July 06, 2020, 08:23:18 PM
Forum Rules: Read This Before Posting

Topic: Explaining the difference in volume of ideal and real gasses  (Read 541 times)

0 Members and 1 Guest are viewing this topic.

Offline Traumatic Acid

  • Regular Member
  • ***
  • Posts: 42
  • Mole Snacks: +2/-3
Greetings all.
So I've been studying gasses, and am writing a report after an experiment where we demonstrated that the gas constant for O2 is not the same as R ideal. Pretty straight forward. I know that the real gas law (Van der Waal's gas equation) takes into consideration the particle size of gasses and inter-molecular forces.
I recently found a graph (hopefully attached) that outlines the difference in volume of an ideal and a real gas (not sure which gas was used). It's pretty obvious that the real gas assumes a smaller volume than the ideal gas at the same temperature and pressure. My question is if that is due to the Van der Waal forces pulling the particles together? and If does anyone know of a good reference I can cite which explains that. For some reason the text book I have doesn't actually go into detail why real gasses occupy a smaller volume.


Offline Enthalpy

  • Chemist
  • Sr. Member
  • *
  • Posts: 3486
  • Mole Snacks: +287/-57
Re: Explaining the difference in volume of ideal and real gasses
« Reply #1 on: May 10, 2019, 07:19:16 AM »
[...] the real gas law (Van der Waal's gas equation) [...]


Van der Waals only provided the first equation with correction to the ideal gas law. His equation is by no means the real gas law. It isn't even very accurate. Dozens of other equations were proposed that are less bad, in all circumstances, or in some domain of pressure and temperature, or for some gases.

Yes, Van der Waals forces attract molecules almost always, and this reduces the gas' volume since at any time, a fraction of the amount of molecules stick together. The effect should not be derived from a mean distance.

An example, for nitrogen, on page 8 there

Sponsored Links