I am confused about this question:
Explain why low temperature and high pressure will cause a real gas to deviate from ideal gas behaviour.
A detailed explanation is appreciated. Cheers!
At low temperatures, gas molecules have less kinetic energy on average, so intermolecular forces have a stronger effect overall on molecular motions. As the intermolecular forces can no longer be neglected, the gases deviate from ideal gas behaviour.
At high pressures (i.e small space, or really high temperatures), gas molecules interact with each other a lot more as the space is relatively small for the amount of energy they have (end up getting closer to each other). Again, intermolecular forces have a greater effect overall.
Or, you can think about it this way. Reducing the temperature and pressurising a gas are the only ways of changing the state of a gas (unless you want to consider supercritical fluids and plasmas as different states of matter...but let's not go there). This implies that in reality, intermolecular forces must be able to hold a gas together if we keep reducing the temperature or pressurise the gas.
Explain why both factors of low temperature and high pressure will cause a real gas to deviate from ideal gas behaviour.
My answer:
That low temperature should decrease the random movement of gas particles, as there is less kinetic energy generated from random collisions, and thereby reduce overall pressure.
When low temperature of a real gas induces high pressure, the real gas deviates from ideal gas behaviour.
Can you help me improve this?
(just because this is the updated version of the quote)
Low temperatures do not induce high pressure. Just saying. You've said that yourself, so you've contradicted yourself.
There is a difference between the two cases. Also, low temperatures do not "generate" less kinetic energy; the particles just have less kinetic energy overall. Your answer doesn't really tell me much.