Chemical equilibrium

Many reactions do not run to completion. Products re-form reactants, and as product builds up that reverse reaction speeds up while the forward one slows. Eventually the two rates match and the composition stops changing. That is equilibrium — and the word "dynamic" attached to it is doing real work: nothing has stopped. Molecules are still reacting in both directions, at identical rates, which is why the concentrations hold steady.

Steady does not mean equal. The equilibrium constant K fixes the ratio of products to reactants, and that ratio can be enormous or tiny. A reaction with K = 10⁶ sits almost entirely on the product side at equilibrium; one with K = 10⁻⁶ has barely proceeded. Equal concentrations of reactants and products is one possible equilibrium among many, not the definition of one.

Disturb a system at equilibrium and it shifts to partly offset the change — Le Chatelier's principle. Add reactant and it consumes some; remove product and it makes more; raise the pressure on a gas mixture and it shifts towards whichever side has fewer molecules. This is what makes industrial synthesis controllable: the Haber process runs at high pressure precisely because ammonia occupies fewer molecules than the nitrogen and hydrogen it comes from. A catalyst does none of this. It lowers the activation barrier for both directions by the same amount, so equilibrium arrives faster but lands in exactly the same place.