Why do real gases not behave exactly like ideal gases?

The molecules have (a) volume and (b) attractive forces

Step-by-step explanation:

At ordinary conditions, the molecules are so far apart that the gases behave almost ideally.

However, if you use high pressure and/or low temperatures, you force the molecules to be close together.

There are two competing effects:

The attractive forces become much stronger at close distances, so the volume is less than that predicted by the Ideal Gas Law. The volume of the molecules becomes a significant portion of the volume of the container. The molecules have less volume in which to can move around, so the pressure is higher than that predicted by the Ideal Gas Law.

Answer: Intermolecular forces of Attraction

Explanation: Ideal gases are the hypothetical gases whose prime formula is derived from the kinetic molecular theory of gases.

These gases have negligible intermolecular forces of attraction.

Real gases have intermolecular forces of attraction.

Real Gases are suppose to behave like ideal gases at high temperature when the the kinetic energy of the gas molecules are its peak and at low pressure when there will no intermolecular forced of attraction as volume will be high.