91Ó°ÊÓ

Le Chatelier's principle is a fundamental concept in chemistry, providing insight into how chemical equilibria respond to external changes. If a system in equilibrium experiences a change in concentration, pressure, or temperature, it will adjust itself to counteract the effect and re-establish equilibrium. For instance, if the pressure of a gaseous reaction is decreased, the system will shift the equilibrium position to increase the number of gas particles, effectively increasing pressure.
This principle is particularly useful in predicting the behavior of reactions under different conditions. It explains why, in our exercise, a reduction in the initial pressure of \(\mathrm{N}_{2} \mathrm{O}_{4}\), leads to a higher percent dissociation. At lower pressures, the shift of equilibrium towards the right side (formation of more gas particles \(\mathrm{NO}_{2}\)) is favored according to Le Chatelier's principle, aiming to restore pressure.

The equilibrium constant for a reaction involving gases, denoted as \(K_p\), is a measure of the ratio of the pressures of products to reactants at equilibrium, each raised to the power of their respective coefficients in the balanced chemical equation. This constant provides crucial information about the position of equilibrium; larger values of \(K_p\) indicate a predominance of products, while smaller values suggest a predominance of reactants at equilibrium.
In the given exercise, the equilibrium constant \(K_p\) for the reaction \[ \mathrm{N}_{2} \mathrm{O}_{4}(g) \rightleftharpoons 2\mathrm{NO}_{2}(g) \] was calculated using the pressures of \(\mathrm{NO}_{2}\) and \(\mathrm{N}_{2} \mathrm{O}_{4}\) at equilibrium. With the relationship \[ K_p = \frac{(P_{\mathrm{NO}_{2}})^2}{P_{\mathrm{N}_{2} \mathrm{O}_{4}}} = \frac{0.16}{0.80} = 0.20 \] it confirms the state of balance between products and reactants.

Percent dissociation is a measure of how much a specific compound breaks down into its constituents compared to its initial concentration or pressure. This is particularly relevant for equilibrium reactions, where not all reactants completely dissociate. The percent dissociation is calculated by dividing the amount dissociated by the initial amount, then multiplying by 100 to convert it to a percentage.
For the reaction \(\mathrm{N}_{2} \mathrm{O}_{4} \rightleftharpoons 2\mathrm{NO}_{2}\), in our exercise, the initial sample has a dissociation percentage that changes with pressure. Initially, at 1.00 atm pressure, 20% of \(\mathrm{N}_{2} \mathrm{O}_{4}\) dissociates at equilibrium. When the pressure is reduced to 0.10 atm, the percent dissociation increases to 45%, illustrating the impact of pressure on equilibrium. This elucidates how equilibrium constants help predict the degree of dissociation under varying conditions.