Dynamic Equilibrium

Definition of Dynamic Equilibrium

Dynamic equilibrium occurs in reversible chemical reactions where the forward and backward reactions happen at exactly the same rate. This means that, although both reactions continue to occur, there is no overall change in the amounts of reactants and products. The concentrations of reactants and products remain constant over time, but the system is still active on a microscopic level.

For example, in a closed system where nitrogen dioxide ($\text{NO}_2$) decomposes into nitrogen monoxide ($\text{NO}$) and oxygen ($\text{O}_2$), the forward and backward reactions eventually balance out. At this point, the system is at dynamic equilibrium.

This is different from a static equilibrium, where nothing happens at all. In dynamic equilibrium, reactions continue but with no net change in concentrations.

Characteristics of Dynamic Equilibrium

• The system must be closed to matter, so no reactants or products can enter or leave. This is essential to maintain constant concentrations.
• Equilibrium is dynamic, meaning both forward and backward reactions continue simultaneously at the same rate.
• There is no net change in the concentrations of reactants and products, even though molecules are constantly reacting.

Because the system is closed, the amounts of substances stay the same, but the reactions keep happening. This dynamic nature is key to understanding many chemical processes.

Conditions for Dynamic Equilibrium

• A closed system is required to prevent loss or gain of substances.
• The reaction must be reversible, so products can react to form reactants again.
• There must be sufficient time for the forward and backward reactions to balance out and reach equilibrium.

If any of these conditions are not met, dynamic equilibrium cannot be established. For example, if the system is open, products may escape and the concentrations will change continuously.

Significance in Chemical Reactions

Dynamic equilibrium is important because it helps predict how reactions behave when conditions change. It forms the basis for understanding Le Chatelier’s Principle, which explains how equilibrium shifts in response to changes in concentration, pressure, or temperature (covered in other topics; see separate notes for details).

In industry, many important processes rely on dynamic equilibrium to maximise product yield, such as the manufacture of ammonia in the Haber process (see separate notes for details).

Understanding dynamic equilibrium allows chemists to control reactions better and improve efficiency in chemical manufacturing.

Example of Dynamic Equilibrium in a Closed System

Consider the reversible reaction:

$$\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$$

Initially, only reactants A and B are present. As the reaction proceeds, products C and D form. After some time, the rate of the forward reaction (A + B → C + D) equals the rate of the backward reaction (C + D → A + B). At this point, the system has reached dynamic equilibrium.

The concentrations of A, B, C, and D remain constant, but molecules continue to react in both directions.

For instance, if the rate of forward reaction is 0.05 mol/s, the backward reaction rate is also 0.05 mol/s at equilibrium.