91Ó°ÊÓ

Understanding the reaction quotient, denoted as \( Q_c \), is crucial in predicting the direction in which a chemical reaction will proceed. The reaction quotient is similar to the equilibrium constant \( K_c \), but it uses the initial concentrations rather than the concentrations at equilibrium. For the given reaction:\[ \mathrm{I}_{2}(g) + \mathrm{Br}_{2}(g) \rightleftharpoons 2 \mathrm{IBr}(g) \] \( Q_c \) is expressed as:

• \( Q_c = \frac{[\text{IBr}]^2}{[\text{I}_2][\text{Br}_2]} \)

This formula comes from the law of mass action, which allows us to calculate \( Q_c \) using the initial amounts of each substance. By understanding \( Q_c \), students can determine how the initial system compares to the equilibrium state and thus predict the reaction's behavior. The calculated \( Q_c = 0.16 \) indicates how the initial products and reactants are distributed before equilibrium is achieved.

The equilibrium constant, denoted as \( K_c \), serves as a cornerstone in studying chemical equilibria. It reflects the ratio of product to reactant concentrations at static equilibrium. It is a fixed value at a particular temperature, depicting the reaction's innate capacity to reach a state where the rates of forward and reverse reactions are equal.For our specific reaction, the given \( K_c \) is \( 1.2 \times 10^{2} \). This large value suggests a position of equilibrium which strongly favors the formation of products, meaning at equilibrium, there are more products than reactants. Understanding \( K_c \) helps in:

• Predicting extent of a reaction: Large \( K_c \) implies product favorability.
• Understanding the system’s position at equilibrium.

Through \( K_c \), students see how the reaction naturally stabilizes in terms of the concentration of substances.

Determining the direction of a reaction to reach equilibrium involves comparing \( Q_c \) with \( K_c \). This comparison indicates whether the system will shift to favor products or reactants.- If \( Q_c < K_c \), the reaction moves in the forward direction to produce more products, achieving equilibrium.- If \( Q_c > K_c \), the reaction shifts in the reverse direction to produce more reactants.- If \( Q_c = K_c \), the system is already at equilibrium.In this exercise, with \( Q_c = 0.16 \) and \( K_c = 120 \), we find that \( Q_c < K_c \). Thus, the reaction will proceed in the forward direction, creating more \( \text{IBr} \) to reach equilibrium.By understanding these principles, students gain insight into reaction dynamics and equilibrium adjustments in response to initial conditions.