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Le Chatelier’s principle is concerned with thermodynamic equilibrium states. The latter are stable in the presence of perturbations that meet specific requirements; this is critical in defining thermodynamic equilibrium.
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It says that changes in a system’s temperature, pressure, volume, or concentration cause predictable and opposing changes in the system in order to reach a new equilibrium state.
To accomplish this, a state of thermodynamic equilibrium is most conveniently described by a fundamental relation that specifies a cardinal function of state, either of the energy or entropy variety, as a function of state variables chosen to correspond to the thermodynamic operations to be perturbed.
In theory, and almost always in some practical scenarios, a body can be in a stationary state with zero macroscopic flows and rates of chemical reaction (for example, when no suitable catalyst is present), but still be out of thermodynamic equilibrium due to being metastable or unstable; in this case, Le Chatelier’s principle does not always apply.
Le Chatelier’s Principles
Equilibrium is a state in which the reactants and products of a chemical process are in equilibrium. Changes in the system’s conditions can upset its equilibrium. When this occurs, the system reacts to re-establish balance. However, the equilibrium position will be altered as a result of the disturbance. In other words, the system responds by altering the concentrations of reactants and products. Some will gain in value while others decline in value until equilibrium is restored.
Henri Le Châtelier (1850–1936), a French scientist, investigated chemical equilibrium, and his explanation of how an equilibrium system responds to a change in conditions became known as Le Châtelier’s principle. When a chemical system is in equilibrium and is perturbed by a stress, this principle states that the system will respond by seeking to restore equilibrium until a new equilibrium is achieved. Stresses on a chemical system can take the form of changes in the concentrations of reactants or products, changes in the system’s temperature, or changes in the system’s pressure. Each of these pressures will be discussed separately. In each case, the change in the equilibrium position favours one of the reactions over the other. When the forward reaction is preferred, the product concentrations grow while the reactant concentrations fall. When the reverse reaction is preferred, the product concentrations decrease while the reactant concentrations increase.
Concentration
Typically, a change in the concentration of one of the reactants or products in an equilibrium system includes the addition or removal of one of the reactants or products. Consider the Haber-Bosch process for manufacturing ammonia industrially from nitrogen and hydrogen gases:
N2(g)+3H2(g)⇌2NH3(g)
When a system’s concentration of a single substance is increased, the system responds by favouring the reaction that consumes that material. When additional N2 is introduced, the forward reaction is preferred because it consumes N2 and transforms it to NH3. Initially, the forward reaction accelerates due to the presence of a larger concentration of one of the reactants, but the reverse reaction remains unchanged. Because the two rates are no longer equal, the system is no longer in equilibrium, and a net shift to the right (generating more NH3) will occur until the two rates are restored to equilibrium. NH3 concentrations increase as N2 and H2 concentrations drop. After a period of time, balance is restored by adjusting the concentrations of all three components. As illustrated in Figure below, the new concentration of NH3 is greater than the initial concentration, as the forward reaction became briefly favoured as a result of the stress. H2 is now at a lower concentration. The final concentration of N2 is greater than it was in the initial equilibrium, but less than it was immediately following the addition of N2. By responding in this manner, the value of the reaction’s equilibrium constant, Keq, remains unchanged as a result of the system’s stress.
In contrast, if additional NH3 is supplied, the reverse reaction is preferred. This “favouring” of a reaction refers to temporarily accelerating it in that direction until equilibrium is restored. Recall that after equilibrium is restored, the rates of forward and reverse reactions are restored to their initial values. By adding NH3, the production of the reactants N2 and H2 would be increased by a factor of two.
The Haber-Bosch process is one in which the reactants (N2 and H2) and the product are in equilibrium (NH3). When more N2 is introduced, the system begins to favour the forward reaction until equilibrium is restored.
Additionally, the whole or partial removal of a reactant or product might cause an equilibrium to be upset. A substance’s concentration is reduced, and the system responds by promoting the process that will take its place. It is observed that as the reaction develops, NH3 is eliminated from the equilibrium system, resulting in an industrial Haber-Bosch process. As a result, the forward reaction is favoured, resulting in the production of greater amounts of NH3. The concentrations of N2 and H2 are decreasing in the atmosphere. When NH3 is removed in a continuous manner, the reaction will eventually be forced to complete itself until all of the reactants have been consumed. If either N2 or H2 were removed from the equilibrium system, the reverse reaction would be favoured, resulting in a reduction in the concentration of NH3 in the system.
The table below summarises the consequences of concentration variations on an equilibrium system
Temperature
When the temperature of a system at equilibrium goes up or down, it puts a lot of stress on the system. In this case, the equation for the Haber-Bosch process is written again below, but this time it is written as a thermochemical equation.
N2(g)+3H2(g)⇌2NH3(g)+91 kJ
Exothermicity is the direction of the forward reaction; the creation of NH3 generates heat. The reverse process is endothermic; heat is absorbed as NH3 decomposes into N2 and H2. When the temperature of a system increases, the direction of the reaction that absorbs heat, or the endothermic direction, becomes more favourable. In this situation, heat absorption alleviates the tension caused by the temperature increase. Increased temperature favours the reverse reaction in the Haber-Bosch process. The system’s NH3 concentration falls, while the N2 and H2 concentrations increase.
In contrast, a drop in the temperature of a system promotes the exothermic direction of the reaction that produces heat. A decrease in temperature favours the forward reaction in the Haber-Bosch process. As a result, the system’s NH3 concentration increases, while the N2 and H2 concentrations fall.
Conclusion
The equilibrium law, commonly known as Le Chatelier’s principles, is used to forecast the effect of certain changes on a system in chemical equilibrium (such as the change in temperature or pressure). The principle is named after Henry Louis Le Chatelier, a French scientist.
According to Le Chatelier, equilibrium modifies the forward and backward reactions to account for changes in the equilibrium conditions.
