Gaseous States Test - 6

r = CP/CV​​​ = 5​/3

= 1.66 

(For Monoatomic as He,Ne,Ar)

Rate of diffusion α 1/√m

white fumes first formed near hydrogen chloride.

In S.I unit, R=8.314J/mol/Kelvin

1J=10⁷ergs

∴ R = 8.314×10⁷ ergs/mol/kelvin

And R=0.0821L atm/mol/K

1cal = 4.18J,R = 1.987Cal/mol/K. = 8.31×10⁷dynescm/mol/K

Gases deviate from the ideal gas behaviour because their molecules have forces of attraction between them.

Assumptions of kinetic theory of gases : 

1). The volume occupied by gas molecules is negligibly small as compared to the volume occupied by the gas.

The first assumption is valid only at low pressures and high temperature, when the volume occupied by the gas molecules is negligible as compared to the total volume of the gas. But at low temperature or at high pressure, the molecules being in compressible the volumes of molecules are no more negligible as compared to the total volume of the gas.

2). The forces of attraction between gas molecules are negligible.

The second assumption is not valid when the pressure is high and temperature is low. But at high pressure or low temperature when the total volume of gas is small, the forces of attraction become appreciable and cannot be ignored.

On increasing pressure, the volume decreases and density increases. So molecules get closer to each other hence mean free path also decreases.

Here volume is constant and mass of H₂ is fixed so the no. of moles of the gas do not change. As temperature increases the pressure also increases. therefore the rate of collision among the gas molecules and their energy also increases.

According to ideal gas equation, pV = nRT

If T is constant and mass is constant, so number of moles (n) also constant [n = mass /molar mass ]

therefore

pV = constant

Real gases behave like ideal gases at high temperature and low pressure and  maximum duration is observed when the condition is reversed means at low  temperature and high pressure

As we know, the ideal gas equation is PV = nRT

Now since N2 is removed the total pressure would be equal to the partial pressure exerted by the O2 gas and vice versa. Therefore, the pressure after the N2 of any gas would be equal to P/2 mm of Hg.

An ideal gas is a hypothetical gas whose pressure, volume and temperature behaviour is completely described by the ideal gas equation. Actually no gas is ideal or perfect in nature. All gases are real gases. Real gases do not obey the ideal gas laws exactly under all conditions of temperature and pressure.

Real gases deviate from ideal behaviour because of mainly two assumptions of "Kinetic theory of gases".

(i) The volume of a gas particle is negligible compared to the volume of the container (while the real gas particle has some significant volume).

(ii) There is no interaction between gaseous particles (while attraction forces exist between real gas particles).

So at lowest pressure and highest temperature, a real gas most closely approaches the behaviour of an ideal gas.