Two separate containers of equal volume V contain ideal gases A and B at the same temperature T. The molar mass of A is twice that of B. The ratio of the densities of A and B ($\rho_A / \rho_B$) is:

At constant temperature, the pressure of an ideal gas is doubled. What happens to its volume?

Which of the following gases is expected to deviate most from ideal behavior at high pressure?

At what temperature will the volume of a gas at 0\,^\circ C double itself, pressure remaining constant?

In the equation of state of an ideal gas $pV = nRT$ the value of the universal gas constant would depend only on

Two separate containers of equal volume V contain ideal gases A and B at the same temperature T. The molar mass of A is twice that of B. The ratio of the densities of A and B ($\rho_A / \rho_B$) is:

Gas equation $PV = nRT$ is obeyed by:

A gas at 298 K is shifted from a vessel of $250\,\,{\rm{c}}{{\rm{m}}^3}$ capacity to that of 1 L capacity. The pressure of the gas will

The equation of state corresponding to 8g of ${{\rm{O}}_2}$ is

When a sample of gas is compressed at constant temperature from 15 atm to 60 atm, its volume changes from $76\,\,{\rm{c}}{{\rm{m}}^3}\,\,{\rm{to}}\,\,20.5\,\,{\rm{c}}{{\rm{m}}^3}$. Which of the following statements are possible explanations of this behaviour? 1. The gas behaves non-ideally. 2. The gas dimerises. 3. The gas is absorbed into the vessel walls.

1.0 L of ${N_2}$ and 7/8 L of ${O_2}$ at the same temperature and pressure were mixed together. What is the relation between the masses of the two gases in the mixture?