Question

The reaction \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)\) is exothermic as written. For the maximum production of ammonia, should this reaction be performed at a lower or a higher temperature? Explain.

Short answer

For the maximum production of ammonia in the exothermic reaction \( N_2(g) + 3 H_2(g) \rightleftharpoons 2 NH_3(g) \), it should be performed at a lower temperature. According to Le Chatelier's principle, lowering the temperature will cause the system to shift to the forward reaction, increasing the production of ammonia.

Step-by-step solution

Recognize the reaction as exothermic

The reaction is exothermic as written, which means that it releases heat as the reaction proceeds. The reverse reaction, which is the decomposition of ammonia into nitrogen and hydrogen gas, is endothermic (i.e., it absorbs heat).

Apply Le Chatelier's principle

According to Le Chatelier's principle, a system at equilibrium will adjust itself to counteract any change in conditions. In the case of an exothermic reaction like the formation of ammonia, increasing the temperature will cause the system to shift to the endothermic (reverse) reaction in order to absorb the excess heat and restore equilibrium. This means that increasing the temperature will actually decrease the production of ammonia. On the other hand, if the temperature is decreased, the system will try to counteract the change by shifting to the exothermic (forward) reaction, which releases heat. This will result in an increase in the production of ammonia.

Conclude the answer

Based on Le Chatelier's principle, for the maximum production of ammonia, this exothermic reaction should be performed at a lower temperature. Lowering the temperature will cause the system to shift to the forward reaction, which is the production of ammonia, therefore maximizing its production.

Key concepts

Exothermic Reactions

In chemistry, an exothermic reaction is one where energy, mostly in the form of heat, is released into the surrounding environment. This release of energy occurs because the total energy of the products is less than the total energy of the reactants. As a result, the difference in energy is expelled as heat. An everyday example of an exothermic reaction is the combustion of gasoline in a car engine, which produces heat and propels the vehicle forward.

Understanding exothermic reactions is crucial in industrial processes and laboratory settings. For instance, when developing a reaction that needs to maintain a consistent temperature, it's important to consider the heat released in exothermic reactions. Too much heat can lead to a runaway reaction, while not enough might mean the reaction does not proceed efficiently.

Chemical Equilibrium

Chemical equilibrium is a balanced state in which a chemical reaction and its reverse process occur at the same rate. At this point, the concentrations of reactants and products remain constant over time, even though both reactions are still occurring. This balance can be visualized as a pair of scales in perfect balance, with neither side outweighing the other.

Le Chatelier's principle is the cornerstone of understanding how a system at equilibrium responds to changes in concentration, temperature, and pressure. It states that if a dynamic balance is disturbed by changing the conditions, the position of equilibrium will shift to counteract the change. For students, visualizing this principle can be facilitated by imagining a crowded room where everyone is trying to move away from an uncomfortably warm spot. Just as people in the room redistribute to restore comfort, chemical systems redistribute to reach a new equilibrium.

Ammonia Production

Ammonia production is commonly executed through the Haber-Bosch process, which combines hydrogen and nitrogen gases under high pressure and temperature in the presence of a catalyst. The reaction is crucial for creating fertilizers, which are vital for food production around the world.

The production of ammonia is a delicate balance. The optimal conditions must be reached to maximize output without expending excessive energy. By applying Le Chatelier's principle to the exothermic reaction of ammonia synthesis, one can determine that lower temperatures should favor the formation of ammonia. However, in practice, a compromise temperature is often chosen to ensure a reasonable reaction rate and an economically viable industrial process.