Le Châtelier’s Principle: If an external stress (change in concentration, pressure, or temperature) is applied to a system at equilibrium, the system adjusts (shifts) in such a way as to partially counteract that stress and reestablish equilibrium. In other words, the equilibrium “moves” to oppose the change imposed.

Adding or removing a reactant or product:
- If you add more reactant, the reaction will shift to the right (toward products) to consume some of the added reactant and restore equilibrium.
- If you remove reactant, the equilibrium shifts to the left (toward reactants) to replace some of the removed species.
- Similarly, adding product shifts the equilibrium to the left; removing product shifts equilibrium to the right.

Example: For the general reaction a A + b B ⇌ c C + d D, if we suddenly double [A], the reaction will proceed in the forward direction (to the right) until a new set of concentrations is reached at which Kc = ([C]^c [D]^d)/([A]^a [B]^b) again. The “extra” A gets used to form more C and D.

Increasing total pressure by volume reduction: When gases are compressed (volume decreases), partial pressures of each gaseous species increase. According to Le Châtelier’s Principle, if the net mole count of gaseous species on the product side differs from that on the reactant side, the equilibrium shifts to the side with fewer moles of gas to reduce total pressure.

Example (Haber process): N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g)
- Initially, consider 1 mol N₂ and 3 mol H₂ (4 mol gas total) producing 2 mol NH₃ (2 mol gas total).
- If you suddenly increase pressure (compress the mixture), the system shifts to the side with fewer moles (2 mol of NH₃) to reduce pressure. Hence, higher pressure favors ammonia formation.

Temperature changes affect equilibrium in two distinct ways:
1. They change the value of K (the equilibrium constant) itself, because K is temperature‐dependent.
2. They cause the system initially at equilibrium to shift either forward or backward until it reaches a new equilibrium consistent with the new K.

• Endothermic versus exothermic reactions:
• For a general reaction, represent the heat term on the appropriate side:
  Exothermic reaction: heat is released in the forward direction. We write it as: A + B ⇌ C + D + heat.
  or, equivalently, A + B → C + D ΔH is negative (heat released).
  Endothermic reaction: heat is absorbed in the forward direction. We write: heat + A + B ⇌ C + D or A + B → C + D ΔH is positive (heat absorbed).
• Raising temperature: Adding heat to the system: If the forward reaction is endothermic (requires heat), adding heat shifts equilibrium to the right (to consume the added heat). K (for that reaction) increases, favoring product formation. If the forward reaction is exothermic (releases heat), adding heat shifts equilibrium to the left (to consume some of the added heat). K decreases, favoring reactants.
• Lowering temperature: Removing heat has the opposite effect: equilibrium shifts toward the exothermic side.

Catalysts increase the rate at which equilibrium is achieved by lowering the activation energy for both the forward and reverse reactions equally.

Important point: While catalysts speed up how fast equilibrium is reached, they do not change the position of the equilibrium or the value of K. The concentrations at equilibrium remain exactly the same as they would be without the catalyst; only the time to reach equilibrium is shortened.