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.