Why do real gases stop obeying PV = nRT at very high pressure or very low temperature?
The ideal-gas law assumes molecules take up no space and don't attract each other.
The ideal-gas law assumes molecules take up no space and don't attract each other. At high pressure the molecules are crammed close, so their own volume is no longer negligible; at low temperature they move slowly enough for intermolecular attractions to matter. Both make the real behaviour drift from PV = nRT, which is why the van der Waals equation adds correction terms for volume and attraction.
Key point
Real gases deviate when molecular volume and attractions matter — high P, low T.
Common mistake
assuming PV = nRT holds under every condition.
Memory tip
ideal gas fails when molecules are crowded (high P) or sluggish (low T).
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