Chemical equilibrium

Many reactions don't go all the way. As products build up, the reverse reaction speeds up, until forward and back run at the same rate. Then the amounts stop changing, though both reactions carry on: dynamic equilibrium, written with a double arrow, ⇌.

At equilibrium the concentrations always satisfy one ratio, the equilibrium constant K. For aA + bB ⇌ cC + dD, K = [C]ᶜ[D]ᵈ ÷ ([A]ᵃ[B]ᵇ): products over reactants, each raised to its coefficient. Pure solids and liquids are left out. A big K (say 10⁵) means mostly products; a small one (10⁻⁵) mostly reactants. K changes only with temperature.

Plug in the concentrations you have right now and you get the reaction quotient, Q. If Q < K there aren't enough products yet, so the reaction goes forward; if Q > K it goes backward; if Q = K it's at equilibrium.

Le Chatelier's principle says that if you disturb an equilibrium, it shifts to partly undo the change. Add a reactant or remove a product and it makes more product. For a gas reaction, raising the pressure favors the side with fewer gas molecules. Heating favors the endothermic direction. The Haber process, N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92 kJ, has 4 gas molecules on the left and 2 on the right, so ammonia plants run at high pressure (around 200 bar). Low temperature would give more ammonia but too slowly, so they compromise at about 450 °C with an iron catalyst and keep removing the ammonia.

K and the standard Gibbs energy are linked: ΔG° = −RT ln K. A negative ΔG° means K > 1. At 25 °C, every 5.7 kJ/mol of ΔG° is a factor of 10 in K, so ΔG° = −11.4 kJ/mol gives K ≈ 100.