Question

Chemists have expended a great deal of effort to maximize the production of ammonia from the elements $$\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)+92 \mathrm{kJ}$$ Suggest three changes that could be made to this equilibrium system, which would tend to increase the yield of ammonia.

Short answer

To increase the yield of ammonia in this equilibrium system, we can: 1. Increase the pressure on the system, as this will shift the equilibrium position to the side with fewer moles of gas (the product side). 2. Decrease the temperature of the system, since the forward reaction is exothermic, and this will shift the equilibrium position to the right, favoring ammonia production. 3. Increase the concentration of one or both reactants (N₂ and H₂), which will shift the reaction towards the right, increasing ammonia production according to Le Chatelier's principle.

Step-by-step solution

Change 1: Increasing Pressure

According to Le Chatelier's principle, when pressure is increased in a system at equilibrium, the equilibrium position will shift to reduce the pressure by favoring the side with fewer moles of gas. In this case, we have 4 moles of reactants and 2 moles of products. So, increasing the pressure on this equilibrium system would shift the reaction towards the side that has fewer moles of gas, i.e., the product side, increasing the yield of ammonia.

Change 2: Decreasing Temperature

The given equation shows that the forward reaction is exothermic, i.e., it releases heat: \[ \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)+92 \mathrm{kJ} \] Thus, when the temperature is decreased, the equilibrium position is shifted toward the side that generates heat, according to Le Chatelier's principle. This means that decreasing the temperature would shift the equilibrium position to the right, favoring the production of ammonia and increasing its yield.

Change 3: Increasing the Reactant Concentrations

According to Le Chatelier's principle, if we increase the concentration of one or more reactants in the system, the equilibrium position will shift to counteract that change, favoring the side that consumes the increased reactant(s). In this case, increasing the concentrations of nitrogen gas and/or hydrogen gas would shift the reaction towards the right, increasing the production and yield of ammonia. In summary, to increase the yield of ammonia in this equilibrium system, we can: 1. Increase the pressure on the system 2. Decrease the temperature of the system 3. Increase the concentration of one or both reactants (i.e., N₂ and H₂)

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that describes how a dynamic equilibrium reacts to external changes. It asserts that if a stress is applied to a system at equilibrium, the system will adjust itself in a way that counteracts the stress and re-establishes equilibrium. This principle can be applied to changes in concentration, pressure, and temperature.

For instance, increasing the pressure of a gaseous equilibrium system favors the side with fewer gas moles, as seen in the production of ammonia. Adding more reactants shifts the equilibrium towards product formation to reduce the concentration of added reactants. Similarly, a change in temperature will shift the equilibrium to absorb the additional heat (in an exothermic reaction) or supply heat (in an endothermic reaction). Understanding this principle is crucial for controlling chemical reactions in industrial processes to maximize yield and efficiency of the desired products.

Chemical Equilibrium

Chemical Equilibrium refers to the state of a reversible chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of reactants and products over time. It is a dynamic balance, not a static one, meaning particles are constantly reacting in both directions, but because their rates are equal, the concentrations remain constant.

Such a state can be represented using an equilibrium constant, denoted as K, which is a numerical value derived from the ratio of product concentrations to reactant concentrations at equilibrium. Manipulating the conditions of a reaction, such as by changing concentration, pressure, or temperature, can shift the position of the equilibrium. This concept is essential in the industrial synthesis of chemicals, such as ammonia, where achieving optimal equilibrium conditions can significantly influence the production rates and economic viability of the process.

Exothermic Reactions

Exothermic Reactions are chemical reactions that release energy in the form of heat or light. The energy released is typically indicated in the chemical equation by a positive number, showing that it is a product of the reaction. A quintessential example is the synthesis of ammonia, where every two moles of ammonia produced release 92 kJ of energy.

In an industrial setting, managing the heat released in exothermic reactions is paramount. Decreasing the temperature of an exothermic reaction, according to Le Chatelier's Principle, will shift the equilibrium towards the products since the system will try to counteract the removal of heat by producing more heat. Thus, in the case of ammonia synthesis, maintaining a lower temperature can increase the yield of ammonia, although too low a temperature can also slow down the reaction rate. Therefore, a balance must be attained to optimize both reaction rate and yield.