Question

Solid \(\mathrm{NH}_{4} \mathrm{HS}\) decomposes by the following endothermic process: $$\mathrm{NH}_{4} \mathrm{HS}(s) \leftrightharpoons \mathrm{NH}_{3}(g)+\mathrm{H}_{2} \mathrm{S}(g)$$ a. What effect will adding more \(\mathrm{NH}_{3}(g)\) have on the equilibrium? b. What effect will adding more \(\mathrm{NH}_{4} \mathrm{HS}(s)\) have on the equilibrium? c. What effect will increasing the volume of the container have on the equilibrium? d. What effect will decreasing the temperature have on the equilibrium?

Short answer

a) Adding more NH3(g) will shift the equilibrium towards the reactants (left), favoring the formation of NH4HS(s). b) Adding more NH4HS(s) will have no effect on the equilibrium position. c) Increasing the volume of the container will shift the equilibrium to the right, favoring the formation of NH3(g) and H2S(g). d) Decreasing the temperature will shift the equilibrium to the right, favoring the formation of NH3(g) and H2S(g).

Step-by-step solution

Understand Le Châtelier's Principle

Le Châtelier's principle states that when a system at equilibrium is subjected to a change in concentration, temperature, or volume, the equilibrium will shift to counteract the change and re-establish equilibrium. Based on this principle, we can predict how the given changes will affect the equilibrium.

Adding more NH3 (g)

According to Le Châtelier's principle, adding more NH3 will cause the equilibrium to shift in the direction that decreases its concentration. In this case, it will shift to the left, favoring the formation of NH4HS(s) and consuming the excess NH3. So, adding more ammonia will shift the equilibrium towards the reactants.

Adding more NH4HS (s)

Since NH4HS is a solid, its concentration does not affect the equilibrium. Therefore, adding more NH4HS will have no effect on the equilibrium position.

Increasing the volume of the container

Increasing the volume of the container will cause the pressure to decrease. According to Le Châtelier's principle, the equilibrium will shift to the direction with more moles of gas to counteract the pressure change. In this reaction, there are more moles of gas on the right side, so the equilibrium will shift to the right. This favors the formation of NH3(g) and H2S(g).

Decreasing the Temperature

The reaction is endothermic, which means it absorbs heat as it proceeds in the forward direction. Decreasing the temperature means removing heat from the system. In this case, the equilibrium will shift in the direction that supplies the heat, which is the forward direction. Therefore, the equilibrium will shift to the right, favoring the formation of NH3(g) and H2S(g).

Key concepts

Chemical Equilibrium

Chemical equilibrium occurs when a reversible reaction's forward and backward rates are equal. This means the concentrations of reactants and products remain stable over time. It's important to note that equilibrium does not mean the reaction stops; the processes continue but at the same rate, so there's no net change.

In our original exercise, the equilibrium is represented by:
\[\mathrm{NH}_{4} \mathrm{HS}(s) \leftrightharpoons \mathrm{NH}_{3}(g)+\mathrm{H}_{2} \mathrm{S}(g)\]
Here, the solid ammonium hydrogen sulfide decomposes into ammonia and hydrogen sulfide gases. At equilibrium, the balanced state is maintained unless changes in concentration, pressure, or temperature occur. Understanding this concept equips us to predict how a system will respond to various changes, consistent with Le Châtelier's principle.

Effect of Concentration Changes

When the concentration of a component in a chemical equilibrium is changed, the equilibrium shifts to counteract this change and establish a new balance.

Let's illustrate this with the provided reaction. When more \(\mathrm{NH}_{3}(g)\) is added to the system, Le Châtelier's principle suggests the equilibrium will shift to reduce this concentration increase. It achieves this by shifting to the left, causing the formation of more \(\mathrm{NH}_{4} \mathrm{HS}(s)\), a solid that doesn't affect equilibrium through concentration changes.

However, adding more \(\mathrm{NH}_{4} \mathrm{HS}(s)\) itself shows no noticeable shift since solids do not alter the equilibrium in this manner. These scenarios emphasize the role of concentration changes in driving equilibrium shifts.

Reaction Volume Changes

Changing the volume of the reaction container influences the pressure, particularly when gases are involved. According to Le Châtelier's principle, increasing volume decreases the system pressure.

In the given chemical reaction, there are no mole gases on the left, but two mole gases on the right (\(\mathrm{NH}_{3}(g)\) and \(\mathrm{H}_{2} \mathrm{S}(g)\)). A volume increase will shift the equilibrium towards the side with more moles of gas to counterbalance the pressure change: in this case, to the right.

• This shift results in increased production of \(\mathrm{NH}_{3}(g)\) and \(\mathrm{H}_{2} \mathrm{S}(g)\).

Modifying the reaction volume exemplifies how equilibrium responds to maintain balance, aligning itself with updated conditions.

Temperature Effect on Equilibrium

Temperature changes impact equilibrium differently, depending on whether a reaction is endothermic or exothermic. For endothermic reactions like the decomposition of \(\mathrm{NH}_{4} \mathrm{HS}\), heat is absorbed as the reaction proceeds in the forward direction.

By reducing temperature, heat is essentially removed from the system, prompting the equilibrium to shift in a heat-producing direction—the forward reaction in this scenario. This adjustment increases concentrations of \(\mathrm{NH}_{3}(g)\) and \(\mathrm{H}_{2} \mathrm{S}(g)\).

Le Châtelier's principle helps visualize how equilibrium strives to accommodate heat changes, ultimately dictating reaction direction under varied thermal conditions. Such insights are pivotal in predicting chemical behavior in fluctuating environments.