Dynamic Equilibrium

When we write a balanced chemical equation of the form 2NO 2 → N 2 O 4 , the meaning seems clear:
Two molecules of NO 2 come together in a synthesis reaction to yield a molecule of N 2 O 4 . From our
discussions earlier on stoichiometry, we know that we can also look at it as 2 moles of NO 2 coming
together to form one mole of N 2 O 4 , or as 5 moles of NO 2 coming together to form two and a half
moles of N 2 O 4 , et cetera. In real life, however, if we put a certain amount of gaseous NO 2 in a vessel,
we will most likely not end up with half that number of moles of N 2 O 4 ; 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 N 2 O 4 , a molecule of N 2 O 4 can also undergo decomposition to give back two molecules
of nitrogen dioxide. In the beginning, because there is no N 2 O 4 around, only the synthesis reaction
takes place; however, as the product of this reaction, N 2 O 4 , 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.