Why does equilibrium not imply equal concentrations of reactants and products?

Equilibrium does not imply equal concentrations of reactants and products because it refers to the rate of reaction, not the concentrations.

In a chemical reaction, equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction. This does not mean that the concentrations of the reactants and products are equal. Instead, it means that the concentrations of the reactants and products remain constant over time because the rate at which they are being produced and consumed is the same.

The concept of equilibrium is based on the principle of dynamic equilibrium. This principle states that in a closed system, a state of equilibrium is reached when the rates of the forward and reverse reactions are equal. At this point, the concentrations of the reactants and products do not change, even though the reactions are still occurring. This is because the amount of reactants being converted into products is exactly balanced by the amount of products being converted back into reactants.

The equilibrium constant, denoted as Kc, is used to express the relationship between the concentrations of the reactants and products at equilibrium. It is calculated by dividing the product of the concentrations of the products by the product of the concentrations of the reactants, each raised to the power of their stoichiometric coefficients. The value of Kc indicates the extent of a reaction at equilibrium. If Kc is greater than 1, the reaction favours the products, and if Kc is less than 1, the reaction favours the reactants. However, regardless of the value of Kc, the concentrations of the reactants and products at equilibrium do not have to be equal.

In conclusion, equilibrium in a chemical reaction does not imply equal concentrations of reactants and products. It simply means that the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products over time.