Introduction to Equilibrium

Drop NO2 gas into a sealed flask and it starts reacting with itself: 2NO2 <→ N2O4. At first the forward reaction runs fast because there's plenty of NO2, and the reverse reaction barely runs because almost no N2O4 exists yet. As NO2 gets used up, the forward rate slows down. As N2O4 builds up, the reverse rate speeds up. Eventually the two rates meet in the middle and become equal. From that moment on, NO2 is destroyed exactly as fast as it's created, so its concentration stops changing. The flask's color stops shifting, but on the molecular level, both reactions are still firing constantly. Nothing has stopped — it's just balanced.

If you graph concentration against time, reactant lines fall and product lines rise, both curving as the rates change, until both go flat and stay flat. That flatness is the signature of equilibrium: no net change, because the two opposing processes cancel out exactly. The height each line settles at has nothing to do with the two amounts being equal to each other — a flask can settle with 90% product and 10% reactant, or the reverse, depending on how strongly the reaction favors one side. That tendency is captured by the equilibrium constant, K, which is where the unit goes next: once you accept that equilibrium is a balance of rates rather than amounts, K gives you the number that tells you exactly where that balance sits, and ICE tables let you calculate the actual resting concentrations.