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## Homework Statement

Ok, firstly, I apologise for posting something which is probably trivial to any physics student, but my understanding of physics is pretty poor, so baby steps would be appreciated!

## Homework Equations

PV=nRT

E = 1/2 mv

^{2}

## The Attempt at a Solution

My understanding of it so far:

At the molecular level, a point particle of an ideal gas (ignoring rotational/vibrational components and forces between particles) has kinetic energy equal to:

E = 1/2 mv

^{2}

But particles are constantly colliding with each other as well as the walls of the container. Hence an expression for average kinetic energy tells us more information. My course notes give this as:

Ebar = 1/2 mv(bar)

^{-2}

an explanation as to what that means and how they got there would be nice :S

And then there's another jump to average molar kinetic energy of an ideal gas, which is given in a different form by different sources. If I have reasoned correctly, the form in my notes for E

_{average, molar}= 3/2 RT = 3/2 PV when n=1 in the ideal gas equation.

but how did they jump from 1/2 mv

^{2}to this? No doubt avogadro's constant comes into play, but its clearly not as simple as taking the second expression and multiplying.

And finally, the jump to graham's law. I understand and accept that kinetic energy is only dependant on the temperature; hence two gases at equal temperatures have equal kinetic energy.

If the rate of diffusion/effusion is a velocity, then how did we get to rate

_{x}= constant/sqrt(molar mass

_{x})?

thanks!

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