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MY textbook says : Real gases show Ideal Behavior under low pressure ,high temperature.

So I wanted to know what happens to real gases at high pressure ,low pressure Do they also show ideal behavior under those conditions?

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  • $\begingroup$ In order to make comparisons you have to specify other conditions i.e temperature and volume. so do the other conditions remain same or not? $\endgroup$
    – user99496
    Oct 3 '20 at 12:09
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When real gases are at high temperature, the kinetic energy prevents any gas particles from interacting via intermolecular forces. With low pressure, the gas particles are separated enough that the intermolecular forces are sparse, therefore, giving rise to the ideal behavior since ideal gases are defined as non-interacting particles.

When real gases are at high pressure or low temperature, they deviate a lot from ideal gases.

If you compress a gas enough, you can induce a phase change.

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The van Der Waal equation is $$\left(p+\frac a{V_\mathrm m^2}\right)(V_\mathrm m-b)=RT$$ Here $V_\mathrm m$ is molar volume. When pressure is low and temperature is very high, we can qualitatively say that the molar volume will be very large. Due to this the volume occupied by the molecules (given by $b$) becomes insignificant.

The pressure is low but the molar volume $V_\mathrm m$ is very large and thus the term $\frac a{V_\mathrm m^2}$ comes out to be much much lower than the already low pressure. $$p+\frac a{V_m^2}\approx p$$ and $$V_\mathrm m-b\approx V_\mathrm m$$ Thus the gas shows ideal behavior at low pressure and high temperature.

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