Problem 25 - Chemistry in 30 days

Question

In the manufacture of ammonia by haber's process $$ \mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{3}(\mathrm{~g})=2 \mathrm{NH}_{3}(\mathrm{~g})+92.3 \mathrm{~kJ} $$ Which of the following conditions is unfavourable? (a) increasing the temperature (b) increasing the pressure (c) reducing the temperature (d) removing ammonia as it is formed

Step-by-Step Solution

Answer

Increasing the temperature is unfavorable.

Step 1: Understand the Reaction

The Haber process for producing ammonia is represented by the equation below:\[\mathrm{N}_2(g) + 3\mathrm{H}_2(g) \rightarrow 2\mathrm{NH}_3(g) + 92.3 \text{ kJ}\] This equation shows that the reaction is exothermic, releasing 92.3 kJ of energy. Additionally, one molecule of nitrogen and three molecules of hydrogen form two molecules of ammonia.

Step 2: Analyze Temperature Changes

Since the reaction is exothermic, increasing the temperature will shift the equilibrium position in the opposite direction, favoring the reactants, according to Le Chatelier's principle. Therefore, higher temperatures would decrease the yield of ammonia, making high temperature undesirable for the synthesis.

Step 3: Analyze Pressure Changes

In the reaction \( \mathrm{N}_2(g) + 3\mathrm{H}_2(g) \rightarrow 2\mathrm{NH}_3(g) \), there are 4 moles of gas on the reactant side and 2 moles of gas on the product side. Increasing the pressure will shift the reaction towards the side with fewer moles of gas, favoring the production of ammonia, which is beneficial.

Step 4: Analyze Reducing Temperature

Reducing the temperature would favor the exothermic reaction, resulting in a higher yield of ammonia. This is usually a favorable condition for maximizing ammonia production in the Haber process.

Step 5: Analyze Removing Ammonia

Removing ammonia as it is formed shifts the equilibrium towards more product formation according to Le Chatelier's principle. This increases the yield of ammonia, making it a favorable practice in the Haber process.

Step 6: Conclusion on Unfavorable Condition

Based on the above analysis, increasing the temperature is the unfavorable condition because it shifts the equilibrium towards the reactants, reducing the ammonia yield.

Key Concepts

Le Chatelier's principle

Le Chatelier's principle is a fundamental concept in chemical equilibrium. It helps predict how a system at equilibrium will respond to changes in concentration, temperature, or pressure. When a change is applied to a system in equilibrium, the system will shift in the direction that counteracts the change to re-establish equilibrium.
For example, if you increase the concentration of reactants in the Haber process, the system will shift to produce more products to reduce the disturbance. Similarly, removing ammonia, which is a product, will shift the equilibrium to produce more ammonia. This behavior of shifting the equilibrium upon changes simply explains how systems prefer to balance themselves out, making Le Chatelier's principle an essential rule for understanding chemical reactions.

Exothermic reaction

An exothermic reaction is a type of chemical reaction that releases energy, usually in the form of heat, into its surroundings.
In the Haber process, the formation of ammonia from nitrogen and hydrogen is exothermic. It releases 92.3 kJ of energy. This tells us that the process of making ammonia gives off heat, contributing to the reaction's overall energy balance.
Since exothermic reactions release heat, increasing the temperature can have unintended effects. According to Le Chatelier's Principle, raising the temperature will push the equilibrium towards the reactants. This is because the system works to absorb the extra heat, resulting in a lower yield of ammonia. Thus, maintaining lower temperatures can often be more advantageous for exothermic reactions like those in the Haber process.

Chemical equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions in a chemical system are equal.
This means that the concentrations of reactants and products remain constant over time. In the Haber process, chemical equilibrium is represented by the balance between the formation of ammonia from nitrogen and hydrogen, and the decomposition of ammonia back into these elements.

• At equilibrium, no net change is observed in the concentration of reactants and products.
• Equilibrium is dynamic, meaning the reactions are ongoing even though macroscopic changes are not visible.

Understanding chemical equilibrium is crucial as it informs how changes in external conditions, like pressure and temperature, will affect the production yields in industrial processes.

Pressure effects

Pressure has significant effects on chemical reactions involving gases. In the Haber process, pressure is a critical factor owing to the involvement of gaseous reactants and products.
Let's look at the reaction: \( \mathrm{N}_2(g) + 3\mathrm{H}_2(g) \rightarrow 2\mathrm{NH}_3(g) \). On the left, there are 4 moles of gas, and on the right, there are 2 moles of ammonia gas. Increasing the pressure causes the system to shift toward the side with fewer moles of gas.
This shift favors the production of ammonia:

• High pressure decreases the volume of the system, driving the equilibrium towards the formation of less gaseous molecules, thus increasing ammonia yield.
• Industrial processes like the Haber process often operate at high pressures to maximize production efficiency.

Therefore, manipulating pressure smartly underlines the industrial viability of the Haber process.