Question

Old fashioned "smelling salts" consist of ammonium carbonate, \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}\). The reaction for the decomposition of ammonium carbonate $$ \left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}(s) \rightleftharpoons 2 \mathrm{NH}_{3}(g)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) $$ is endothermic. What would be the effect on the position of this equilibrium if the reaction were performed at a higher temperature?

Short answer

Increasing the temperature of the endothermic decomposition of ammonium carbonate shifts the position of the equilibrium to the right, resulting in more products (ammonia, carbon dioxide, and water vapor) being formed.

Step-by-step solution

Identify the type of reaction

The given reaction is endothermic: \[ (\mathrm{NH}_{4})_{2}\mathrm{CO}_{3}(s) \rightleftharpoons 2\mathrm{NH}_{3}(g)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(g) \]

Apply Le Chatelier's principle

According to Le Chatelier's principle, an increase in temperature for an endothermic reaction will favor the forward reaction, causing the system to shift in the direction that absorbs the heat (to the right).

Determine the effect on the position of the equilibrium

Since the forward reaction is favored, the equilibrium will shift to the right. This causes more of the reactant, ammonium carbonate, to decompose and produces more of the products (ammonia, carbon dioxide, and water vapor) in the gaseous state. This results in the following change: \[ (\mathrm{NH}_{4})_{2}\mathrm{CO}_{3}(s) \xrightarrow[\text{Increased T}]{} 2\mathrm{NH}_{3}(g)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(g) \] In conclusion, increasing the temperature of the endothermic reaction, the decomposition of ammonium carbonate, shifts the position of the equilibrium to the right, resulting in more products (ammonia, carbon dioxide, and water vapor) being formed.

Key concepts

Le Chatelier's Principle

Le Chatelier's principle is a cornerstone concept in chemical equilibrium that describes how a system at equilibrium reacts to external changes. Under this principle, if a dynamic equilibrium is disturbed by changing the conditions, the position of the equilibrium will shift to counteract the change and re-establish equilibrium.

One way to visualize this principle is by thinking of a see-saw perfectly balanced in the middle. If you add more weight to one side, the see-saw will tilt in that direction; then to rebalance it, you'd either have to add the same weight to the other side or remove some weight from the heavier side. Similarly, in chemical reactions, changes in concentration, pressure, volume, or temperature can cause the equilibrium to shift in such a way as to minimize the effect of that change.

Understanding Le Chatelier's principle helps predict the direction a reaction shifts when exposed to various perturbations, which is invaluable for industries and laboratories where precise chemical conditions are required.

Endothermic Reactions

Endothermic reactions are chemical reactions that absorb heat from their surroundings during the process. The term 'endothermic' comes from the Greek words 'endo,' meaning 'within,' and 'thermos,' meaning 'heat.' These reactions often feel cold to the touch because they are taking in heat from the environment.

Analogous to melting ice cubes absorbing heat to turn into water, endothermic reactions require energy input to proceed. The additional energy is usually supplied in the form of heat, and as a result, these reactions are strongly influenced by temperature changes. In the classroom, endothermic reactions are often demonstrated with processes like the dissolution of ammonium nitrate in water or the reaction of barium hydroxide with ammonium chloride, both of which result in a significant temperature decrease of the solution.

One essential aspect of these reactions is that when the temperature is increased, the reaction rate typically speeds up, producing more products as the reaction seeks to absorb the added energy. This can be seen in the decomposition of ammonium carbonate, where the reaction absorbs heat and shifts the equilibrium toward the products.

Ammonium Carbonate Decomposition

The decomposition of ammonium carbonate (\((\text{NH}_4)_2\text{CO}_3\text(s)\)) into ammonia (\(\text{NH}_3\text(g)\)), carbon dioxide (\(\text{CO}_2\text(g)\)), and water vapor (\(\text{H}_2\text{O}\text(g)\)) is a classic example of an endothermic process. This means that the substance absorbs heat from its surroundings to break down into simpler molecules. Ammonium carbonate is often used in smelling salts because when it decomposes, it releases ammonia gas, which is pungent and can induce inhalation, potentially reviving someone who has fainted.

In terms of chemical equilibrium, this decomposition is reversible, and the balance between the solid and gaseous products can be altered by changing the temperature. According to Le Chatelier's principle, increasing the temperature will drive the equilibrium to the right, thus favoring the decomposition. This reaction is utilized in practical applications such as leavening agents in cooking and as a component of fire extinguishing compounds. Understanding the reaction's thermodynamic properties is crucial for effectively harnessing its benefits in various industrial and consumer processes.