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Since the pressure of a real gas is less than that of the ideal gas and its volume is more than that of ideal gas, I am assuming that the real gas is difficult to compress in comparison to an ideal gas.

Is my logic right? Also I'm looking for a clearer explanation!

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Let's use the van Der Waals equation for this. (Real gasses also described elegantly by the Viral equation)

enter image description here

  • You were false that "pressure of a real gas is less than that of the ideal gas". It is greater. Also the volume is smaller.

  • Thus if it has a higher pressure and smaller volume, it typically is harder to compress (i.e. the compression factor increases) but it also depends on temperature. Here is a nice plot of compression factors vs temp for some common gasses from the cambridge chem guide.

    http://www.chemguide.co.uk/physical/kt/realgases.html enter image description here

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  • $\begingroup$ Your assessment of the values of the pressure of real and ideal gases seems backward here. If Pideal = Preal + an^2/V^2, then Preal must be less than Pideal, since the second term is positive. That makes sense, since the IMFs are pulling the gas molecules together, causing them to hit the walls of the container less forcefully. Similar reasoning yields a smaller Videal relative to Vreal, since the ideal volume does not account for the additional space taken up by the gas particles. $\endgroup$ – Jason Patterson Sep 22 '14 at 19:39

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