Physics · Kinetic Theory · NEET
| Obeys PV = nRT | Exactly, at all P and T | Only approximately; deviates at high P, low T |
| Molecular volume | Zero (point-size molecules) | Finite; molecules take up real space |
| Intermolecular force | None (no attraction or repulsion) | Present (molecules attract each other) |
| Exists in reality | No; it is only a model | Yes; all actual gases are real |
| Best behaviour | Always ideal by definition | Nearly ideal at low pressure and high temperature |
| Can it liquefy | Never (no attraction to pull molecules together) | Yes, at low temperature and high pressure |
No. NCERT states directly: 'An ideal gas is a simple theoretical model of a gas. No real gas is truly ideal.' Ideal gas is only a model. Real gases like hydrogen, helium, oxygen and nitrogen come CLOSE to ideal behaviour at low pressure and high temperature, but none obeys PV = nRT exactly at all conditions. Helium and hydrogen behave most ideally because their molecules are small and have very weak attractions.
An ideal gas assumes molecules are point-size and do not attract each other. At LOW pressure the gas is spread out, so molecules are far apart. At HIGH temperature they move fast. In both cases the molecules are far apart, so their own volume is tiny compared to the container and the attraction between them is negligible. NCERT: 'At low pressures or high temperatures the molecules are far apart and molecular interactions are negligible. Without interactions the gas behaves like an ideal one.'
Two assumptions fail. (1) Molecules of an ideal gas have zero volume, but real molecules have a real size, so at high pressure their own volume matters. (2) Molecules of an ideal gas do not attract or repel each other, but real molecules DO attract each other (intermolecular forces). These two facts cause real gases to deviate from PV = nRT.
A real gas deviates most at HIGH pressure and LOW temperature. At high pressure the molecules are packed close, so their finite size and the attractions between them become important. At low temperature the molecules move slowly, so weak attractions can pull them together and even cause the gas to turn into a liquid. This is the opposite of ideal conditions (low P, high T).
Yes. Since an ideal gas obeys PV = nRT exactly, it also follows Boyle's law (PV = constant at fixed T) and Charles's law (V proportional to T at fixed P) perfectly. A real gas only follows them approximately. NCERT notes the agreement between experimental P-V curves and Boyle's law 'is good at high temperatures and low pressures' - again the near-ideal region.
The volume occupied by the molecules contained in 4.5 kg of water at STP, if the intermolecular forces vanish away, is:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
An ideal gas obeys PV = nRT exactly at all pressures and temperatures and has molecules with zero volume and no intermolecular force. A real gas does not obey PV = nRT exactly, because its molecules have real size and attract each other.
Only approximately, and only near low pressure and high temperature. Under those conditions molecules are far apart, so the ideal gas equation gives good results for real gases. At high pressure or low temperature real gases clearly deviate.
Hydrogen and helium behave most like ideal gases. Their molecules are very small and their intermolecular attractions are very weak, so they follow PV = nRT closely over a wide range of conditions.
NEET asks conceptual questions on when PV = nRT can be applied. You must know that ideal gas is a model, that no real gas is truly ideal, and that a real gas approaches ideal behaviour at low pressure and high temperature. Numericals still use PV = nRT when the question ignores molecular size and forces.
Two things: the finite volume of the molecules themselves (they are not points) and the attractive intermolecular forces between them. Both become important when molecules are close together, that is at high pressure and low temperature.
