chemical equilibrium
Picture a crowded room with two doors between two halls. People keep walking from hall A into hall B, and others keep walking back from B into A. After a while the count in each hall stops changing — not because anyone has stopped moving, but because just as many people cross one way as the other. The doorways are still busy; the totals are simply steady. A chemical reaction reaches the same kind of standstill, and we call it chemical equilibrium.
More precisely, chemical equilibrium is the state of a reversible reaction in which the forward reaction (reactants turning into products) and the reverse reaction (products turning back into reactants) proceed at exactly the same rate, so the amounts of every substance no longer change with time. It is reached spontaneously in a closed system at constant temperature and pressure, and it is the point where the system's Gibbs free energy is at its minimum — there is no longer any net push in either direction.
The crucial caveat is that equilibrium does not mean the reaction has stopped, and it does not mean reactants and products are present in equal amounts. Both reactions are still happening underneath; only the net change is zero. The actual mix at equilibrium can be almost all products, almost all reactants, or anything in between, depending on the particular reaction and conditions. Equilibrium is about balance of rates, not equality of quantities.
Seal some brown nitrogen dioxide gas in a tube; some of it joins up into colourless dinitrogen tetroxide. At first the colour fades, then it settles to a fixed shade and stops changing — yet molecules are still constantly forming and breaking apart. The colour is steady because the two reactions now run at the same speed.
Steady colour, busy molecules — equilibrium is a balance of rates.
Two common mistakes: thinking the reaction has stopped (it has not), and thinking equilibrium means equal amounts of reactants and products (it usually does not).