Problem 24 - 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 temperature (option a) is unfavorable.

Step 1: Understanding the Reaction

The given reaction is the Haber process: \( \text{N}_2(g) + 3\text{H}_2(g) \rightarrow 2\text{NH}_3(g) + 92.3 \, \text{kJ} \). It is an exothermic reaction, meaning it releases heat as it proceeds.

Step 2: Analyze Temperature Changes

As the Haber process is exothermic, increasing the temperature will shift the equilibrium to the left (towards the reactants) based on Le Chatelier's Principle. Conversely, reducing the temperature will favor the formation of ammonia by shifting the equilibrium to the right.

Step 3: Analyze Pressure Changes

The reaction involves a reduction in the number of moles of gas from 4 (1 \( \text{N}_2 \) and 3 \( \text{H}_2 \)) to 2 \( \text{NH}_3 \). Increasing the pressure will favor the reaction forming ammonia, as it shifts toward the side with fewer moles of gas.

Step 4: Consider Removal of Ammonia

Removing ammonia as it is formed shifts the equilibrium to the right, favoring the continuous production of more ammonia to maintain equilibrium.

Step 5: Determine the Unfavorable Condition

By evaluating each option: (a) increasing the temperature is unfavorable as it shifts equilibrium to the left. Options (b), (c), and (d) enhance ammonia production. Thus, option (a) is unfavorable.

Key Concepts

Exothermic Reaction

The Haber process, critical for producing ammonia, exemplifies an exothermic reaction. In such reactions, more energy is released than absorbed during product formation, resulting in heat being released into the surroundings.
This particular process involves combining nitrogen and hydrogen to produce ammonia along with a release of 92.3 kJ of energy.

Why is this important? Understanding that the Haber process is exothermic helps determine how temperature adjustments could affect the reaction.

• If the temperature increases, the reaction releases less heat, pushing the balance back toward reactants.
• Conversely, lower temperatures drive the reaction toward products, thereby favoring ammonia production.

The key takeaway is that exothermic reactions, by releasing heat when forming products, can be influenced significantly by temperature changes.

Le Chatelier's Principle

Le Chatelier's Principle offers insight into how equilibrium reactions respond to external changes, helping to predict outcomes of adjustments like temperature or pressure changes. It states that a system at equilibrium will shift to counteract any applied change disrupting its balance.

In the context of the Haber process, understanding this principle is essential. Because the reaction is exothermic:

• An increase in temperature makes the reaction shift toward the reactants since the system seeks to absorb the excess heat.
• Conversely, a decrease in temperature pushes the equilibrium toward more product formation (more ammonia), which aligns with our goals.

This principle allows us to predict and optimize conditions under which ammonia can be most efficiently produced.

Equilibrium Shift

At the heart of chemical reactions like the Haber process is the concept of equilibrium, where the forward and reverse reactions occur at the same rate.
Equilibrium can shift in response to conditions such as temperature and pressure changes or reactant/product concentration alterations.

The Haber process demonstrates how manipulating these conditions impacts reaction outcomes:

• Increasing temperature shifts equilibrium toward reactants, while decreasing temperature pushes it toward products according to **Le Chatelier's Principle**.
• Increasing pressure pushes the equilibrium toward the side with fewer gas moles, favoring ammonia production because fewer moles means less space occupation.

This understanding helps manage the reaction's direction and optimize conditions for maximum ammonia yield.

Gas Moles Reduction

A notable aspect of the Haber process is the reduction in moles of gas as reactants form products. Initially, 4 moles of gas (1 nitrogen and 3 hydrogen) shift to form just 2 moles of ammonia.
As per **Le Chatelier's Principle**, increasing the pressure favors the side of the reaction with fewer gas molecules, effectively driving the reaction toward more ammonia production.

Understanding this:

• Increasing the pressure during the reaction shifts equilibrium to the right (product side), capitalizing on the reduction in gas moles.
• This is desirable, as fewer moles equate to a smaller gaseous volume, thus reducing system pressure and encouraging more product.

This highlights the efficiency in focusing on mole reduction within gaseous reactions to enhance yield, a critical factor in the Haber process.