These are reasonable assumptions for real gases at relatively high pressure  and low temperatures, the gases become less and less  ideal as condensation occurs. If the temperature is high and the pressure is low, so the molecules are moving around fast but its low pressure so there is less bouncing around ,there is some volume taken out and because of the high temperature, it is going so fast there is no time for molecular interactions. Since the pressure is low there are not much collisions, also the container is real big there would not be too many collisions, so the ideal gas law can be applied. When the temperature is low, an ideal gas the temperature wouldn't matter but since there is a low temperature the molecules would move slow, there will be more time therefore colliding. Also since the pressure is so low there would be more collision than bumping into walls. Therefore if the volume of the container, the pressure for the real gas would be lower, if the temperature decreases it could begin condensation therefore not suitable for the equation. If the volume is kept the same as well as the pressure a real gas would require a larger volume. Therefore, the law works for most gases until there is a high temperature and low pressure, things begin to go wrong. 

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Ideal Gases

• No intermolecular interactions- for the ideal gas law we assume that there are no intermolecular interactions 

• The gas molecules volume is negligible

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Real gas

• The particles occupy large volumes 

• When particles collide with one another or the sides of the container, there is energy transfer. 

• Particles to interact with one another through intermolecular forces