Real vs. ideal gases — A truly ideal gas cannot be liquefied. The primary reason is that, in the ideal-gas model,

Introduction:
Liquefaction of gases depends on intermolecular attractions strong enough to hold molecules together as a condensed phase. This question contrasts ideal-gas assumptions with the behavior required for liquefaction.

Given Data / Assumptions:

• Ideal gas model neglects molecular size and attractions (no a, b terms as in van der Waals).
• Liquefaction requires a finite critical temperature and pressure determined by real interactions.

Concept / Approach:
In the ideal-gas limit, attractive forces are zero. Without attractions, there is no driving force for condensation, so an ideal gas cannot exhibit a liquid phase or a finite critical point. Real gases deviate from ideality precisely because of these interactions, enabling liquefaction below the critical temperature.

Step-by-Step Solution:
Recall ideal-gas postulates: no intermolecular forces; point-like molecules.Relate liquefaction to attractions that create a potential well facilitating condensation.Conclude the key reason: absence of intermolecular attractions in the ideal model.

Verification / Alternative check:
The van der Waals equation incorporates parameters a (attractions) and b (co-volume). Setting a = 0 removes the possibility of a liquid phase.

Why Other Options Are Wrong:

• (a) Small size alone does not prevent liquefaction; noble gases are small yet liquefy.
• (b) Critical temperature varies by gas and is not linked to 0°C specifically.
• (d) Gases do not “skip” the liquid phase under normal conditions; sublimation pertains to solids.
• (e) Molecular rotation is not the governing factor for liquefaction.

Common Pitfalls:
Assuming liquefaction depends on absolute temperature markers like 0°C; conflating idealized models with real behavior.

Final Answer:
intermolecular forces are negligible (no attractions)