Why during free-expansion, temperature changes for real-gas while the same doesn't happen so for ideal gas?

The reason that the temperature changes in an expansion/compression of a real gas at constant internal energy can be quite intuitively understood.

Suppose that the interactions between the gas molecules are repulsive. If we compress the gas while fixing its energy, the average intermolecular distance decreases, leading to the increase in the potential energy. Accordingly, the kinetic energy, as well as the temperature, has to decrease (energy conservation).

Similar lines of reasoning apply for other cases, i.e., expansion or compression of a gas with repulsive or attractive interactions.

The internal energy of a real gas depends on both temperature and pressure. So, if U remains constant and pressure changes, the temperature must change. In the ideal gas limit of very low pressures, the pressure dependence of real gases weakens and approaches zero.