States of Matter (Gases and Liquids)NEET Chemistry · Class 11 · NCERT Chapter 13

An ideal gas is a hypothetical gas that obeys PV = nRT perfectly at all temperatures and pressures. It assumes: (1) gas molecules have no volume (point masses), (2) there are no intermolecular forces between molecules, and (3) all collisions are perfectly elastic. Real gases deviate from ideal behaviour because: (a) at high pressures, the actual volume of molecules becomes significant compared to container volume, (b) at low temperatures, intermolecular attractive forces (van der Waals forces) become important and molecules are not truly non-interacting. Real gases approach ideal behaviour at high temperature and low pressure (conditions where molecules are far apart and moving fast).

The van der Waals equation is (P + an²/V²)(V − nb) = nRT. The constant "a" accounts for intermolecular attractions: an²/V² is the pressure correction (molecules attract each other, reducing the pressure on the walls below the ideal value). A larger "a" means stronger intermolecular attractive forces. The constant "b" accounts for the finite volume of gas molecules: nb is the volume correction (the actual free volume available for movement is less than V because molecules occupy space). A larger "b" means larger molecular size.

Critical temperature (Tc) is the temperature above which a gas cannot be liquefied by pressure alone, no matter how high the pressure is applied. Above Tc, the kinetic energy of molecules is too high for intermolecular forces to hold them in a liquid state. To liquefy a gas, you must first cool it below its Tc, then apply pressure. For example, CO₂ has Tc = 31.1 °C — it can be liquefied at room temperature by applying pressure. But N₂ has Tc = −147 °C, so you must cool it below −147 °C before it can be liquefied. This is why the liquefaction of air (N₂, O₂) requires cooling by the Linde process.

Graham's law states that the rate of diffusion (or effusion) of a gas is inversely proportional to the square root of its molar mass at the same temperature and pressure: rate ∝ 1/√M. For comparing two gases: rate₁/rate₂ = √(M₂/M₁). Lighter gases diffuse faster. For example, H₂ (M = 2) diffuses 4 times faster than O₂ (M = 32) because √(32/2) = √16 = 4. Graham's law has applications in isotope separation (uranium hexafluoride enrichment) and in explaining why a balloon filled with H₂ deflates faster than one filled with CO₂.

Surface tension is the force per unit length (or energy per unit area) acting along the surface of a liquid, arising from the unequal forces experienced by surface molecules compared to interior molecules. Interior molecules are attracted equally from all sides; surface molecules experience a net inward pull. This makes the surface contract to a minimum area (spherical droplets, bubbles). Surface tension (γ) has SI unit N/m or J/m². It decreases with increasing temperature because thermal energy reduces the cohesive forces between molecules. Surfactants (soaps, detergents) reduce surface tension by disrupting intermolecular attractions at the surface.

In liquids, viscosity arises from intermolecular attractive forces between layers of molecules. As temperature increases, molecules gain more kinetic energy, overcoming these attractions more easily, so viscosity decreases. In gases, viscosity arises from momentum transfer between gas molecules colliding between layers. As temperature increases, molecules move faster and collide more frequently, transferring more momentum between layers and increasing viscosity. So for liquids: viscosity decreases with temperature; for gases: viscosity increases with temperature. This is a key distinction for NEET.