Le Chatelier's Principle
There are three major factors that alter the state of equilibrium. They are concentration, temperature and pressure. The addition of a catalyst has no effect on the state of equilibrium. Its presence merely hastens the approach of the equilibrium.
Le Chatelier's Principle
According to this principle, if a system at equilibrium is subjected to a disturbance or stress, then the equilibrium shifts in the direction that tends to nullify the effect of the disturbance or stress. Let us consider the effects of changes in temperature, concentration and pressure, on the equilibrium reactions and the predictions of Le Chatelier's principle.
Effect of change of concentration
Consider the following equilibrium reaction
N2(g) + O2(g) -- > < -- 2NO(g)
At the equilibrium conditions the reaction mixture contains both the reactant and product molecules, that is, N2, O2 and NO molecules. The concentrations of reactant and product molecules are constant and remain the same as long as the equilibrium conditions are maintained the same. If a change is imposed on the system by purposely adding NO into the reaction mixture then the product concentration is raised. Since the system possesses equilibrium concentrations of reactants and products, the excess amount of NO react in the reverse direction to produce back the reactants and this results in the increase in concentrations of N2 and O2. Similarly if the concentration of reactants such as N2 and O2 are purposely raised when the system is already in the state of equilibrium, the excess concentrations of N2 and O2 favour forward reaction. Concentration of NO is raised in the reaction mixture.
In general, in a chemical equilibrium increasing the concentrations of the
reactantsresultsin shifting the equilibriumin favour of the productswhile increasing the concentrations of the products results in shifting the equilibrium in favour of the reactants.
Effect of change of temperature
A chemical equilibrium actually involves two opposing reactions. One favouring the formation of products and the other favouring the formation of reactants. If the forward reaction in a chemical equilibrium is endothermic (accompanied by absorption of heat) then the reverse reaction is exothermic (accompanied by evolution of heat).
Let us consider the example
N2O4(g) -- > < -- 2NO2(g) ; DH = +59.0 kJ/mole
In this equilibrium, the reaction of the product formation (NO2) is endothermic in nature and therefore, the reverse reaction of reactant formation (N2O4) should be exothermic. If the above equilibrium reaction mixture is heated then its temperature will be raised. According to Le Chatelier's principle, the equilibrium will shift in the direction which tends to undo the effect of heat. Therefore,the equilibrium will shift towards the formation of NO2 and subsequently dissociation of N2O4 increases. Therefore, generally, when the temperature is raised in a chemical equilibrium, among the forward and reverse reactions, the more endothermic reaction will be favoured. Similarly, if the temperature of the
equilibrium is decreased i.e., cooled, then the exothermic reaction among the forward and reverse reaction of the equilibrium will be favoured.
Effect of change of pressure
If a system in equilibrium consists of reactants and products in gaseous state, then the concentrations of all components can be altered by changing the total pressure of the system. Consider the equilibrium in the gaseous state such as
N2O4(g) -- > < -- 2NO2(g)
Increase in the total pressure of the system in equilibrium will decrease the volume proportionately. According to Le Chatlier's principle, the change can be counteracted by shifting the equilibrium towards decreasing the moles of products.
Hence, the reaction of combination of NO2 molecules to N2O4 formation will be favoured.
In case of a gas phase equilibrium which is accompanied by decrease in number of moles of products formed, the effect of pressure can be considered as follows,
N2(g) + 3H2(g) -- > < -- 2NH3(g)
If the pressure is increased then the volume will decrease proportionately.
Consequently, the equilibrium will shift in the direction in which there is a decrease in the total number of moles, ie., favours the formation reaction of NH3. Here from four moles of reactants two moles of NH3 are formed. Thus at higher
pressures, the yield of ammonia will be more.
Ammonia is mainly used as a source of nitrogen fertiliser, in nitric acid production and in nitrogen containing pharmaceuticals. Ammonia is commercially produced in industries from the gaseous elements nitrogen and hydrogen in air by means of Haber's process. Ammonia formation reaction is an equilibrium reaction.
N2(g) + 3H2(g) -- Fe-- > < -- Fe -- 2NH3(g) DH0f = -22.0 kcal/mole
The forward reaction is accompanied by decrease in the number of moles of reactants and according to Le Chatlier's principle, an increase in pressure favours such a reaction and shifts the equilibrium towards the product formation direction. Therefore, nearly 300-500 atm pressure is applied on 3:1 mole ratio of H2:N2 gas mixture in the reaction chamber for maximum yield of ammonia. The ammonia formationreactionisexothermic.ByLeChatlier'sprinciple,increaseintemperature favours decomposition reaction of ammonia. However, at low temperature the time to reach the equilibrium becomes very long. Hence an optimum temperature close to 500 o C-550 o C is maintained. Iron catalyst is chosen to speed up the attainment of the equilibrium concentration of ammonia. In order to maintain the equilibrium conditions, steam is passed to remove away the ammonia as and when it is formed so that the equilibrium remains shifted towards the product side. The maximum yield of ammonia is nearly 37%.
This process involves the equilibrium reaction of oxidation of SO2 gas by gaseous oxygen in air to manufacture large quantities of SO3 gas.
2SO2(g) + O2(g) -- v2o5 -- > < -- v2o5 -- 2SO3(g) DH0f = -47 kcal/mole
The formation reaction of SO3 involves a decrease in the overall moles of the reactants. By Le Chatlier's principle, when large pressure is applied, forward reaction is favoured. 700 atm - 1200 atm pressure is maintained on the 2:1 mole ratio mixture of pure SO2 and O2 gases in the reaction chamber. SO3 production is an exothermic reaction. Hence, increase in temperature favours SO3 dissociation. However, lowering of temperature prolongs the time of attainment of equilibrium. Therefore, an optimum temperature at nearly 400 o C to 450 o C is maintained to favour the equilibrium.
The most widely used catalyst for SO3 production is porous vanadium pentoxide (V2O5). Presence of moisture deactivates the catalyst. Only dry and pure SO2 and O2 gases are used over the catalyst. Since oxidation of SO2 is a slow process, presence of V2O5 speeds up the equilibrium process and high yield of SO3 is achieved in a short period. SO3 is the anhydride of H2SO4.
Therefore, SO3 from contact process along with steam is used in oleum and H2SO4 manufacturing processes in contact process, the yield of SO3 is nearly 97%.
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