Real gases?

Is it possible that "P" is on the wrong side?

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In the ideal gas law, molecules of gas are treated like singular points (no volume) with only the repulsion of the outer electron shells (against each other) providing the pressure = the sum of all repulsions.

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Van der Wall's equation corrects for the fact that real molecules are not singular points (they have a physical volume that is not zero), and for the fact that there may be SOME small amount of attraction between molecules (it is not only electron repulsion). The corrections are not big but, in some cases, they may make a difference.

The full equation, for pressure, is

P = nRT/(V-nb) - a(n/V)^2

P is the pressure of the gas, depending on the elements on the right

V = volume to which the gas is confined

n = number of moles of gas

a = attraction between individual gas particles

b = average volume of individual gas particles

R = ideal gas constant = 0.08206 L·atm/mol·K

T = absolute temperature

On the right of the equal sign, there are two terms. Your question deals with the first term (correcting for the fact that molecules have a volume instead of being ideal points).

The second term deals with the correction for the bit of attraction that balances part of the repulsion. The more attraction there is between molecules, the less pressure you need to keep the gas confined (that is why the term is negative).

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The equation, AS YOU HAVE IT WRITTEN, would work only if there were no attraction (in other words, when you set a=0)

P = nRT/(V - nb) - 0

P = nRT/(V - nb)

move the denominator:

P(V - nb) = nRT

or

(V - nb) = nRT/P

I had some from the burritos last nite but I'm good now

Ideal gases vs real gases ....

Ideal gases follow the tenets of the kinetic molecular theory, which says, among other things, that gases do not attract each other and that gas molecules are points, with no volume. Neither is true of real gases. The van der Waals equation takes that into account. This is where the van der Waals forces (*) come from.

P = nRT / (V - nb) - (an²/V²)

The variable "a" and the "(an²/V²)" term take into account the attraction between molecules. "(V - nb)" corrects for the finite volume of molecules.

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The three van der Waals forces:

1. Keesom forces (dipole-dipole attraction)

2. Debye forces (induced attraction)

3. London dispersion forces (present between all molecules)

========== Follow up ==========

You have written, " ( V − n b ) = n R T P "

If you leave out the (an²/V²) correction for the attraction between molecules, then the van der Waals equation becomes:

P = nRT / (V - nb)

Solving for (V - nb) gives you:

(V - nb) = nRT / P

Therefore, P is in the denominator, not the numerator. The equation you wrote is not correct..

volume (of something) - noble boron= noble boron* (recipricate of terrestrial plate)

As opposed to fictional gases?