Consider the system $$\mathrm{SO}_{3}(g) \rightleftharpoons \mathrm{SO}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \quad \Delta H=98.9 \mathrm{~kJ}$$ (a) Predict whether the forward or reverse reaction will occur when the equilibrium is disturbed by (1) adding oxygen gas. (2) compressing the system at constant temperature. (3) adding argon gas. (4) removing \(\mathrm{SO}_{2}(g)\). (5) decreasing the temperature. (b) Which of the above factors will increase the value of \(K ?\) Which will decrease it?

Chemical equilibrium is a state in a reversible chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant over time, not because the reactions have stopped, but because they are occurring at equal rates.

Understanding chemical equilibrium is crucial, as it explains how different factors can influence the concentrations of substances involved in a reaction. It's like a scale balanced perfectly; any change to the system can tip the scale in favor of one side or the other. The Le Chatelier's Principle explains precisely how a reaction will respond to such changes to maintain this balance. For instance, adding a component like oxygen gas to the system will shift the balance, causing the reaction to adjust accordingly—in this case, by producing more sulfur trioxide (SO₃) to reduce the excess oxygen.

The reaction quotient, denoted as Q, is a measure that compares the concentrations of reactants and products in a reaction at any given point in time. It's calculated in the same way as the equilibrium constant (K), but it isn't limited to conditions at equilibrium.

By comparing the reaction quotient Q to the equilibrium constant K, we can predict the direction in which a reaction will proceed to reach equilibrium. If Q is less than K, the forward reaction is favored to produce more products. Conversely, if Q is greater than K, the reverse reaction is favored to produce more reactants. This is what happens when a system at equilibrium is disturbed -- the reaction quotient changes, and the system shifts to re-establish equilibrium.

The equilibrium constant, K, is a number that expresses the ratio of the concentrations of products to reactants at equilibrium for a reversible reaction at a constant temperature. Each substance's concentration is raised to the power of its coefficient in the balanced chemical equation.

K offers insight into the position of equilibrium; a large K indicates that the equilibrium favors products, while a small K shows a reactant-favored equilibrium. It's essential to note that K is only affected by changes in temperature; alterations in concentrations, volumes, or pressures (except for changes due to temperature) do not affect the value of K, although they may shift the position of equilibrium as dictated by Le Chatelier's Principle. Hence, understanding K helps students grasp why only changing the temperature alters this constant, as evidenced in our textbook problem.

Endothermic and exothermic reactions are two types of processes that either absorb or release energy, respectively. In an endothermic reaction, the system absorbs heat from its surroundings, resulting in a cooling effect. Conversely, in an exothermic reaction, heat is released into the surroundings, causing a warming effect.

The textbook exercise involves a reaction with a positive \( \Delta H \), indicating it's endothermic. Decreasing the temperature favors the exothermic process -- in this case, the reverse reaction. Recognizing whether a reaction is endothermic or exothermic is essential to predicting how changes in temperature will affect the equilibrium, as per Le Chatelier's Principle.