le-Chatelier principle
Le Chatelier’s principle explains how equilibrium changes in response to factors such as concentration and pressure. It states that if an equilibrium is subjected to any change, it will shift in the forward or backward direction to counteract that change. This shift occurs in order to restore balance within the system. Changes can include variations in concentration, pressure, temperature, or the addition of catalysts and inert gases. This principle has more applications in industry. An equilibrium system can be shifted to favor the formation of product. Haber process used to synthesize NH3 industrially. Haber optimized the required conditions to get the maximum yield of ammonia.
Effect of concentration change
If the equilibrium is disturbed by the addition of reactant or product to the equilibrium, then is shifts to either forward or backward direction in order to nullify the effect of concentration changed to maintain the same equilibrium constant. Addition of reactant or the removal of the product shifts the equilibrium forward reaction (favors the formation of product). It helps to increase the yield of product in the industry. For example in the synthesis of CaO(s) from CaCO3(s), continuous removal of one of the product, CO2(g). equilibrium shifts to product side to form more product, CaO(s). Similarly in the synthesis of NH3(g), the ammonia gas produced is condensed and removed from the equilibrium, enhances the formation of NH3 further until the reaction goes to the end. Removal of the reactant or the addition of the product shifts equilibrium to backward direction.
Let us consider the equilibrium: H2(g) + I2(g) ⇌ 2HI(g). If H2(g) added to this equilibrium, then equilibrium point is disturbed. To restore the equilibrium the added H2(g) should be consumed. Therefore, H2(g) reacts with I2(g) to form HI(g), therefore equilibrium shifts to the forward direction and new equilibrium is established.
Experiment:
Fe3+(yellow) + SCN–(Colorless)⇌ [Fe(SCN)]2+(Deep Red color). If this equilibrium is disturbed either by the removal of Fe3+(By adding oxalic acid) or by the removal of SCN – (by the addition of HgCl2) then equilibrium shifts forward direction (Can be observed through the decrease in the intensity of red color). If the oxalic acid is added, then oxalate forms a stable complex of [Fe(C2O4)3]3- which reduces the concentration of Fe3+ ions at equilibrium. Therefore, equilibrium shifts backward direction. If HgCl2 is added to the equilibrium, then Hg2+ forms a stable complex of [Hg(SCN)4]2- which decreases the concentration of SCN– ions at the equilibrium. Therefore, equilibrium shifts backward direction. If KSCN(Potassium thiocyanate) is added to the equilibrium (Increases the concentration of reactant SCN–) then equilibrium shifts to the forward direction, which can be observed through the increase in intensity of red color.
Question:
Following is the equilibrium for the formation of CH3OH(g), Predict the direction to which the given equilibrium shifts in the following cases
2H2(g) + CO(g) ⇌ CH3OH(g)
- addition of H2
- addition of CH3OH
- removal of CO
- removal of CH3OH
Ans: Addition of H2(reactant): Forward direction
Addition of CH3OH(product): Backward direction
Removal of CO(reactant): Backward direction
Removal of CH3OH(product): Forward direction.
Addition of product or Removal of reactant ⟹ Equilibrium shifts backward directionAddition of reactant or removal of product ⟹ Equilibrium shifts forward direction
Effect of pressure change
In case of equilibrium systems with Δng ≠ 0 (Number of gaseous reactants are different from the number of moles of gaseous products), if we change the pressure of the equilibrium then to counteract this change the equilibrium will shift either forward or back ward reaction. If the pressure increased then the equilibrium shifts to the direction which has less number of gaseous moles of species. Example: CO(g) + 3H2(g) ⇌ CH4(g) + H2O(g). Let us consider this equilibrium in a cylinder with a movable piston. If the gases are compressed at constant temperature to half of the volume, then the pressure becomes double(Because PV = constant at constant temperature). Now the system is not at equilibrium therefore to re-establish the equilibrium, it shifts to product side(forward direction), which has less number of moles than reactant.
