Question about real gas behavior vs ideal gas behavior....can anybody help me understand the difference between these two plz?
There are 2 assumptions we make for an ideal gas (neither of which are ever REALLY satisfied, and thus ideal gases do not really exist):
1) gas molecules have no volume
2) there are no attractive forces between molecules
although neither assumption is 100% true, under favorable conditions we can make them more true. the conditions which allow a gas to act more ideal are as follows:
1) low pressures - this condition allows for gas molecules to act more ideal by better satisfying the first assumption above. We know that pressure is inversely related to volume when holding all else constant, thus low pressures allow for a high volume in which the gas molecules can move around in. This greater amount of space for the gas molecules to move around in give the gas molecules themselves smaller volumes when compared to the very large volume of the container, and helps satisfy the first assumption. (Chad says compare 10 marbles bouncing around in a water bottle to 10 marbles bouncing around in a classroom)
2) High temperatures - this gives the molecules more kinetic energy as they move around faster. since the molecules move around at much higher speeds with increased temps, they have less time to participate in any attractive forces with other molecules, thus making the gas act more ideal by satisfying the second assumption above.
The ideal gas law PV=nRT is important because we can determine different variables of an
ideal gas when given others. But, since ideal gases do not really exist, this equation lets us estimate variables for
real gases, and lets us do this best when they act most like ideal gases (under low pressures and high temps).
I think i just went on a rant but I hope this is what you were asking? Looking back at your question now...
ideal gas behavior = molecules have no volume and no attractive forces
real gas behavior = molecules have volumes and interact with each other
