Le Chatelier's Principle Study Notes

Le Chatelier's Principle

• States that if a stress is applied to a system in equilibrium, the system will shift to relieve that stress.
• The stress moves the reaction out of its equilibrium state.
  • Concentrations or partial pressures are no longer in the equilibrium ratio.
  • The equilibrium ratio itself changes due to a temperature change.
• The reaction responds by reacting in the direction (forward or reverse) that re-establishes equilibrium.

Changes in Concentration

• Adding or removing reactants or products moves the reaction from its minimum energy state.
• The ratio of products to reactants is no longer equal to the equilibrium ratio, meaning $Q<em>c \neq K</em>{eq}$.
• If reactants are added or products are removed, $Q<em>c < K</em>{eq}$, and the reaction spontaneously reacts in the forward direction until $Q<em>c = K</em>{eq}$.
• If reactants are removed or products are added, $Q<em>c > K</em>{eq}$, and the reaction spontaneously reacts in the reverse direction until $Q<em>c = K</em>{eq}$.
• The system reacts away from the added species or toward the removed species.
• Le Chatelier's principle can be used to improve the yield of chemical reactions.
• In industrial production, products are removed as they form to prevent the reaction from reaching equilibrium, driving the reaction forward.
• Starting with high concentrations of reactants can also drive the reaction forward, increasing the absolute quantities of products.

Changes in Pressure and Volume

• Only reactions involving gaseous species are affected by changes in pressure and volume because liquids and solids are essentially incompressible.
• Compression decreases volume and increases total pressure.
• Increase in total pressure increases the partial pressures of each gas, resulting in $Q<em>p \neq K</em>{eq}$.
• The system moves toward the side with the lower total number of moles of gas.
• Ideal gas law: there is a direct relationship between the number of moles of gas and the pressure of the gas.
• Increasing pressure causes the system to decrease the total number of gas moles, decreasing the pressure.
• Expansion increases volume and decreases total pressure and partial pressures.
• The system reacts toward the side with the greater number of moles of gas to restore the pressure.
• Example:
  • $N<em>2(g) + 3H</em>2(g) \rightleftharpoons 2NH_3(g)$
  • Left side: 4 moles of gas
  • Right side: 2 moles of gas
  • Increased pressure: reaction shifts to the right, forming more ammonia.
  • Decreased pressure: reaction shifts to the left, reforming nitrogen and hydrogen gas.

Changes in Temperature

• Changing temperature causes the system to react to return to equilibrium.
• Unlike concentration or pressure changes, temperature changes result in a change in $K<em>{eq}$, not $Q</em>c$ or $Q_p$.
• $Q$ immediately after the temperature change is the same as before the change; thus, $Q \neq K_{eq}$.
• The system moves in the direction that allows it to reach its new equilibrium state at the new temperature.
• The direction is determined by the enthalpy of the reaction, $\Delta H$.
  • Endothermic reaction: \Delta H > 0, heat functions as a reactant.
  • Exothermic reaction: \Delta H < 0, heat functions as a product.
• Think of heat as a reactant or product.
• Example with an endothermic reaction:
  • $N<em>2O</em>4(g) \rightleftharpoons 2NO_2(g)$
  • Adding heat and increasing the temperature shifts the reaction to the right, increasing $NO_2$ (reddish brown).
  • Removing heat and decreasing the temperature shifts the reaction to the left, increasing $N<em>2O</em>4$ (more transparent).
• The equilibrium shifts in the direction that consumes energy.