91Ó°ÊÓ

Chemical equilibrium is a state in a reversible chemical reaction where the rate of the forward reaction equals the rate of the backward reaction. At this point, the concentrations of reactants and products remain constant over time, even though both reactions are still occurring. It's important to note that equilibrium does not mean the reactants and products are equal in concentration, but rather stable in their proportions.

In the context of the given exercise, the chemical equilibrium is represented by the reaction between hydrogen sulfide (\(2 \mathrm{H}_2 \mathrm{S}(g) \rightleftharpoons 2 \mathrm{H}_2(g) + \mathrm{S}_2(g)\)). At equilibrium, the concentrations of hydrogen sulfide, hydrogen, and sulfur are stable when the system is maintained at a constant temperature of 1130°C.

Le Chatelier's Principle is a helpful guideline for predicting how a chemical system at equilibrium responds to disturbances, such as changes in concentration, temperature, or pressure. According to this principle, if any of these conditions change, the equilibrium will shift in the direction that tends to reduce the impact of the change.

In the provided reaction, suppose we increased the concentration of hydrogen sulfide, \(\mathrm{H}_2 \mathrm{S}\). According to Le Chatelier's Principle, the system would adjust to counteract this change by favoring the forward reaction to produce more hydrogen and sulfur. Conversely, if hydrogen or sulfur were increased, the equilibrium would shift toward the formation of more \(\mathrm{H}_2 \mathrm{S}\), partially alleviating the increase in \(\mathrm{H}_2\) or \(\mathrm{S}_2\) concentrations.

Through recognizing these shifts, chemists can manipulate conditions to increase the yield of desired products.

The reaction quotient, sometimes denoted as \(Q_c\), serves as a valuable indicator in determining the current state of a reaction system compared to its equilibrium state. It is calculated using the same expression as the equilibrium constant \(K_c\) but with current concentrations instead of equilibrium concentrations.

For the reaction specified in the exercise, the expression for the reaction quotient \(Q_c\) is:

• \(Q_c = \frac{[\mathrm{H}_2]^2[\mathrm{S}_2]}{[\mathrm{H}_2 \mathrm{S}]^2}\)

Comparing \(Q_c\) with \(K_c\) can tell us whether the system is at equilibrium:

• If \(Q_c = K_c\), the system is at equilibrium.
• If \(Q_c < K_c\), the forward reaction will be favored to reach equilibrium.
• If \(Q_c > K_c\), the backward reaction will be favored.

In this exercise, although the actual concentrations at a given moment were not used to calculate \(Q_c\), understanding this relationship is crucial for predicting how changes affect the position of equilibrium.