How do changes in concentration affect equilibrium?

Changes in concentration can shift the equilibrium position of a reaction, either to the right (products) or left (reactants).

In a chemical reaction, the equilibrium position refers to the balance between the reactants (substances that start the reaction) and the products (substances that are produced). According to Le Chatelier's Principle, if a system at equilibrium is subjected to a change, the system will adjust itself to counteract that change. This principle applies to changes in concentration.

If the concentration of a reactant is increased, the system will try to decrease it by producing more products, thus shifting the equilibrium to the right. Conversely, if the concentration of a product is increased, the system will try to decrease it by producing more reactants, shifting the equilibrium to the left.

Similarly, if the concentration of a reactant is decreased, the system will try to increase it by converting more products into reactants, shifting the equilibrium to the left. If the concentration of a product is decreased, the system will try to increase it by converting more reactants into products, shifting the equilibrium to the right.

It's important to note that these changes in concentration do not alter the equilibrium constant, which is a measure of the ratio of the concentrations of products to reactants at equilibrium. The equilibrium constant remains the same unless the temperature changes.

In summary, changes in concentration can cause the equilibrium position of a reaction to shift in order to counteract the change, but they do not affect the equilibrium constant. Understanding this concept is crucial for predicting how a system at equilibrium will respond to changes, which is a key aspect of many chemical processes.