Why do we add, and not subtract, the correction term for pressure in [Van der Waals] equation?
Since the pressure of real gases is lesser than the pressure exerted by (imaginary) ideal gases, shouldn’t we subtract some correction term to account for the decrease in pressure?
I mean, that’s what we have done for the volume correction: Subtracted a correction term from the volume of the container V since the total volume available for movement is reduced.
asked Sep 30 ’16 at 15:20
— Physics Stackexchange
Ideal gas law:
However, since in a real gas, there are attractions between molecules, so the measured value of pressure P is smaller than that in an ideal gas:
Also, since the gas molecules themselves occupy some space, the measured value of the volume V is bigger that the real gas really has:
If we substitute onto the LHS, since , the LHS will be smaller than the RHS:
So in order to maintain the equality, a correction term to the pressure must be added:
If we substitute onto the LHS, since that volume is bigger that actual volume available for the gas molecules to move, the LHS will be bigger than the RHS:
So in order to maintain the equality, a correction term to the pressure must be subtracted:
In other words,
— Me@2018-05-13 03:37:18 PM
Why? I still do not understand.
— Me@2018-05-13 03:22:54 PM
The above is wrong.
The “real volume” has 2 possible different meanings.
One is “the volume occupied by a real gas”. In other words, it is the volume of the gas container.
Another is “the volume available for a real gas’ molecules to move”.
To avoid confusion, we should define
Or even better, avoid the terms and altogether. Instead, just consider the relationship between and that between .
Whether is bigger or smaller than ultimately depends on the assumptions and definitions used in the derivation of the ideal gas equation itself.
— Me@2018-05-13 04:15:34 PM
2018.05.13 Sunday (c) All rights reserved by ACHK