States of Matter: Question 1

Worked solution

Step 1: Recall the assumptions of the ideal gas model

The ideal gas model assumes that (i) the particles themselves have negligible volume compared with the volume of the container, and (ii) there are no intermolecular forces of attraction or repulsion between particles, so all collisions are perfectly elastic. Real gases only approximate this model closely when their molecules are small and their intermolecular forces are very weak.

Step 2: Compare the intermolecular forces in each gas

• Helium (A): a monatomic noble gas with only extremely weak, instantaneous dipole–induced dipole (London/van der Waals) forces. This makes it one of the closest real gases to ideal behaviour.
• Nitrogen (B): a small, non-polar diatomic molecule with weak van der Waals forces, slightly stronger than helium’s because $\text{N}_2$ has more electrons, but still weak overall.
• Methane (C): a slightly larger non-polar molecule, so its van der Waals forces are a little stronger than nitrogen’s, but methane is still non-polar with no permanent dipole.
• Ammonia (D): a polar molecule that forms hydrogen bonds between molecules (N–H···N), which are far stronger than the van der Waals forces present in A, B and C.

Step 3: Link intermolecular force strength to deviation from ideality

The ideal gas equation ignores intermolecular forces entirely. The stronger the real intermolecular forces in a gas, the more the actual pressure and volume differ from the values predicted by $pV = nRT$, because attractive forces pull molecules together (reducing the pressure the gas actually exerts) and take up a proportionally larger share of the total behaviour at a molecular level. Since ammonia has by far the strongest intermolecular forces (hydrogen bonding) of the four gases listed, it deviates the most from ideal gas behaviour.

Why the other options are wrong

• A (helium): helium is the gas that behaves most ideally, not least, because its intermolecular forces are the weakest of all four.
• B (nitrogen): only very weak van der Waals forces are present; nitrogen is commonly used as a benchmark “nearly ideal” gas in calculations.
• C (methane): methane’s van der Waals forces are stronger than nitrogen’s but still far weaker than the hydrogen bonding present in ammonia.

Final answer

• The gas expected to deviate most from ideal behaviour is $\boxed{\text{ammonia, } \text{NH}_3}$, option D.