ideal gas law deviations

[banner18]

I've been reading the past threads about gas law deviations but im only getting more confused....

I get that deviations only occur at low temperatures or high pressure.

SO, the law ignores volume and attractive forces of the gases and these will be the variables that deviate from the ideal conditions.

Pressure: if you decrease volume a lot (keeping temp. const.) you will have a lower than expected pressure reading due to the attractive forces of the gas, lowering collisions against the container.

Volume: if you increase pressure a lot (keeping temp. const) you will have lower than expected volume because the gas molecules take up a volume that is significant at high pressure, the space that the gas molecules occupy will take away from the volume of the container??? Or will it increase more than expected because the gases volume needs to accounted for on top of the containers volume???

If you decrease temp. volume will be lower than expected due to intermolecular attractions. Will this mean that pressure will decrease as well since temp. and pressure are proportional?? Can you predict how volume would change if you decrease temp and hold pressure constant?

CONFUSED

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[movax]

I've been reading the past threads about gas law deviations but im only getting more confused....

I get that deviations only occur at low temperatures or high pressure.

There are two assumptions about ideal gas to attack here:
1. gas molecules themselves occupy no volume,
2. intermolecular forces are 0.

Ideal gases are most ideal at high temperatures and low pressure (think about 273K and 1atm). Low T and high P cause deviations.

Increasing the pressure will result in the gas occupying slightly more volume than it does at ideal/STP because we now have to take into account the volume of the gas molecules themselves. (I think I just did this on CBT 4R or 7, can't remember which...).

That is van der Waals' factor b and attacks assumption #1.

Now playing with temperature, we attack assumption #2. If there were no intermolecular forces, gases could never condense. At low temperature, gases will condense to form liquids, which is obviously a deviation. Additionally, with intermolecular forces in play, real pressure will be slightly less than ideal pressure.

van der Waals' factor a corrects for real pressures being lower than ideal pressures when we think about temperature.

So, high pressures keeping all else constant will result in a larger than expected volume. Low temps keeping all else constant will result in a lower than expected pressure.

 

[banner18]

So, high pressures keeping all else constant will result in a larger than expected volume.

In the gas law, the Volume refers to the volume of the container that the gas is in, not the gas it itself. So when you have high pressure the volume of the container will be less than expected because the volume of the gas that is occupying the container cant be ignored. But the intermolecular attraction forces of the gases being so close together will cause the volume of the container to be higher since the gas particles arent moving around as freely.

So is volume lower than expected or higher than expected?

 

[movax]

In the gas law, the Volume refers to the volume of the container that the gas is in, not the gas it itself.

😕 Gases will always expand to fill their container, yes, but this statement confuses me a bit.

So is volume lower than expected or higher than expected?

Every source I've looked at indicates that the volume occupied by gas is greater than expected when temperature is held constant and pressure increased from "low" to "high". Look at #25 on AAMC CBT 7R.

 

[justhanging]

In the gas law the volume represents the volume of the gas, but if the gas is in a container then its the volume of the container.

 

[MedChallenge]

I'm confused on this, too, but Rabolisk has a good explanation here:
http://forums.studentdoctor.net/showthread.php?t=804701