Dynamic Equilibrium

A reversible reaction that takes place in a closed container will eventually reach a state called dynamic equilibrium. This is when the forward and reverse reactions occur at the same rate, so the concentrations of the reactants and products remain constant.

An illustration depicting two beakers with different states of gas formation. On the left, a beaker labelled "evaporation" shows red gas particles rising from a pink liquid, indicated by a blue upward arrow. On the right, a beaker labelled "equilibrium" has red gas particles both suspended in the pink liquid and rising above it, with both upward and downward blue arrows, suggesting a balance between the rising and falling of particles.

Dynamic equilibrium can only be reached in a closed system. In an open system, gases formed as products can escape and disrupt the equilibrium.

To understand how the forward and reverse reactions reach equilibrium, look at the diagram below.

A graph illustrating the relationship between 'Rate' on the vertical axis and 'Time' on the horizontal axis. The graph displays two curves: a red curve labelled 'Forward Rate' that starts high and decreases over time, and a blue curve labelled 'Reverse rate' that begins lower and gradually increases. At the point where the two curves intersect, there is a note stating "Equilibrium is reached when both rates become equal."

Initially, only the reactants are present, so the forward reaction is at its highest rate. As the reaction progresses, the concentration of the reactants decreases, causing the rate of the forward reaction to also decrease. On the other hand, the concentration of the products increases, leading to an increase in the rate of the reverse reaction.

Eventually, the rates of the forward and reverse reactions become equal, and at this point, equilibrium is reached.