MCQ Detailed View

The concept of an ideal gas is based on two key assumptions: first, that gas molecules occupy negligible volume compared to the container, and second, that there are no intermolecular attractions or repulsions between the gas molecules. In reality, all gases deviate from this behavior to some extent, but some come closer than others under ordinary conditions.

At 0 °C, the degree to which a gas behaves ideally depends mainly on molecular size and intermolecular forces. Smaller, lighter, and non-polar molecules generally exhibit behavior closest to the ideal gas law.

Among the given gases, helium (He) shows the most ideal behavior. Helium is a noble gas with a monoatomic structure. Its atoms are extremely small, non-polar, and interact only through very weak London dispersion forces. These minimal interactions make helium nearly an ideal gas, even at relatively low temperatures such as 0 °C.

Hydrogen (H₂), while light and small, still has slightly stronger intermolecular interactions than helium due to being a diatomic molecule. Methane (CH₄) is larger in size and has more electrons, which increases dispersion forces, making it deviate further from ideal gas behavior. Ammonia (NH₃) deviates the most because it is a polar molecule capable of strong hydrogen bonding, which significantly increases intermolecular attraction.

Therefore, at 0 °C, helium is the gas that shows the most ideal behavior. This is why helium and other noble gases are often used as examples in ideal gas law problems. Their weak intermolecular forces and small atomic size allow them to follow gas laws such as Boyle’s law, Charles’ law, and the ideal gas equation with minimal deviation.