What is the Effect of Increasing Pressure on a Reaction at Dynamic Equilibrium?

Increasing the pressure on a reaction at dynamic equilibrium shifts the equilibrium to the side with fewer moles of gas.

This principle, based on Le Chatelier's principle, explains how a system at equilibrium responds to changes in conditions. When pressure increases, the system will try to relieve that stress by favoring the reaction that produces fewer gas molecules. This reduction in gas molecules decreases the overall pressure exerted by the system.

Le Chatelier's Principle and Pressure Changes

Le Chatelier's principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. In the case of pressure changes, the "stress" is the increased pressure. The system will shift to reduce this stress.

How it Works

• Identify the Number of Gas Moles: Examine the balanced chemical equation to determine the number of moles of gaseous reactants and gaseous products. Only gaseous species are considered when assessing pressure effects.
• Compare Sides: Compare the total number of moles of gas on the reactant side to the total number of moles of gas on the product side.
• Predict the Shift:
  • If the reactant side has more moles of gas than the product side, increasing the pressure will shift the equilibrium towards the products.
  • If the product side has more moles of gas than the reactant side, increasing the pressure will shift the equilibrium towards the reactants.
  • If the number of moles of gas is the same on both sides, changing the pressure will have no effect on the equilibrium.

Examples

1. N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

   • Reactants: 1 mole N₂ + 3 moles H₂ = 4 moles of gas
   • Products: 2 moles NH₃
   • Effect: Increasing pressure shifts the equilibrium to the right (towards NH₃) because there are fewer moles of gas on the product side.

2. H₂(g) + I₂(g) ⇌ 2HI(g)

   • Reactants: 1 mole H₂ + 1 mole I₂ = 2 moles of gas
   • Products: 2 moles HI
   • Effect: Changing the pressure has no effect on the equilibrium because there are equal moles of gas on both sides.

3. CaCO₃(s) ⇌ CaO(s) + CO₂(g)

   • Reactants: Only CaCO₃ is a solid, so it doesn't count. 0 moles of gas
   • Products: Only CO₂ is a gas, so 1 mole of gas
   • Effect: Increasing pressure shifts the equilibrium to the left (towards CaCO₃) because there are fewer moles of gas on the reactant side (considering only gases).

Important Considerations

• Only Gases Matter: Pressure changes only affect reactions involving gases. Solids and liquids are considered incompressible, so their amounts do not influence the effect of pressure.
• Inert Gases: Adding an inert gas (like helium or argon) at constant volume does not change the partial pressures of the reactants and products, so it does not affect the equilibrium position. However, if the addition of an inert gas increases the volume of the system (while maintaining constant pressure of reactants and products), it can affect equilibrium.

Conclusion

In summary, increasing pressure on a system at dynamic equilibrium shifts the equilibrium to the side of the reaction with the fewest moles of gas, attempting to alleviate the pressure increase. This principle is crucial for optimizing chemical reactions, particularly in industrial processes.