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Getting a little mixed up with Real vs. Ideal gasses. Could someone let me know if im thinking about it correctly?
PV = nRT is the ideal gas law
At STP(1 atm, 273 K), 1 mole of an ideal gas occupies 22.4L of volume. We assume molecular volume is negligible, intermolecular forces are negligible.
[P-(an^2)/(V^2)][V-nb] = nRT for a real gas, which is van der waals equation.
Real gas has a larger volume than an ideal gas because molecular volume is NOT negligible.
Real gas has a lower pressure than an ideal gas because of intermolecular forces(attractive).
However, I was reading up on something about total volume vs free volume and that's where my confusion comes in. When we use the ideal gas law, the V is equal to the total volume which is equal to the free volume. When we use van der waals, does the V refer to the total volume or the free volume?
PV = nRT is the ideal gas law
At STP(1 atm, 273 K), 1 mole of an ideal gas occupies 22.4L of volume. We assume molecular volume is negligible, intermolecular forces are negligible.
[P-(an^2)/(V^2)][V-nb] = nRT for a real gas, which is van der waals equation.
Real gas has a larger volume than an ideal gas because molecular volume is NOT negligible.
Real gas has a lower pressure than an ideal gas because of intermolecular forces(attractive).
However, I was reading up on something about total volume vs free volume and that's where my confusion comes in. When we use the ideal gas law, the V is equal to the total volume which is equal to the free volume. When we use van der waals, does the V refer to the total volume or the free volume?
