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Sometimes I don't know which method I should use to calculate the volume of the gas in a system, I'm not sure if I should use the table values for the specific volume, or if I should calculate by using the Ideal Gases equation. Table values must give me more accurate answers.

But this brings me the question: at certain conditions, gases will act like ideal gases, so the values from the ideal equation are similar to the tabled values, however, at other conditions they are not, because gases stop acting like ideal gases.

How can I know when do they stop behaving like ideal gases? Does it have to do if they are in the SuperHeated region, or Saturated region? Can they even stop being ideal gases on the saturated region?

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  • $\begingroup$ No physical gas ever behaves exactly like an ideal gas; the ideal gas law (like almost every law of physics) is an approximation that holds under certain idealized assumptions. The degree to which a gas sample deviates from the ideal gas law will grow and shrink continuously depending on the state parameters, but the deviations will never go away entirely; there's nowhere that it "starts" or "stops" behaving like an ideal gas. $\endgroup$ Commented Apr 24, 2017 at 18:40
  • $\begingroup$ @MichaelSeifert ok, there is always a deviation. Could I say that in saturated region the deviation would be smaller than in SuperHeated region? Or it simply doesn't matter, maybe for some parameters of saturated region it can have more deviation than some parameters of superheated region? $\endgroup$ Commented Apr 24, 2017 at 18:57
  • $\begingroup$ @MichaelSeifert you also have to be careful not to attribute non ideal behavior to causes like heat transfer. Other physical processes need to be considered $\endgroup$
    – docscience
    Commented Apr 24, 2017 at 19:11
  • $\begingroup$ Are you familiar with the Law of Corresponding States, and the graph of compressibility factor Z as a function of reduced temperature and pressure, where the reduced temperature is the actual temperature normalized by the critical temperature and the reduced pressure is the actual pressure normalized by the critical pressure? $\endgroup$ Commented Apr 24, 2017 at 23:45
  • $\begingroup$ @yes, I wanted to know if they are related with their region on the graph. $\endgroup$ Commented Apr 25, 2017 at 0:00

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One way to determine if you need to ideal gas behavior is to calculate the Compressibility Factor using the virial equation, a model that accounts for interactions of molecules (the ideal gas law assumes no interactions other than perfectly elastic) according to the level of temperature. For $Z$ too different than 1.00 or according to your accuracy requirements, you have a deciding metric and a correction factor for the ideal gas equation.

$$PV=ZRT$$

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  • $\begingroup$ The key phrase here is (as always) "according to your accuracy requirements". It's always about the level of deviation you can tolerate. $\endgroup$ Commented Apr 24, 2017 at 18:46
  • $\begingroup$ @dmckee in my application area the ideal gas law always applies for my purpose. Low pressure, moderate temperature $\endgroup$
    – docscience
    Commented Apr 24, 2017 at 18:55
  • $\begingroup$ ok, so I'll use only the table. What about the phases of the gas? Does it have to do with the deviation of "ideal gas behaviour"? $\endgroup$ Commented Apr 24, 2017 at 19:00
  • $\begingroup$ @VitorAguiar68 Ideal gasses don't exhibit condensation to liquids. The usual choice for a simple model that includes condensation is the van der Waals gas. $\endgroup$ Commented Apr 26, 2017 at 20:40

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