Removing a product from a chemical reaction shifts the equilibrium towards the product side, thereby increasing the rate of the forward reaction.

In a chemical reaction, equilibrium refers to the state in which the concentrations of reactants and products remain constant over time. It is essential to note that this does not imply that the reaction has ceased; rather, both the forward and reverse reactions occur at the same rate. This phenomenon is explained by Le Chatelier’s Principle, which states that if a dynamic equilibrium is disrupted by a change in conditions, the system will adjust to counteract that change and restore a new equilibrium.

When a product is removed from the system, the equilibrium is disturbed. According to Le Chatelier’s Principle, the system responds by attempting to replace the removed product. This adjustment is achieved by shifting the equilibrium position to the right, towards the products side. As a result, the rate of the forward reaction increases, leading to the production of more product to compensate for the loss.

For example, consider the reversible reaction in which nitrogen and hydrogen react to form ammonia:

If some of the ammonia ($\text{NH}_3$) is removed from the system, the equilibrium will shift to the right in order to produce more ammonia. This shift means that additional nitrogen and hydrogen will react, thereby increasing the rate of the forward reaction.

This principle is particularly significant in industrial processes that require continuous production. By continuously removing a product, the reaction can be maintained in the desired direction, maximizing the yield of the product. For instance, in the Haber process for ammonia production, the generated ammonia is continuously removed through cooling and liquefaction. This removal drives the reaction toward the right, enhancing the overall yield of ammonia.