The main difference between an ideal gas and a real gas is that an ideal gas follows all the assumptions of the kinetic theory of gases, while a real gas deviates from these assumptions.
The assumptions of the kinetic theory of gases are: 1) gas particles are in constant, random motion; 2) gas particles are point masses with no volume; 3) gas particles do not interact with each other; 4) gas particles only interact with the walls of the container through elastic collisions; and 5) the average kinetic energy of gas particles is directly proportional to the temperature.
Real gases deviate from the assumptions of the kinetic theory in several ways. Firstly, real gas particles have volume and therefore do not behave as point masses. Secondly, real gas particles do interact with each other, leading to intermolecular forces. Thirdly, real gases can also deviate from ideal behavior at high pressures and low temperatures.
The properties of an ideal gas, such as pressure, volume, and temperature, follow the ideal gas law exactly at all conditions. On the other hand, the properties of a real gas can deviate from the ideal gas law due to the factors mentioned in the previous questions, such as intermolecular forces and non-zero volume of gas particles.
Under certain conditions, such as low pressures and high temperatures, a real gas can behave closely to an ideal gas. This is known as the ideal gas limit. However, at higher pressures and lower temperatures, the deviations from ideal behavior become more significant and the gas cannot be considered an ideal gas.