Chemical Equilibrium and Le Châtelier’s Principle

Chemical Equilibrium and How Systems Respond to Change

What Is Chemical Equilibrium?

• In a chemical reaction, reactants are converted into products.
• Reversible reactions: Products can react to reform the original reactants.
• Chemical equilibrium: The rate of the forward reaction (reactants → products) equals the rate of the reverse reaction (products → reactants).
• Dynamic balance: Both reactions continue, but no net change occurs in concentrations of substances.
• Example: Synthesis of ammonia: N2(g) + 3H2(g) ightharpoons 2NH_3(g)
  • At equilibrium, nitrogen, hydrogen, and ammonia are all present and their concentrations remain constant over time.
• Conditions for Equilibrium:
  • The reaction must be reversible.
  • It must occur in a closed system (no substances entering or leaving).
  • The temperature and pressure must be stable.
  • The forward and reverse reactions must occur at equal rates.

Rate of Reaction vs. Concentration at Equilibrium

• Rate: How fast the reactants turn into products and vice versa.
• Concentration: How much of each substance is present at a given time.
• At equilibrium:
  • The rates of the forward and reverse reactions are equal.
  • The concentrations of the substances are constant, but not necessarily equal.
• Important Note: Equilibrium does not mean equal amounts of reactants and products, only that they are stable.

Le Châtelier’s Principle: How Systems Respond to Change

• When a system at equilibrium experiences a disturbance (change in concentration, pressure, or temperature), the system will shift to oppose the change and restore equilibrium.

Effect of Concentration Changes

• Adding more of a reactant or product: the system shifts away from the added substance to use it up.
• Removing a substance: the system shifts toward the removed substance to replace it.
• Example: Fe^{3+}(aq) + SCN^−(aq) ightharpoons FeSCN^{2+}(aq)
  • Adding more $SCN^−$: the system shifts right to form more product.
  • Removing $FeSCN^{2+}$: the system shifts right to make more.

Effect of Pressure Changes (Gases Only)

• Changes in pressure only affect reactions involving gases, and only when the number of gas particles is different on each side of the reaction.
• Increasing pressure: the system shifts toward the side with fewer gas molecules.
• Decreasing pressure: the system shifts toward the side with more gas molecules.
• Example: N2(g) + 3H2(g) ightharpoons 2NH_3(g)
  • Left side: 4 moles of gas
  • Right side: 2 moles of gas
  • If pressure increases, the system shifts right (fewer moles).
  • If pressure decreases, it shifts left (more moles).
• When Pressure Has No Effect:
  • If the number of gas molecules is equal on both sides, changing pressure does not shift the equilibrium.

Effect of Temperature Changes

• Temperature changes shift equilibrium based on whether the reaction is exothermic or endothermic.
• Exothermic reaction (releases heat): heat is treated like a product.
• Endothermic reaction (absorbs heat): heat is treated like a reactant.
• Adding heat: The system shifts away from heat.
• Removing heat: The system shifts toward heat.
• Example (Exothermic): N2(g) + 3H2(g) ightharpoons 2NH_3(g)
  • $\Delta$ = -56kJ
  • Adding heat → shift left (toward reactants)
  • Removing heat → shift right (toward products)

Summary of Direction of Shifts

Final Thoughts:

• Understanding chemical equilibrium and how it responds to stress is essential for predicting how reactions behave in real-world systems—from industrial processes to biological systems.
• By applying Le Châtelier’s Principle, we can manipulate conditions to maximize product yield, control reactions, and design effective chemical systems.