# Ideal Gas Law Deviations?

Discussion in 'MCAT Discussions' started by ilovelucy, Apr 18, 2007.

1. ### ilovelucy This is Lucy.

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Does anyone have a good way to remember when ideal gases deviate from the ideal gas law?

I was taught that you can always use PV=nRT with the caveat that when it is not ideal, the number you calculate using PV=nRT for a gas deviating is always slightly higher than the "real" number

ideal>real. However, this doesn't always seem to work....

Thanks in advance for the help!

3. ### djones473

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I think its easiest to understand why it happens and then you wont even really have to memorize anything.

Two assumptions of the Ideal Gas Law are:

1. Ideal gases take up no volume
2. Ideal gases don't have intermolecular forces (no attraction)

Real gases deviate at high pressures and low temperatures.

Pressure: When pressure is moderately high, volume is less than predicted because there are intermolecular attractions so the molecules attract together and thus occupy less volume. At VERY high pressures (more than 500atm or so) the SIZE of the gas molecules becomes important and the volume would be greater. This is because at an infinitely high pressure the volume of the gas would approach zero (P and V are inversely related). Of course its not possible for the gases to take up no volume, so thus the ideal gas predicts a smaller volume than actual.

Temperature: When the temperature is lower (close to condensation point) intermolecular forces play an important role and so the volume of the real gas will be less than predicted by ideal gas.

So just remember: Intermolecular forces of attraction make the volume smaller than predicted. This attraction occurs at moderately high pressure (a few hundred atm) and low volume.

and

The size of the molecules makes the REAL volume greater than the volume predicted by the ideal gas law at VERY HIGH pressures.

hope that helps..

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4. ### The Force

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The ideal gas law assumes that the gas molecules have no volume. Since real gases have volumes, real volume is greater than ideal.

It also assumes that there are no intermolecular forces between the gas molecules. Since real gases do have intermolecular interactions, molecules aren't as free to strike the walls of the container. Less molecules striking/and or less force per molecule striking the container leads to a pressure lower than ideal.

Its a simplistic explanation but it should help.

5. ### ilovelucy This is Lucy.

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Thanks to both of you -- both posts were super helpful!

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