Le Chatelier’s Principle
Effects of External Factors on Chemical Equilibrium (HSC Chemistry)
Once a reversible reaction reaches chemical equilibrium, if the external conditions (temperature, pressure, inert gas addition, and concentration) remain unchanged, the equilibrium state will continue indefinitely. Controlling temperature, pressure, inert gas addition, and concentration governs the equilibrium state. A change in any one of these factors shifts the equilibrium state, resulting in a change in the concentrations of reactants and products.
To explain the effect of changing these parameters on chemical reactions, the French scientist Henri Le Chatelier proposed a fundamental principle known as Le Chatelier’s Principle. Using this principle, the influence of temperature, pressure, inert gas addition, and concentration on chemical equilibrium can be accurately understood, which is widely utilized in industrial processes to optimize chemical reactions in minimum time with maximum yield.
Statement of Le Chatelier’s Principle
If a chemical system at equilibrium is subjected to a change in any of the equilibrium conditions (such as temperature, pressure, or concentration), the system will adjust itself in such a direction so as to nullify or minimize the effect of that change.
Effects of Temperature, Pressure, and Concentration on Chemical Equilibrium
Effect of Temperature, Pressure & Concentration on Equilibrium Constant of Reaction
1. Effect of Temperature:
The effect of temperature on a system at equilibrium is two-fold:
- a. Effect on the position of equilibrium.
- b. Effect on the equilibrium constant (Kc / Kp).
a) Effect on the Position of Equilibrium:
To evaluate this effect, we examine two types of reversible reactions:
This is an endothermic reaction, meaning heat is absorbed during the forward reaction. According to Le Chatelier’s principle, an increase in temperature favors the endothermic forward reaction to absorb the added heat and counteract the temperature rise. Consequently, increasing temperature increases the yield of NO (from N2 and O2), shifting the equilibrium position from left to right. Thus, for all endothermic reversible reactions, increasing temperature shifts the equilibrium position to the right (forward), while decreasing temperature shifts it to the left (backward).
| Applied Change | System’s Response | Direction of Shift | Yield of Product |
|---|---|---|---|
| Temperature Increase | Absorbs Heat (Lowers Temp) | Forward Direction | Increases |
| Temperature Decrease | Releases Heat (Raises Temp) | Backward Direction | Decreases |
This is an exothermic reaction, meaning heat is released during the forward process. According to Le Chatelier’s principle, an increase in temperature shifts the equilibrium position from right to left (backward) to absorb the excess heat and neutralize the temperature increase. Consequently, the yield of NH3 decreases and the amounts of reactants (N2 and H2) increase. Conversely, lowering the temperature shifts the equilibrium to the right (forward), increasing the yield of NH3. The same behavior applies to all exothermic reactions.
| Applied Change | System’s Response | Direction of Shift | Yield of Product |
|---|---|---|---|
| Temperature Increase | Absorbs Heat (Lowers Temp) | Backward Direction | Decreases |
| Temperature Decrease | Releases Heat (Raises Temp) | Forward Direction | Increases |
At low temperatures, the equilibrium mixture contains a substantial concentration of NH3, and the equilibrium constant (Kc) is high. As temperature rises, NH3 dissociates back into N2 and H2, shifting the equilibrium to the left and decreasing Kc. That is, for exothermic reactions, increasing temperature decreases the value of the equilibrium constant.
2. Effect of Pressure on Equilibrium:
Pressure has no significant effect on reactions involving only pure liquids or solids. However, for reactions taking place in the gaseous phase, pressure plays a vital role. When pressure is applied to a gaseous equilibrium mixture, according to Le Chatelier’s principle, the equilibrium shifts in the direction that relieves the applied pressure.
When external pressure is applied to a gas, the system tries to compress its volume to relieve the pressure. Therefore, applying pressure shifts the equilibrium toward the side with fewer total moles (smaller volume). However, if a reaction contains an equal number of gaseous moles on both the reactant and product sides, the gaseous volumes are equal. Consequently, pressure has no effect on such equilibrium systems. For example:
In the dissociation of HI to produce I2 and H2, the number of moles of reactants (2 moles HI) equals the total moles of products (1 mole I2 + 1 mole H2 = 2 moles). Hence, change in pressure has no effect on this equilibrium.
