Question

\(N_{2}+3H_{2}\rightleftharpoons2NH_{3}\) favours the following conditions

• A. Low Pressure. High Temperature
• B. High Pressure. Low Temperature
• C. Low Pressure. Low Temperature
• D. High Pressure. High Temperature

Hint:

Logic Tip: 4 gas molecules shrink into 2, so squeeze them with High Pressure! The reaction generates heat, so cool it down with Low Temperature to keep the reaction moving forward!

Answer 1

The Correct Option is B

Concept:
Le Chatelier's Principle states that if a change in conditions (pressure, temperature) is imposed on a system at equilibrium, the equilibrium position will shift to counteract that change. This principle dictates how to maximize product yield in reversible reactions.

Step 1:
Look at the molar ratio of gases: 1 mole of $N_2$ + 3 moles of $H_2$ react to form 2 moles of $NH_3$. This is 4 moles of reactant gas turning into 2 moles of product gas.

Step 2:
An increase in pressure shifts the equilibrium towards the side with fewer moles of gas to reduce the pressure. Since the product side has fewer moles (2 vs 4), High Pressure favors the forward reaction.

Step 3:
The synthesis of ammonia (Haber process) is an exothermic reaction, meaning it releases heat ($\Delta H$ is negative).

Step 4:
For an exothermic reaction, lowering the temperature causes the system to try and produce more heat by shifting equilibrium to the right (the exothermic direction). Thus, Low Temperature favors the forward reaction.

Step 5:
To maximize the production of $NH_3$, the system theoretically requires High Pressure (to squeeze 4 molecules into 2) and Low Temperature (to draw out the released heat).