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Van der Waals had stated that at high pressure, pressure exerted by ideal gas on the walls of container is more than real gas.

However in this graph, one can clearly see that, the at high pressure volume of real gas tends to be more than volume of ideal gas. Hence real gas should be exerting more pressure on the walls of the container, as it occupies more volume. So why does this not hold true?

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It's because in ideal gases there are no intermolecular forces whereas in a real gas the force with which it collides is reduced due to intermolecular forces.Now for the volume occupying part ideal gas particles are point sizes but in a real gas the molecules do occupy some volume.That's why some corrections were made to volume and pressure in the ideal gas equation. Note that the above were assumptions of the Kinetic theory of gases.

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  • $\begingroup$ But see the graph above, it says that volume occupied by the Real gas is more than the volume occupied by ideal gas . So thus , the pressure exerted by real gas must be more than that of ideal gas , simply because it occupies more volume. $\endgroup$
    – Jajati
    Sep 18, 2017 at 6:04
  • $\begingroup$ Can you tell me how more volume means more pressure? Because they are inversely proportional its if there's very high pressure wouldn't the molecules clump together(i know repulsion) and the volume decreases Aka Boyle's law? $\endgroup$
    – WhiteHole
    Sep 18, 2017 at 7:05
  • $\begingroup$ In a same container( keeping the volume constant) , a real gas will tend to occupy more volume than ideal gas. So since we are restricting it's volume, the real gas should exert more pressure on the walls of the container. It's that simple $\endgroup$
    – Jajati
    Sep 18, 2017 at 7:20
  • $\begingroup$ i don't quite understand first you say volume is constant , then real gas occupies more volume ? Is volume property of the container or molecule what's the reasoning behind "real gas "should" exert more pressure"? $\endgroup$
    – WhiteHole
    Sep 18, 2017 at 7:29
  • $\begingroup$ See the graph properly and the line I drew to mark the pressure. At a constant pressure, the volume occupied by real gas is more than that of ideal gas. Now if we try to confine the both gases in a container, the gas that tends to acquire larger volume will tend to exert more pressure. ( Try to draw a line from the volume axis to the graph ,it will cut the real gas and ideal gas graphs at different points. It means that at a constant and confined volume real gas exerts more pressure on wall than ideal gas) $\endgroup$
    – Jajati
    Sep 18, 2017 at 7:35

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