Question

For a reaction \[ A(g)+\frac{1}{2}B(g)\rightleftharpoons C(g)+\text{heat} \] favorable conditions for the reaction to occur in the forward direction are

• A. Low \(T\) and Low \(P\)
• B. Low \(T\) and High \(P\)
• C. High \(T\) and Low \(P\)
• D. High \(T\) and High \(P\)

Hint:

According to Le Chatelier’s principle: \[ \text{Exothermic reaction } \rightarrow \text{ favored by low temperature} \] and Equilibrium shifts toward fewer gaseous moles at high pressure.

Answer 1

The Correct Option is B

Step 1: Write and analyze the given reaction.
$A(g) + \frac{1}{2}B(g) \rightleftharpoons C(g) + $ heat. The reaction is exothermic (heat is a product).
Step 2: Apply Le Chatelier's principle for temperature.
Since the forward reaction releases heat, it is exothermic. By Le Chatelier's principle, decreasing temperature (low T) shifts equilibrium to the right (forward direction) to produce more heat. Therefore, LOW temperature favors the forward reaction.
Step 3: Count moles of gas on each side.
Reactant side: $1 + \frac{1}{2} = 1.5$ moles of gas. Product side: $1$ mole of gas. The forward reaction decreases the number of moles of gas.
Step 4: Apply Le Chatelier's principle for pressure.
Increasing pressure shifts equilibrium toward the side with fewer moles of gas. Since the product side has fewer moles (1 vs 1.5), HIGH pressure favors the forward reaction.
Step 5: Identify the favorable conditions.
Forward reaction is favored by: (i) Low temperature (exothermic reaction), and (ii) High pressure (fewer moles on product side).
Step 6: Match with options.
The condition that maximizes forward reaction yield is: Low Temperature + High Pressure.
\[ \boxed{\text{Low temperature and High pressure (option 2)}} \]