When we write a balanced chemical equation of the form 2NO2 → N2O4, the meaning seems clear: Two molecules of NO2 come together in a synthesis reaction to yield a molecule of N2O4. From our discussions earlier on stoichiometry, we know that we can also look at it as 2 moles of NO2 coming together to form one mole of N2O4, or as 5 moles of NO2 coming together to form two and a half moles of N2O4, et cetera. In real life, however, if we put a certain amount of gaseous NO2 in a vessel, we will most likely not end up with half that number of moles of N2O4; in other words, the reaction is not seen to go to completion. Instead, what we would have is a mixture of both gases. What we have failed to take into account is that just as two molecules of nitrogen dioxide can combine to form a molecule of N2O4, a molecule of N2O4 can also undergo decomposition to give back two molecules of nitrogen dioxide. In the beginning, because there is no N2O4 around, only the synthesis reaction takes place; however, as the product of this reaction, N2O4, accumulates, the decomposition reaction starts to “kick in” and works to undo what the combination reaction has done. In general, for every reaction there is a reverse reaction that takes place simultaneously in opposition to it. In the long run, a state is eventually reached where the two reactions, while still going on, reach a stalemate so that no one side is gaining any net ground. From a macroscopic perspective (to our naked eye, if you will) no change is occurring in the composition of the system. This state is known as a dynamic equilibrium and will persist until the surrounding conditions are disturbed.