Equilibrium in Chemical Processes: The Dynamic Nature of Equilibrium

The same dynamic character seen in physical equilibria carries over, unchanged in principle, to

chemical reactions. When a reaction is carried out in a closed vessel, it frequently does not run to

completion: instead of every last trace of reactant being converted to product, the system settles

into a state where the concentrations of all reactants and products stop changing with time. This

state is called chemical equilibrium.

To see why this happens, imagine starting with only the reactants of a reversible reaction, written

$A \rightleftharpoons B$. At the very first instant, the forward reaction $A \to B$ proceeds at its

maximum rate (concentration of $A$ is highest) while the reverse reaction $B \to A$ cannot occur at

all (no $B$ exists yet). As the reaction proceeds, $A$ is consumed, so the forward rate falls; at the

same time $B$ accumulates, so the reverse rate rises from zero. Eventually the two rates become

exactly equal. From that moment on, $A$ is converted to $B$ just as fast as $B$ is converted back to

$A$, so the net concentrations of $A$ and $B$ no longer change — this is chemical equilibrium.

The word "dynamic" is essential and is not a mere figure of speech. At equilibrium, individual

molecules of $A$ are still being transformed into $B$, and individual molecules of $B$ are still

being transformed back into $A$, continuously and simultaneously — it is only the net, bulk

concentrations that appear frozen. This was demonstrated experimentally using isotopically labelled

atoms: if a reaction such as $\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g)$ is

allowed to reach equilibrium and then a small amount of radioactively tagged iodine is introduced,

the label is soon found distributed throughout both the remaining $\text{I}_2$ molecules and the

$\text{HI}$ molecules, even though the total, measured concentrations of $\text{H}_2$, $\text{I}_2$

and $\text{HI}$ never change. This could only happen if the forward and reverse reactions were both

still actively occurring.

Chemical equilibrium can be approached from either direction. Starting with pure reactants and

letting the forward reaction proceed, or starting with pure products and letting the reverse reaction …