ANS:D - None of the above.

The application of vacuum will help in a reaction where the number of gas molecules in the product side is more than that on the reactant side. Here's why:

1. Le Chatelier's Principle: According to Le Chatelier's Principle, when a system at equilibrium is disturbed by changes in temperature, pressure, or concentration, the system will adjust to counteract the change and restore equilibrium. In a chemical reaction involving gases, changing the pressure by applying vacuum (lowering the pressure) can cause the equilibrium to shift in the direction that reduces the total number of gas molecules.
2. Increase in Volume: When vacuum is applied, the pressure is decreased, which effectively increases the volume of the reaction vessel. In a reaction where the number of gas molecules in the product side is more than that on the reactant side, shifting the equilibrium towards the side with fewer gas molecules will occupy the larger volume, thereby reducing the pressure and counteracting the vacuum.
3. Example: For example, consider the reaction: N2(g) + 3H2(g) ⇌ 2NH3(g) In this reaction, the number of gas molecules on the product side (2 molecules of NH3) is less than the total number of gas molecules on the reactant side (1 molecule of N2 + 3 molecules of H2 = 4 molecules). Applying vacuum would help shift the equilibrium towards the product side, favoring the formation of NH3, which reduces the total number of gas molecules and occupies the larger volume created by the vacuum.