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Why is potential energy = 0 for ideal gases and not for non ideal gases ?

The reason is in my textbook that average between the atoms increases or decreases due to bonds being broken.So , my Q is that in case of ideal gases. Let us take two gases , $H_2$ and $He$. So , here. For hydrogen , I can say that bonds of H and H can break and average separation can increase or decrease.p even when if I consider hydrogen as ideal gas. I checked online also if I can consider $H_2$ has ideal gas and it said yes. Whereas , in case of He. I am not getting with what He would form bond with and average separation between atoms would change.

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  • $\begingroup$ Search for fugacity and real gas theory. It becomes significant for high pressures and low temperatures (so has quite a bit of industrial application). Also note that even as an ideal gas, H2 and He differ in behavior because of the different gamma (ratio of specific heats) due to He being monoatomic. That will, for example, change the relationship between T and P when fast (adiabatic) expansion or compression happen. I don't know enough to say about potential energy. $\endgroup$
    – Pete W
    May 16 '21 at 14:46
  • $\begingroup$ Ok. Thanks a lot. More to my Q , the example I took about He and H2. I didn’t understand whether I choose the bond breaking as I explained correct or not. Could you help in that $\endgroup$
    – Rider
    May 16 '21 at 14:51
  • $\begingroup$ I'd look for lecture notes for a first year grad school course in physical chemistry. Here's a free text ... you may want to back up to earlier chapters too. Note that the P-Chem and Engineering Thermodynamics curriculums break it down it in a different way. $\endgroup$
    – Pete W
    May 16 '21 at 14:58
  • $\begingroup$ Is there are potential energy change only during phase change $\endgroup$
    – Rider
    May 19 '21 at 15:41
  • $\begingroup$ see Gibbs Free Energy. If I understand correctly, it is conserved in a reversible process (adiabatic + isentropic), but many processes are not. (adiabatic arguably depends on definition of system boundary, entropy you just can't fix) $\endgroup$
    – Pete W
    May 19 '21 at 15:52
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The text below is excerpted from this wiki articke, hope it helps.

"Deviations from ideal behavior of real gases

Since the ideal gas law neglects both molecular size and inter molecular attractions, it is most accurate for monatomic gases at high temperatures and low pressures. The neglect of molecular size becomes less important for lower densities, i.e. for larger volumes at lower pressures, because the average distance between adjacent molecules becomes much larger than the molecular size. The relative importance of intermolecular attractions diminishes with increasing thermal kinetic energy, i.e., with increasing temperatures. More detailed equations of state, such as the van der Waals equation, account for deviations from ideality caused by molecular size and intermolecular forces."

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