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Chemical equilibrium is a state of balance in a reversible chemical reaction. When a reaction reaches this state, the rate of the forward reaction equals the rate of the backward reaction. This does not necessarily mean that the concentrations of reactants and products are equal, but rather that they remain constant over time.
The equilibrium constant, denoted as \( K_c \), is a numerical value that characterizes the ratio of concentrations of products to reactants for a reaction at equilibrium. Each concentration is raised to the power of its coefficient in the balanced equation. It is important to note that this constant changes with temperature and pressure.
In our example reaction, the equilibrium constant expression is derived from the balanced chemical equation: \( 3 \mathrm{~A}(g)+\mathrm{B}(s) \rightleftharpoons 2 \mathrm{C}(aq)+\mathrm{D}(aq) \). Since \( \mathrm{B} \) is a solid, it is omitted from the expression for \( K_c \). Thus, \( K_c = \frac{[\mathrm{C}]^2[\mathrm{D}]}{[\mathrm{A}]^3} \), where the square brackets denote molarity or concentration in moles per liter.

Le Chatelier's Principle provides insight into how a system at equilibrium responds to external changes. It states that if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system will adjust itself to partially counteract the effect of the change.
For instance, if the concentration of reactants in a reaction is increased, the equilibrium will shift towards the products to reduce the impact of this change. Conversely, increasing the concentration of products will push the equilibrium towards the reactants.
In temperature changes, increasing the temperature of an exothermic reaction results in the equilibrium shifting towards the reactants as the system tries to absorb the extra heat. This principle not only helps in understanding the nature of equilibriums but also in predicting how equilibriums can be manipulated effectively.

Reversible reactions are chemical reactions where the conversion of reactants to products and the conversion of products back to reactants occur simultaneously. These reactions are represented by the equilibrium symbol \( \rightleftharpoons \), indicating that reactions can proceed in both directions.
Reversibility is a common characteristic in many natural and industrial chemical processes. In reversible reactions, both reactants and products are present in the system, and the reaction reaches a state of equilibrium over time.
An example of a reversible reaction is the one mentioned in the exercise, \( 3 \mathrm{~A}(g)+\mathrm{B}(s) \rightleftharpoons 2 \mathrm{C}(aq)+\mathrm{D}(aq) \). Such reactions highlight the dynamic nature of equilibrium, where molecules change continuously, yet the overall concentrations remain constant, illustrating the balance that defines chemical equilibrium.