Le Chatelier's Principle is a powerful qualitative tool used by chemists to predict how a system at equilibrium responds to external changes. By understanding these shifts, industrial chemists can optimise reaction conditions to maximise the yield of valuable products.
Key Principle
When a system in dynamic equilibrium is subjected to an external change, the system shifts its position of equilibrium in a direction that tends to oppose the change.
If a factor (such as temperature, pressure, or concentration) is changed in a system at dynamic equilibrium, the position of equilibrium shifts to counteract the change.
A reversible reaction in which all the reactants and products are in the same physical state or phase (for example, all are gases or all are dissolved in aqueous solution).
Factors Affecting the Position of Equilibrium
We can manipulate three main external factors to shift the position of a homogeneous equilibrium: concentration, pressure, and temperature. A catalyst, however, behaves differently.
1. Effect of Concentration
If you increase the concentration of a reactant, the system responds by trying to decrease it. It does this by shifting the equilibrium to the right (towards products) to consume the added reactant.
- Increase reactant concentration: Equilibrium shifts to the right to decrease reactant concentration.
- Decrease product concentration: Equilibrium shifts to the right to replace the lost product.
- Increase product concentration: Equilibrium shifts to the left to consume the added product.
Step 1: Identify the change. Adding NaSCN increases the concentration of thiocyanate ions (SCN⁻) on the reactant side.
Step 2: Apply Le Chatelier's Principle. The system will oppose the increase by consuming SCN⁻ ions.
Step 3: Deduce the shift and final observation. The position of equilibrium shifts to the right (forward direction) to consume the added reactants. Consequently, the concentration of the deep red complex [Fe(SCN)]²⁺ increases, and the solution turns a deeper red colour.
2. Effect of Pressure
Pressure only affects equilibria that involve gases and where there is a difference in the number of moles of gas between the reactant and product sides of the balanced equation.
- Increase total pressure: Equilibrium shifts to the side with fewer moles of gas to decrease the pressure.
- Decrease total pressure: Equilibrium shifts to the side with more moles of gas to increase the pressure.
If a reaction has the same number of gaseous moles on both sides, such as H₂(g) + I₂(g) ⇌ 2HI(g), changing the pressure has no effect on the position of equilibrium. Make sure to count the gas stoichiometric coefficients carefully before writing your exam answer.
3. Effect of Temperature
To predict the effect of temperature, you must know whether the forward reaction is exothermic (releases heat) or endothermic (absorbs heat):
- Increase temperature: Equilibrium shifts in the endothermic direction to absorb the added heat energy.
- Decrease temperature: Equilibrium shifts in the exothermic direction to release heat energy and counteract the cooling.
4. Effect of a Catalyst
Adding a catalyst has no effect on the position of equilibrium. A catalyst speeds up the rates of the forward and reverse reactions by the exact same factor, by providing an alternative pathway with a lower activation energy. Therefore, the concentrations remain unchanged, but the system reaches its equilibrium state much faster.
- The temperature is increased.
- The overall pressure is increased.
Part a: Temperature Increase
The forward reaction is exothermic (ΔH is negative). Increasing the temperature causes the position of equilibrium to shift to the left (the endothermic direction) to absorb the added heat energy. Therefore, the yield of SO₃ decreases.
Part b: Pressure Increase
There are 3 moles of gas on the reactant side (2SO₂ + 1O₂) and 2 moles of gas on the product side (2SO₃). Increasing the pressure causes the position of equilibrium to shift to the right (the side with fewer moles of gas) to lower the pressure. Therefore, the yield of SO₃ increases.
Industrial Compromises: Haber and Contact Processes
In industry, chemical plants must balance the thermodynamic yield of a product against the kinetic rate of the reaction and the financial/safety costs of the equipment.
The Haber Process
Ammonia is produced via the Haber process reaction:
According to Le Chatelier's Principle:
- A low temperature would give a high equilibrium yield of ammonia (exothermic forward reaction). However, low temperatures result in an extremely slow rate of reaction. A compromise temperature of 450 °C is used to achieve an acceptable yield at a reasonable rate.
- A high pressure increases both the yield (shifts to the side with fewer gas moles: 2 vs 4) and the rate of reaction. However, generating and safely containing high pressures is extremely expensive and hazardous. A compromise pressure of 200 atm is used.
- An iron catalyst is used to increase the rate, enabling the plant to run efficiently at the compromise temperature of 450 °C.
When discussing compromise conditions, you must mention both the thermodynamic effect (equilibrium yield) and the kinetic effect (rate of reaction). For example: "A lower temperature increases the yield of ammonia because the forward reaction is exothermic, but decreases the rate. Therefore, 450 °C is a compromise temperature that provides a sufficient rate while maintaining a viable yield."
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