Introduction to Equilibrium

Understanding the Concept of Equilibrium

Equilibrium is a state in a chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction. This balance means the concentrations of the reactants and products remain constant over time.
Chemical equilibrium can only occur in a closed system, meaning no substances are added or removed during the reaction.
The point at which a reaction reaches equilibrium is dependent on the concentration of reactants, temperature, and pressure.

Equilibrium Expressions and the Equilibrium Constant, K

For a balanced chemical equation, the equilibrium constant expression (K) describes the relationship of the mathematical product of the concentrations of the products to the mathematical product of the concentrations of the reactants.
In a general reaction represented as aA + bB ⇌ cC + dD, the equilibrium constant expression (K) is given by K = ([C]^c [D]^d) / ([A]^a [B]^b) where [A], [B], [C], and [D] represent the molar concentrations of the species in the equilibrium reaction, and a, b, c, and d are respective stoichiometric coefficients. Each concentration is raised to the power of its stoichiometric coefficient.
The constant K is unitless, and its value is specific for a particular reaction at a particular temperature.

The Role of Le Chatelier’s Principle

Le Chatelier’s Principle states that if a system at equilibrium is subjected to a change, the system will shift its equilibrium position to counteract the effect of the change.
Changes which impact the position of equilibrium include changes in concentration, pressure, and temperature.
If additional reactant or product is added, the system will shift to use up the added substance. If reactant or product is removed, the system will shift to replace it.
If pressure is increased (for gaseous reactions), the system will shift in the direction that decreases the total number of gas molecules. If pressure is decreased, the system will shift in the direction that increases the total number of gas molecules.
If temperature is increased, the equilibrium shifts in the endothermic direction to absorb the added heat. If temperature is decreased, the equilibrium shifts in the exothermic direction to replace the lost heat.

Applications and Implications of Equilibrium

Equilibria are important in a number of chemical systems including those in industrial, environmental, and biological contexts.
Learning to predict the effects of varying conditions on a system at equilibrium allows us to control reactions to our advantage, for example, to maximize yields in chemical manufacturing processes.
Understanding the basic principles of equilibrium is crucial for more advanced studies in chemistry and related fields.