Why do real gases not behave exactly like ideal gases?

The theory assumes that collisions between gas molecules and the walls of a container are perfectly elastic, gas particles do not have any volume, and there are no repulsive or attractive forces between molecules .

At high pressure and low temperature, real gas deviates from ideal gas. Compressibility factor is used to measure the deviation of real gas from ideal gas.

The compressibility factor is "n" which equals PV/RT.This is equal to one in case of ideal gas and this condition does not hold good for real gas.

At high pressures, the molecules in the gas tend to become closer, resulting in greater intermolecular forces. Internal molecular force holds these molecules together,prevents movement and collision and thus reduces the pressure of the gas much below than that of an ideal gas.

At low temperatures, the kinetic energy of the gas molecules also becomes less thus this results in less collision of the gas molecules and thus finally this results in the decrease in pressure of the gas molecules which would be much lower than that of an ideal gas.

For an ideal gas, the translational kinetic energy is 3/2kT. This condition does not hold good for a real gas.

a gas, having a molecular weight of 13 kg/kmol and a kinematic viscosity of 0.25 cm2/s, flows through a pipe 0.25 m internal diameter and 5 km long at the rate of 0.4 m3/s and is delivered at atmospheric pressure.

what is the question here. i think you messed up.