If Δng < 0 ⇒ Pressure increases equilibrium shifts forward direction
Pressure decreases equilibrium shifts backward direction
If Δng > 0 ⇒ Pressure increases equilibrium shifts forward direction
Pressure decreases equilibrium shifts forward direction
If Δng = 0 ⇒ There is no effect of pressure
Question:
Does the number of moles of reaction products increase, decrease or remain same when each of the following equilibria is subjected to a decrease in pressure by increasing the volume?
(a) PCl5 (g) ⇋ PCl3 (g) + Cl2 (g)
(b) CaO (s) + CO2 (g) ⇋ CaCO3 (s)
(c) 3Fe (s) + 4H2O (g) ⇋ Fe3O4 (s) + 4H2 (g)
Ans:
If the pressure decreases, then equilibrium shifts more number of gaseous moles side and if the number moles of gaseous reactants are same as that of gaseous products there is no effect of pressure on equilibrium.
(a) Total number of gaseous products are two(1 mole of PCl3 and 1 mole of Cl2) and the number of moles of gaseous reactants is 1 (1 mole of PCl5). Therefore, equilibrium shifts product side(Forward direction).
(b) Number of moles of gaseous products = 0; Number of moles of gaseous reactants = 1. Therefore, decrease in pressure shifts reactant side (Backward reaction).
(c) Number of moles of gaseous products = 4; Number of moles of gaseous reactants = 4. Therefore, decrease in pressure does not affect the equilibrium.
Effect of temperature
Effect of temperature on the equilibrium constant depends on the sign of ∆H. Therefore, it id different for both exothermic and endothermic equilibrium. In case of exothermic reactions as the temperature increases, reaction shifts backward direction to nullify that change because the energy added to the system should be used and the backward reaction is endothermic and absorbs the added heat.
Experiment: Let us consider an equilibrium which is exothermic: 2NO2(Brown) ⇋ N2O4(Colorless) +57.2 kJmol-1. Fill two test tubes with NO2 gas and keep both the tubes for about 10 minutes in a beaker which has normal water and then one of the test tube to the container which has freezing mixture and another to the container which has hot water. The test tube in freezing mixture the brown color of NO2 becomes less intense because as the temperature decreases equilibrium shifts forward direction. The test tube in the container which has hot water the brown color of NO2 becomes more intense because as the temperature increases reaction shifts backward direction.
For exothermic reaction: ⇒As the temperature increases reaction shifts backward directionAs the temperature decreases reaction shifts forward directionFor endothermic reaction: ⇒As the temperature increased reaction shifts forward directionAs the temperature decreases reaction shifts backward direction
Effect of addition of inert gas
Addition of inert gas like Argon to the equilibrium at constant volume does not change the position of the equilibrium because inert gas does not participate in the reaction. Therefore, addition of inert gas to the equilibrium does not change the partial pressure or molar concentration of reactants and products. Therefore, reaction quotient remains constant which results there is no shifting of equilibrium either forward or backward reaction
Addition of inert gas to the equilibrium at constant volume does not disturb the equilibrium.
Effect of a catalyst
Catalyst is a substance which increases the rate of reaction by changing pathway of the reaction with lower activation energy. Catalyst lowers the activation energy of both forward and backward reactions with same extent. Therefore, catalyst do not change the equilibrium composition of both reactants and products.
Example:
Let us consider the synthesis of ammonia by the Haber process: N2 + 3H2 ⇋ 2NH3 + 92.4 kJ. This highly exothermic process shifts backward at high temperature and lowers the yield of NH3 and at low temperature rate of reaction is low and slowly it reaches to equilibrium. Therefore, a catalyst like Fe increases rate of reaction and favors the formation of ammonia. Haber optimised the reaction conditions (Fe catalyst, 500oC temperature, 200 atm of pressure) to maximize the yield of ammonia.
Addition of catalyst to the equilibrium do not change the composition of reactants or products. It increases rate of forward and backward reactions to the same extent.