If a gaseous reaction results in a decrease in total moles (products have fewer moles than reactants), increasing pressure increases the product yield. For example, 1 mole of N2 and 3 moles of H2 (total 4 moles of reactants) produce 2 moles of NH3 gas:
Here, the volume of products is less than that of the reactants. According to Le Chatelier’s principle, increasing the pressure shifts the equilibrium to the right (forward) to decrease volume and relieve the pressure. Thus, the production of NH3 increases while the amounts of reactants decrease. Conversely, reducing pressure shifts the equilibrium to the left (backward), decreasing the NH3 yield. For this reason, ammonia is commercially synthesized under a high pressure of approximately 200 atm.
| Applied Change | System’s Response | Direction of Shift | Yield of Product |
|---|---|---|---|
| Pressure Increase | Decreases Pressure (Volume) | Forward Direction | Increases |
| Pressure Decrease | Increases Pressure (Volume) | Backward Direction | Decreases |
If the number of gaseous moles of products is greater than that of reactants, increasing pressure decreases the product yield. In such reactions where product volume is greater, the pressure effect is reversed. For example:
In this reaction, the gaseous products total 5 moles (4 NO2 + 1 O2) while the reactants total 2 moles. Thus, product volume is greater than reactant volume. According to Le Chatelier’s principle, increasing pressure shifts the equilibrium to the left (backward) toward the smaller volume to minimize the applied pressure. As a result, the formation of N2O5 increases and the product amounts decrease. Conversely, decreasing pressure shifts the equilibrium to the right (forward), increasing the yields of 4NO2 and O2. Therefore, this type of reaction is favored at low pressures.
| Applied Change | System’s Response | Direction of Shift | Yield of Product |
|---|---|---|---|
| Pressure Increase | Decreases Pressure (Volume) | Backward Direction | Decreases |
| Pressure Decrease | Increases Pressure (Volume) | Forward Direction | Increases |
3. Effect of Concentration:
The effect of component concentration on a reversible equilibrium system can be explained using Le Chatelier’s principle. For example, when PCl5 is heated in a closed vessel, it dissociates reversibly into PCl3 and Cl2:
In terms of molar concentrations, the equilibrium constant is: Kc = [PCl3][Cl2] [PCl5]
In this reaction, PCl5 partially dissociates, and eventually an equilibrium is established between the products (PCl3 and Cl2) and undissociated PCl5. All three components coexist in the equilibrium mixture.
If additional PCl5 is added to the mixture, the concentration of PCl5 increases. While this would alter the ratio, the equilibrium constant (Kc) remains constant at a fixed temperature. To nullify the effect of the added PCl5, the excess reactant dissociates, increasing product concentrations and shifting the equilibrium to the right (forward).
Conversely, if Cl2 is added to the equilibrium mixture, the concentration of Cl2 increases. To maintain a constant Kc, the added Cl2 reacts with PCl3 to form more PCl5, thereby neutralizing the increase in Cl2 and shifting the equilibrium to the left (backward).
Furthermore, if PCl3 or Cl2 is continuously removed from the reaction vessel, more PCl5 dissociates to replenish the deficit and restore equilibrium, shifting the position to the right. Thus, changing concentrations shifts the position of equilibrium, but the equilibrium constant (Kc) remains unchanged at a constant temperature.
| Applied Change | System’s Response | Direction of Shift | Yield of Product |
|---|---|---|---|
| Addition of PCl5 (Increase) | Consumes PCl5 (Decrease) | Forward Direction | Increases |
| Addition of Cl2 (Increase) | Consumes Cl2 (Decrease) | Backward Direction | Decreases |
Describe the effects of Temperature, Pressure, and Concentration on the following reactions:
- H2(g) + N2(g) ⇌ 2NH3(g) + Heat
- 2SO2(g) + O2(g) ⇌ 2SO3(g) + Heat
- N2(g) + O2(g) ⇌ 2NO(g) – Heat
- N2O4(g) ⇌ 2NO2(g) + Heat
- 2N2O5(g) ⇌ 4NO2(g) + O2(g) + Heat
