Reversible reactions and equilibriumIB MYP Chemistry: Revision notes

In many reactions the reactants change into products and that is the end of it. In a reversible reaction the products can react to make the reactants again, so the reaction can go in both directions. We show this with the symbol ⇌.

Example: heating blue hydrated copper(II) sulfate gives white anhydrous copper(II) sulfate and water. Adding water to the white powder gives the blue crystals back, and heat is given out.

CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O

The left-to-right change is the forward reaction and the right-to-left change is the reverse reaction. If the forward reaction is exothermic, the reverse reaction is endothermic.

A closed system is one where nothing can get in or out, for example a sealed flask. In a closed system a reversible reaction reaches dynamic equilibrium.

At equilibrium:

• the forward and reverse reactions are both still happening
• they happen at the same rate
• so the concentrations of reactants and products stay constant, and observable properties such as colour do not change

Example: 2NO₂(g) ⇌ N₂O₄(g). Brown NO₂ turns into colourless N₂O₄ until the colour stops changing, but molecules keep changing in both directions.

If you change the conditions of a system at equilibrium, the system responds to oppose the change. This simple idea is called Le Chatelier's principle. The position of equilibrium shifts to the left (more reactants) or to the right (more products).

• Concentration: adding more of a reactant shifts the position to the right, to use it up. Removing a product also shifts it to the right.
• Pressure (gases only): increasing pressure shifts the position towards the side with fewer gas molecules. Decreasing pressure shifts it to the side with more.
• Temperature: raising the temperature shifts the position in the endothermic direction, which absorbs heat. Lowering it favours the exothermic direction.
• Catalyst: a catalyst makes equilibrium happen faster but does not change the position of equilibrium.

The Haber process makes ammonia for fertilisers: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). The forward reaction is exothermic. Nitrogen comes from the air and hydrogen from natural gas.

Typical conditions are 450 °C, 200 atmospheres and an iron catalyst. These are compromises:

• Low temperature gives more ammonia but a slow rate, so 450 °C balances rate and yield.
• High pressure gives more ammonia (4 gas molecules become 2) but costs more and is more dangerous, so 200 atmospheres is used.
• The iron catalyst speeds up the reaction without changing the yield.

The ammonia is cooled until it liquefies and is removed, which shifts the equilibrium to the right. Unreacted nitrogen and hydrogen are recycled into the reactor.