Reversible reactions and dynamic equilibrium：練習問題

Consider the following reversible reaction involving hydrated copper(II) sulfate:
\(blue \ hydrated \ copper(II) \ sulfate \rightleftharpoons white \ anhydrous \ copper(II) \ sulfate + water\)

If the forward reaction is endothermic, what will be observed when water is added to the white anhydrous copper(II) sulfate?

A student investigates the following equilibrium in a sealed container:
\(N_{2}O_{4}(g) \rightleftharpoons 2NO_{2}(g)\)

The forward reaction is endothermic. If the student decreases the temperature of the system, what will be observed according to Le Chatelier's Principle?

Consider the following reversible reaction at equilibrium in a closed system:
\(2SO_{2}(g) + O_{2}(g) \rightleftharpoons 2SO_{3}(g)\)
The forward reaction is exothermic. According to Le Chatelier’s Principle, what is the effect of increasing the temperature?

In a reversible chemical reaction, if the forward reaction is exothermic, what can be concluded about the reverse reaction according to the syllabus principles?

For the following gaseous reaction at equilibrium:
\(2NO(g) + O_{2}(g) \rightleftharpoons 2NO_{2}(g)\)

The forward reaction is exothermic. According to Le Chatelier’s Principle, what happens if the pressure of the system is increased?

(HT only) Consider the following reversible gaseous reaction at equilibrium:
\(2SO_{2}(g) + O_{2}(g) \rightleftharpoons 2SO_{3}(g)\)

The forward reaction is exothermic. Predict and explain the effect of decreasing the temperature on the position of equilibrium and the yield of \(SO_{3}\) according to Le Chatelier’s Principle.

(HT only) Consider the following reversible reaction in a closed system at equilibrium:
\(2SO_{2}(g) + O_{2}(g) \rightleftharpoons 2SO_{3}(g)\)

The forward reaction is exothermic.

a) Use Le Chatelier's Principle to predict and explain the effect of decreasing the temperature on the position of equilibrium and the yield of sulfur trioxide (\(SO_{3}\)).
b) Predict the effect of increasing the total pressure on the position of equilibrium. Justify your answer by referring to the number of molecules of gas on each side of the balanced equation.
c) In an industrial setting, a catalyst is used. Explain why the catalyst does not affect the final position of equilibrium or the percentage yield of \(SO_{3}\).
d) A student suggests that the concentration of \(SO_{3}\) could be increased by constantly removing it from the reaction vessel. Explain why this would prevent the system from reaching a state of dynamic equilibrium.

The production of ammonia in the Haber process involves the following reversible reaction:
\( \text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g}) \)
The forward reaction is exothermic.

a) Use Le Chatelier's Principle to explain the effect of increasing the pressure on the percentage yield of ammonia.
b) Explain why an industrial pressure of \( 200\text{ atmospheres} \) is used rather than a much higher pressure, considering safety and cost.
c) Explain why the reaction is carried out at a temperature of \( 450^\circ\text{C} \) even though a lower temperature would produce a higher yield of ammonia.