Problem 156 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following will favour the formation of \(\mathrm{NH}_{3}\) by Haber's Process? (a) Increase of temperature (b) Increase of pressure (c) Addition of catalyst (d) Addition of promoter

Step-by-Step Solution

Answer

Increase of pressure favours ammonia formation in Haber’s Process.

Step 1: Understand the Reaction

Haber's Process is the synthesis of ammonia (\(\text{NH}_3\)) from nitrogen (\(\text{N}_2\)) and hydrogen (\(\text{H}_2\)), represented by the equation \(\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3\). This reaction is exothermic.

Step 2: Analyze the Effect of Temperature

Since the reaction is exothermic, according to Le Chatelier's principle, an increase in temperature would shift the equilibrium to the left, reducing the formation of ammonia. Therefore, a lower temperature favours formation, though too low a temperature can reduce reaction rate.

Step 3: Analyze the Effect of Pressure

The equation shows a decrease in moles of gas: from 4 moles (\(1 \text{N}_2 + 3 \text{H}_2\)) to 2 moles of \(\text{NH}_3\). Increasing pressure shifts the equilibrium towards the side with fewer moles, thus favouring the formation of ammonia.

Step 4: Consider the Role of Catalyst

A catalyst speeds up the rate at which equilibrium is reached but does not shift the position of the equilibrium. Therefore, while a catalyst makes the reaction occur faster, it doesn't necessarily favour ammonia formation in terms of equilibrium.

Step 5: Consider the Role of Promoter

A promoter enhances the effectiveness of a catalyst. While it aids the catalyst in speeding up the reaction, it also does not shift equilibrium towards the formation of ammonia; it merely improves the catalyst's efficiency.

Key Concepts

Effect of Temperature on Equilibrium

In chemical reactions, temperature can significantly affect the position of equilibrium. For the Haber Process, which synthesizes ammonia, the reaction is exothermic. This means it releases heat as it proceeds. According to Le Chatelier's principle, if you increase the temperature, the system will adjust to counteract that change by moving in the direction that absorbs heat. In this case, increasing temperature shifts the equilibrium to the left, away from ammonia formation, reducing the yield of \(\text{NH}_3\). Thus, lower temperatures favor ammonia formation. However, there's a catch. At too low temperatures, the reaction rate also slows, making the process inefficient. Therefore, in industry, a compromise is struck at moderately low temperatures to balance between a good yield and a reasonable reaction speed.

Pressure and Chemical Equilibrium

Pressure is a vital factor in many chemical equilibria involving gases. In the Haber Process, the equation \(\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3\) depicts a shift from 4 moles of reactant gases to 2 moles of product gas. According to Le Chatelier's principle, increasing the pressure of the system will favor the side with fewer gas moles, in this case, the formation of ammonia \(\text{NH}_3\). Therefore, high pressure shifts the equilibrium towards the production of \(\text{NH}_3\), maximizing yields. Despite this advantage, extremely high pressures require robust equipment and can be costly, so it's crucial to find a practical balance between pressure levels and economic feasibility.

Role of Catalysts in Chemical Reactions

Catalysts play a critical role in the Haber Process by speeding up the reaction without altering the equilibrium position. They work by providing an alternative pathway with a lower activation energy for the reaction, which means less energy is needed for the reaction to proceed. While catalysts do not favor the formation of more ammonia in terms of equilibrium, they allow the process to reach equilibrium faster. In the industrial Haber Process, iron-based catalysts are typically used to improve reaction rates at reasonable temperatures and pressures. By enhancing the reaction speed, catalysts make the process more efficient and economically viable, even though they do not change the final concentrations of reactants and products at equilibrium.