Factors affecting Equilibrium: Le Chatelier’s Principle
Changes of almost any kind can make this balance to shift. A system whose state of balance has been disturbed makes necessary adjustments to restore equilibrium. When equilibrium has been restored, however, the position of equilibrium is different from its original position; that is, the number of products or reactants may have increased or decreased. Such a difference is called the shift in the position of equilibrium. The law was put forth by Henri Le Chatelier who studied shifts in the position of equilibrium that result from changing conditions. According to him, if a stress is applied to a system in dynamic equilibrium, the system changes to relieve the stress. Examples of such stress include the concentration of reactants or products, change in temperature, and change in pressure. The following examples of applications of Le Chatelier’s Principle involve reversible reactions. Therefore, the product of the forward reaction is the reactant of the reverse reaction. In short, the components to the left of the reaction arrow will be considered the reactants and the elements on the RHS of the arrow are always regarded as products.
Table Of Contents
2. Concentration and equilibrium
3. Temperature and the equilibrium
Concentration and equilibrium
Changing the amount of any reactant or product in a system at equilibrium will cause the system to shift. In most cases, the system will adjust to minimize the effects of the change. For example, H2CO3 in aqueous solution will decompose to form carbon dioxide gas and water. At equilibrium, the amount of H2CO3 is less than 2%.
Adding more CO2 disturbs the equilibrium. For instance, the added CO2 changes the ratio CO2: H2CO3. Additional CO2 leads to the formation of more carbonic acid. The system, therefore, shifts to the left to set up some of the added carbon dioxides and minimize the stress. On the other hand, the removal of carbon dioxide in the system will lead to a decrease in the ratio of CO2: H2CO3. Carbonic acid decomposes to minimize the imposed stress on the system. The removal of products is usually used to increase the yield of the desired product. As products are removed from a reaction mixture, the system continually changes to restore equilibrium by producing more products.
Temperature and the equilibrium
Increasing temperature causes the equilibrium position of a reaction to the shift in the direction that absorbs heat. The heat absorption reduces the applied temperature stress. For instance, the formation of SO3 is an example of an exothermic reaction.in which SO2 reacts with O2 to yield SO2.
Heating the reaction mixture at equilibrium pushes the equilibrium position to the left, favouring the backward reaction. As a result, the product yield of the products decreases. When you cool or remove heat, the equilibrium shifts to the right, and more product will be formed.
Pressure and the equilibrium
We should note that a change in pressure on a system affects only an equilibrium that has an unequal number of moles of gaseous reactants and products. A good example is the formation of ammonia gas A catalyst is added to speed up the rate of reaction.
Generally, the system can relieve some of the pressure by reducing the number of gas molecules. For every two moles formed, four molecules of hydrogen and nitrogen are used up. In this reaction, the equilibrium will shift for more ammonia to be formed. Furthermore, the system equilibrium position for the reaction can favour the reactants by increasing the volume of the gas. To restore, the initial pressure, the system needs more gas molecules which can be generated by decomposition of a given number of ammonia molecules. In summary, one way to favour the reactants is by lowering the pressure on the system.
Explain the effect of the change in concentration on systems at equilibrium with the help of Le- Chatelier's principle.
The effect of the change in concentration on systems at equilibrium can be explained as follows:
When the concentration of the reactant(s) is increased, the system tries to reduce its concentration by favouring the forward reaction.
When the concentration of product(s) is increased, the system tries to reduce its concentration by favouring the backward reaction.
When the reactant(s) is decreased, the system tries to increase its concentration by favouring the backward reaction.
When the concentration of the product(s) is decreased, the system tries to increase its concentration by favouring the forward reaction.