// CHEMISTRY L4 — EQUILIBRIA
A reaction that never finishes.
In a sealed flask, many reactions reach a balance point — going forwards and backwards at the same rate. Poke that balance and it fights back. Run the Haber process and watch the position shift.
// FIG. 01 — THE HABER PROCESS · N₂ + 3H₂ ⇌ 2NH₃ (forward reaction is exothermic)
◀ SHIFTS LEFT = LESS AMMONIA
YIELD OF NH₃: {{ yieldPct }}%
SHIFTS RIGHT = MORE AMMONIA ▶
{{ explanation }}
// DYNAMIC EQUILIBRIUM
A reversible reaction can go both ways: ⇌. Seal it in a closed container and it reaches dynamic equilibrium — the forward and backward reactions are still happening, but at equal rates, so the amounts of everything stop changing. Nothing has stopped; it's a busy stalemate. The mixture can settle anywhere between "mostly reactants" and "mostly products" — that balance point is the position of equilibrium.
// LE CHATELIER'S PRINCIPLE
Change a condition and the equilibrium shifts to oppose that change. Add more reactant → it shifts right to use it up. Raise the pressure → it shifts toward the side with fewer gas molecules. Raise the temperature → it shifts in the endothermic direction to absorb the heat. Add a catalyst and — importantly — nothing shifts: a catalyst speeds both directions equally, so it reaches the same equilibrium faster but doesn't change the yield.
// THE INDUSTRIAL COMPROMISE
The Haber process shows the real-world tension. High pressure and low temperature give the best yield of ammonia — but low temperature makes the reaction agonisingly slow. So industry compromises: about 200 atm, 450 °C, and an iron catalyst. Unreacted gases are recycled. It's a deliberate trade of yield against rate against cost — the same balancing act behind fertiliser for half the planet's food.
// REFERENCES & FURTHER READING
Cross-check against your own exam board's specification (AQA, Edexcel, OCR, SQA).