While studying Gaseous state in books, there was a common question which I had difficulty answering -

Why do real gases deviate from ideal behavior (3 points)?


closed as off-topic by orthocresol, Jan, F'x, M.A.R. ಠ_ಠ, Klaus-Dieter Warzecha Aug 31 '16 at 10:31

This question appears to be off-topic. The users who voted to close gave this specific reason:

If this question can be reworded to fit the rules in the help center, please edit the question.

  • 3
    $\begingroup$ The (3 points) part of your question is a pretty good giveaway that you copied this question from somewhere. Could you share what your thoughts on this question are? $\endgroup$ – Ben Norris Aug 31 '16 at 10:31
  • $\begingroup$ Actually I was studying my textbook which is the NCERT for class 11 in India. In which there is this question. I added 3 points because my teacher asked the same question - Give 3 points while taking a viva examination. $\endgroup$ – S Aditya Aug 31 '16 at 12:00
  • $\begingroup$ Further it didn't click me that I had to write about the assumptions that were taken by Maxwell and Boltzmann while formulating the kinetic theory of gases... I was trying to relate this question to the trend of compressibility factor of real gases vs ideal gases. $\endgroup$ – S Aditya Aug 31 '16 at 12:06

Real gases deviate from ideal behaviour because their particles (atoms for inert gases or molecules) occupy some finite space and do exert interactive forces among them. Completely ideal behaviour is hypothetical because of the reasons above. At low pressure and high temperature, real gases behave approximately as ideal gases. In ideal behaviour, gas particles don't occupy space and do not have any interaction, as assumed in the kinetic theory of gases. But in reality this is not the case: we get errors by applying the ideal gas law. That's why van der Waals corrected it by introducing suitable constants.


Not the answer you're looking for? Browse other questions tagged or ask your own question.