Reversible Reactions and Equilibria

What "position of equilibrium" means

• The position of equilibrium describes how much of each species is present once dynamic equilibrium has been reached
• The position lies "to the right" if products are favoured (more product, less reactant at equilibrium)
• The position lies "to the left" if reactants are favoured (less product, more reactant)
• Changing the conditions of the system shifts where this balance sits

Le Chatelier's principle

• Le Chatelier's principle states that if a closed system at equilibrium is disturbed, the equilibrium position shifts so as to oppose the change imposed on it
• Le Chatelier's principle predicts the shift direction caused by:
  • Changing the temperature
  • Changing the pressure (gases)
  • Changing the concentration of a species (not assessed numerically in IGCSE but the direction is)

Effect of temperature

• Identify whether the forward reaction is exothermic or endothermic; the reverse direction is then the opposite
• Increasing temperature shifts the equilibrium in the direction of the endothermic reaction (the system "absorbs" the extra heat to oppose the rise)
• Decreasing temperature shifts the equilibrium in the direction of the exothermic reaction (the system releases heat to oppose the fall)
• Worked example. In the Haber process N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g), the forward reaction is exothermic.
  • Raising the temperature drives the equilibrium to the left — ammonia decomposes — so the yield of NH₃ falls
  • Industry compromises by running the reaction at about 450 °C: hot enough for a fast rate, cool enough that the yield is still useful

Effect of pressure (gaseous reactions only)

• Count the number of moles of gas on each side of the equation; only species in the gaseous state count
• Increasing pressure shifts the equilibrium toward the side with fewer moles of gas — that side occupies a smaller volume, opposing the pressure rise
• Decreasing pressure shifts the equilibrium toward the side with more moles of gas
• If the two sides have the same number of moles of gas, changing the pressure has no effect on the position of equilibrium (it does still alter the rate at which equilibrium is reached)
• Worked example. In the Contact process 2 SO₂(g) + O₂(g) ⇌ 2 SO₃(g):
  • Left side has 3 mol of gas (2 SO₂ + 1 O₂); right side has 2 mol of gas (2 SO₃)
  • Raising the pressure shifts the equilibrium to the right, producing more SO₃
  • The industrial reaction runs at slightly above atmospheric pressure because the yield is already high; the extra cost of higher-pressure equipment is not worth the small gain

Effect of concentration

• Adding more of a reactant shifts the equilibrium toward the products to consume the added material — position shifts to the right
• Adding more of a product shifts the equilibrium toward the reactants — position shifts to the left
• Removing a product as it forms (e.g. condensing it out of a gas mixture) drains material from the right side, so the system makes more product to compensate — equilibrium keeps shifting to the right
• This last trick is heavily used in industry: continuously removing the product is a way to push reactions that would otherwise stop at a moderate yield

Effect of a catalyst

• A catalyst speeds up both the forward and the reverse reactions by the same factor
• It lets equilibrium be reached faster, but it does not change the position once equilibrium has been reached
• The amounts of reactant and product at equilibrium are the same with or without the catalyst
• Industrially, catalysts are still used because reaching a useful yield in minutes instead of hours saves enormous amounts of energy