91Ó°ÊÓ

Chemical equilibrium occurs in a reversible chemical reaction when the rate of the forward reaction equals the rate of the reverse reaction. This results in the concentrations of the reactants and products remaining constant over time. When a system reaches this state, it does not mean that reactions have stopped; rather, they are occurring at equal and opposite rates.
Understanding this concept is crucial in predicting the behavior of a chemical reaction. The position of equilibrium gives us insight into the concentrations of reactants and products when equilibrium is achieved. The equilibrium constant, denoted as \( K_c \), is vital in quantifying this state. It helps us determine the ratio of product to reactant concentrations at equilibrium.
The magnitude of the equilibrium constant can tell us a lot about the reaction's characteristics at equilibrium. If \( K_c \) is large, it suggests that the products are favored, and the reaction proceeds significantly to the right. Conversely, a small \( K_c \) indicates that reactants are favored and little reaction occurs.

The extent of a reaction refers to how far a reaction progresses toward forming products before reaching equilibrium. It is closely related to the concept of the equilibrium constant \( K_c \).
In reactions with a large \( K_c \), like the reaction of sulfur dioxide (\( \ ext{SO}_2 \)) with oxygen (\( \ ext{O}_2 \)) to form sulfur trioxide (\( \ ext{SO}_3 \)), the extent of the reaction is significant. This occurs because a large \( K_c \) indicates that the equilibrium position is far to the right, favoring a nearly complete conversion of reactants to products.
This information is crucial when predicting the concentrations of substances in a system at equilibrium. In practical scenarios, knowing the extent of a reaction helps chemists optimize conditions to maximize the desired products.

Sulfur dioxide (\( \ ext{SO}_2 \)) is a colorless gas known for its sharp smell. It is a significant air pollutant and is produced by volcanic eruptions and industrial processes, such as the burning of fossil fuels.In the reaction \( 2\ \ ext{SO}_2(g) + \ ext{O}_2(g) \rightleftharpoons 2\ \ ext{SO}_3(g) \), sulfur dioxide reacts with oxygen to form sulfur trioxide. This reaction is notable for its large equilibrium constant at room temperature, indicating that when equilibrium is reached, the concentration of \( \ ext{SO}_2 \) is extremely low.Understanding the behavior of \( \ ext{SO}_2 \) in chemical reactions allows us to better comprehend its impact on the environment and its role in industrial chemical processes. This knowledge is helpful in controlling emissions and minimizing its impact on air quality.

The reaction quotient, denoted as \( Q_c \), is similar to the equilibrium constant but applies to non-equilibrium conditions. It is calculated using the concentrations of reactants and products at any stage prior to reaching equilibrium.
For a reaction \( a\ A + b\ B \rightleftharpoons c\ C + d\ D \), the reaction quotient \( Q_c \) is represented as: \[Q_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}\]
Comparing \( Q_c \) to \( K_c \) helps us predict the direction in which a reaction must proceed to reach equilibrium:

• If \( Q_c < K_c \), the reaction will proceed to the right, favoring products.
• If \( Q_c > K_c \), the reaction will shift to the left, favoring reactants.
• If \( Q_c = K_c \), the system is at equilibrium.

In our exercise, we find that almost all sulfur dioxide and oxygen are converted to sulfur trioxide, consistent with the very large \( K_c \). This indicates that the initial \( Q_c \) was much lower than \( K_c \), driving the reaction strongly toward product formation.