Increasing the pressure by decreasing the volume leads to an immediate increase in concentration of all species. The system will then shift the direction to the side with the least number of particles, as this will decrease overall pressure, partially opposing the original increase.
For an aqueous solution, adding water dilutes the solution, immediately decreasing the concentration of all species. The system will then shift to the side with more particles to partially offset this change.
When we look at reactions that are exothermic or endothermic, we can consider energy to act as a reactant or product. Decreasing the temperature of an exothermic reaction will favour a forward reaction, decreasing reactant concentrations and increasing product concentrations.