Problem 28 - Chemistry and Chemical Reactivity

Question

The decomposition of \(\mathrm{NH}_{4} \mathrm{HS}\) $$ \mathrm{NH}_{4} \mathrm{HS}(\mathrm{s}) \rightleftarrows \mathrm{NH}_{3}(\mathrm{g})+\mathrm{H}_{2} \mathrm{S}(\mathrm{g}) $$ is an endothermic process. Using Le Chatelier's principle, explain how increasing the temperature would affect the equilibrium. If more \(\mathrm{NH}_{4} \mathrm{HS}\) is added to a flask in which this equilibrium exists, how is the equilibrium affected? What if some additional \(\mathrm{NH}_{3}\) is placed in the flask? What will happen to the pressure of \(\mathrm{NH}_{3}\) if some \(\mathrm{H}_{2} \mathrm{S}\) is removed from the flask?

Step-by-Step Solution

Answer

Increasing temperature shifts right. Adding \( \mathrm{NH}_{4} \mathrm{HS} \) does nothing. Adding \( \mathrm{NH}_{3} \) shifts left. Removing \( \mathrm{H}_{2} \mathrm{S} \) increases \( \mathrm{NH}_{3} \) pressure.

Step 1: Understanding the Reaction

We have the reaction \( \mathrm{NH}_{4} \mathrm{HS(s)} \rightleftharpoons \mathrm{NH}_{3(g)} + \mathrm{H}_{2} \mathrm{S(g)} \). This is an endothermic reaction, meaning heat is absorbed when the reaction proceeds in the forward direction.

Step 2: Effect of Temperature Increase

According to Le Chatelier's principle, raising the temperature of an endothermic reaction shifts equilibrium to the right toward the products to absorb more heat. Therefore, the concentration of \( \mathrm{NH}_{3} \) and \( \mathrm{H}_{2} \mathrm{S} \) will increase.

Step 3: Adding More Solid \( \mathrm{NH}_{4} \mathrm{HS} \)

Adding more \( \mathrm{NH}_{4} \mathrm{HS} \), the solid reactant, does not affect the equilibrium position. The equilibrium will not shift because the concentration of the solid does not appear in the equilibrium expression.

Step 4: Adding \( \mathrm{NH}_{3} \)

Adding \( \mathrm{NH}_{3} \) will shift the equilibrium to the left to consume some of this added \( \mathrm{NH}_{3} \). Consequently, less \( \mathrm{NH}_{3} \) and \( \mathrm{H}_{2} \mathrm{S} \) will be produced.

Step 5: Removing \( \mathrm{H}_{2} \mathrm{S} \)

Removing \( \mathrm{H}_{2} \mathrm{S} \) will shift the equilibrium to the right to replace the \( \mathrm{H}_{2} \mathrm{S} \) that was removed. This will result in an increase in \( \mathrm{NH}_{3} \) pressure.

Key Concepts

Equilibrium

In chemistry, equilibrium refers to a state where the rates of the forward and reverse reactions are equal. In the case of the decomposition of \( \mathrm{NH}_{4} \mathrm{HS} \), both the decomposition of the solid into \( \mathrm{NH}_{3} \) and \( \mathrm{H}_{2} \mathrm{S} \) gases and the recombination of these gases into the solid occur at the same rate. This balance means the concentrations of reactants and products remain constant over time.Equilibrium can be dynamic, meaning the reactions don't stop; they just occur at an equal rate, keeping the system balanced. If the system is disturbed, for example, by changes in concentration, pressure, or temperature, it tends to shift in a direction to counteract the disturbance and re-establish equilibrium, according to Le Chatelier's principle.When more \( \mathrm{NH}_{4} \mathrm{HS} \) is added, because it is a solid, it does not affect the equilibrium position. Solids and pure liquids do not appear in the equilibrium expression as their concentrations are constant. Thus, changes in the amount of solid do not influence the equilibrium directly. However, adding \( \mathrm{NH}_{3} \), a gaseous product, to the system will shift the equilibrium toward the left. This causes the system to attempt to consume some of the added \( \mathrm{NH}_{3} \), moving toward the reactant side.

Endothermic Reactions

Endothermic reactions are chemical processes that absorb heat from their surroundings. In simpler terms, energy in the form of heat is required for the reaction to proceed. The decomposition reaction \( \mathrm{NH}_{4} \mathrm{HS(s)} \rightarrow \mathrm{NH}_{3(g)} + \mathrm{H}_{2} \mathrm{S(g)} \) is such an endothermic process, meaning heat helps push it toward the formation of products.According to Le Chatelier's Principle, if the temperature of the reaction system is increased, the equilibrium will shift in the direction that absorbs the excess heat. For endothermic reactions, this means the equilibrium will shift towards the products. In our exercise, higher temperatures result in more \( \mathrm{NH}_{3} \) and \( \mathrm{H}_{2} \mathrm{S} \) gases being formed.This concept is crucial in industrial processes where controlling the direction of the reaction by temperature adjustments can optimize production of desired products.

Chemical Decomposition

Chemical decomposition is a type of reaction where one compound breaks down into two or more simpler substances. This process is essential in various natural and industrial processes. In the case of \( \mathrm{NH}_{4} \mathrm{HS} \), it decomposes into \( \mathrm{NH}_{3} \) and \( \mathrm{H}_{2} \mathrm{S} \), both gases, under the right conditions.This decomposition is driven by heat in an endothermic reaction, meaning sufficient heat must be provided for the transformation to take place efficiently. The importance of decomposition reactions is evident in the breaking down of more complex molecules into simpler forms, which can then be used in further chemical reactions. Decomposition reactions are often influenced by factors such as temperature and pressure. Removing products like \( \mathrm{H}_{2} \mathrm{S} \) encourages the forward reaction by shifting equilibrium to replace the removed gas, demonstrating how the principles of chemical decomposition and equilibrium are interrelated.