91Ó°ÊÓ

The reaction quotient, often denoted as \( Q \), is a measure used to determine the direction in which a chemical reaction will proceed to reach equilibrium. You can calculate \( Q \) using the concentrations of the reactants and products at any point in time using the formula: \[ Q = \frac{[\text{products}]}{[\text{reactants}]} \] This is similar to the equilibrium constant \( K \), but \( Q \) can be used at any stage of the reaction, not just at equilibrium. In the given exercise, we calculated \( Q \) for the reaction \( A + B \rightleftharpoons C + D \) using the initial concentrations:

• \( Q = \frac{[C] \times [D]}{[A] \times [B]} = \frac{3 \times 4}{1 \times 2} = 6 \).

Comparing this value to the equilibrium constant helps in predicting the reaction's direction.

The equilibrium constant, denoted as \( K_c \), provides a snapshot of the ratio of product concentrations to reactant concentrations when a reaction is at equilibrium. Each chemical reaction has its unique \( K_c \) value, which remains constant at a given temperature. In the exercise, we know that \( K_c \) for the reaction \( A + B \rightleftharpoons C + D \) is 10 at 25°C. This signifies that, at equilibrium, the ratio \[ K_c = \frac{[C] \times [D]}{[A] \times [B]} = 10 \] must hold. The equilibrium constant tells us about the position of equilibrium and whether a reaction mixture will have more products or reactants at equilibrium. A large \( K_c \) suggests a product-favored reaction, while a small one suggests reactant-favored. Comparing this to the calculated \( Q \) allows us to predict the shift needed to reach equilibrium.

Determining the direction of a reaction is essential for understanding how a system reaches equilibrium. By comparing \( Q \) to \( K_c \), we can predict this direction:

• If \( Q < K_c \), the reaction will proceed from left to right (forming more products).
• If \( Q > K_c \), the reaction will proceed from right to left (forming more reactants).
• If \( Q = K_c \), the reaction is already at equilibrium.

In our specific problem, since \( Q = 6 < K_c = 10 \), the reaction will move from left to right, favoring the production of \( C \) and \( D \), hence moving towards equilibrium. Knowing the direction is crucial for controlling reactions in chemistry and industrial processes.