- [Voiceover] Imagine you had a monatomic ideal gas
in this cylinder here, an there was this tightly fitted
piston above it that prevented any gas from getting out.
Well we know that the total internal energy
for a monatomic ideal gas is just three
halves P times V or three halves NKT,
or three halves little nRT,
and we know that saying U internal
the internal energy is really just code for the total
kinetic energy of the monatomic ideal gas,
that these are the same thing.
When we talk about internal energy
we're talking about how fast are these particles moving,
in other words, what's the total kinetic energy of
all of these particles added up.
My question is how do we go about
changing the internal energy?
Let's say we wanted to increase the total kinetic energy,
what could we do?
Well, we could say increase the pressure,
or the volume, or the temperature, yeah but I mean
physically, actually, in the lab, what do we do?
And there's basically two ways to change
the internal energy.
If you want to add internal energy,
i.e. get these particles moving faster,
we can heat it up so put this above a flame
or on a hot plate,
and heat will flow into the gas which will
cause these particles to move faster and faster.
That's one way to do it, to add heat.
The other way to do it is to do work on the gas.
I could take this piston and push it down,
and if you push this down hard enough
it will squash this gas together,
and those impacts with this piston while it's moving down
will cause them to start moving faster and faster,
that will also add internal energy to the gas.
So if we wanted to write down a formula that told us
how you could get a change in the internal energy,
if I want to change the internal energy,
Delta U, which is really just saying
changing the kinetic energy,
well there's two ways to do it.
I can add heat.
If I added ten joules of heat
I'd add ten joules to the internal energy,
but I've also got to take into account this work being done,
and so I can do plus the work done on the gas,
and that's it, this is actually the first law
of thermodynamics, it's the law of conservation of energy
it says there's only two ways to add energy,
internal energy to a gas.
Let me talk a little bit more about this work done though
cause getting the sign right is important.
If you're doing work on the gas compressing it
you're adding energy to the gas,
but if you let the gas push up on the piston
and this gas expands pushing the piston up,
then the gas is doing the work,
that's energy leaving the system.
So if the gas does work you have to subtract
work done by the gas.
If the outside force does work on the gas
you add that to the internal energy.
So you got to pay attention to
which way the energy is flowing.
Work done on gas, energy goes in.
Work done by the gas, energy goes out,
and you'd have to subtract that over here.
Let's say the gas did expand.
Let's say the gas in here was under so much pressure
that the force it exerted on this piston was enough
to push that piston upward by a certain amount.
So let's say it started right here and it went up to here
so that piston went from here to there
no gasses escaping cause this is tightly fitting,
but the gas was able to push it up a certain distance d,
how much work was done?
You know the definition of work.
Work is defined to be the force times the distance
through which that force was applied.
So the work that the gas did was F times d,
but we want this to be in terms of thermal quantities
like pressure, and volume, and temperature.
So what could we do?
We can say that this volume,
not only did the piston raise up,
but there was an extra volume generated within here
that I'm going to call Delta V,
and I know that this Delta V has got to equal
the area of the piston times the distance
through which that piston moved
because this height times that area gives me
this volume right in here.
Why am I doing this?
Cause look I can write d as equal to Delta V over A,
and I can take this, I can substitute this formula for d
into here, and something magical happens,
I'll get work equals F times Delta V over A,
but look F over A, we know what F over A is,
that's pressure so I get that the work done by the gas
is the pressure times Delta V.
This is an equation that I like because it's in terms
of thermodynamic quantities that we're already dealing with.
So work done you can figure out by taking P times Delta V
but strictly speaking this is only true if this
pressure remained constant, right?
If the pressure was changing,
then what am I supposed to plug in here,
the initial pressure, the final pressure?
If the pressure's staying constant
this gives you an exact way to find the work done.
You might object and say wait,
how is it possible for a gas to expand and remain
at the same pressure?
Well, you basically have to heat it up
while the gas expands,
that allows the pressure to remain constant
as the gas expands.
And now we're finally ready to talk about heat capacities.
So let's get rid of this,
and heat capacity is defined to be,
imagine you had a certain amount of heat being added.
So a certain amount of heat gets added to your gas.
How much does the temperature increase?
That's what the heat capacity tells you.
So capital C is heat capacity
and it's defined to be the amount of heat
that you've added to the gas,
divided by the amount of change
in the temperature of that gas.
And actually, something you'll hear about often
is the molar heat capacity,
which is actually divided by an extra n here.
So instead of Q over Delta T
it's Q over n the number of moles times Delta T.
Pretty simple but think about it.
If we had a piston in here,
are we going to allow that piston to move
while we add the heat,
or are we not going to allow the piston to move?
There's different ways that this can happen,
and because of that there's different heat capacities.
If we don't allow this piston to move,
if we weld this thing shut so it can't move
we've got heat capacity at constant volume,
and if we do allow this piston to move freely while
we add the heat so that the pressure
inside of here remains constant,
we'd have the heat capacity at constant pressure.
And these are similar but different,
and they're related, and we can figure them out.
So let's clear this away, let's get a nice,
here we go, two pistons inside of cylinders.
We'll put a piston in here,
but I'm going to weld this one shut.
This one can't move.
We'll have another one over here, it can move freely.
So over on this side,
we'll have the definition of
heat capacity, regular heat capacity,
is the amount of heat you add divided by
the change in temperature that you get.
So on this side we're adding heat, let's say
heat goes in, but the piston does not move
and so the gas in here is stuck,
it can't move, no work can be done.
Since this piston can't move,
external forces can't do work on the gas,
and the gas can't do work and allow energy to leave.
Q is the only thing adding energy into this system,
or in other words,
we've got heat capacity at constant
volume is going to equal,
well, remember the first law of thermodynamics said
that Delta U, the only way to add internal energy,
or take it away is that you can add or subtract heat,
and you can do work on the gas.
In this case, Q, if I subtract W from both sides
I get Delta U minus W over Delta T,
but since we're not allowing this piston to move
the work done has got to be zero.
So there's no work done at all so the heat capacity
at constant volume is going to be Delta U
over Delta T, what's Delta U?
Let's just assume this is a monatomic ideal gas,
if it's monatomic we've got a formula for this.
Delta U is just three halves, P times V over Delta T.
That's not the only way I can write it.
Remember I can also write it as three halves NK Delta T
over Delta T, and something magical happens,
check it out the Delta T's go away
and you get that this is a constant.
That the heat capacity for any monatomic ideal gas
is just going to be three halves,
Capital NK, Boltzmann's constant,
N is the total number of molecules.
Or you could have rewrote this as little n R Delta T.
The T's would still have cancelled
and you would have got three halves, little n,
the number of moles, times R, the gas constant.
So the heat capacity at constant volume for any
monatomic ideal gas is just three halves nR,
and if you wanted the molar heat capacity
remember that's just divide by an extra mole here
so everything gets divided by moles everywhere
divided by moles, that just cancels this out,
and the molar heat capacity at constant volume
is just three halves R.
So that's heat capacity at constant volume,
what about heat capacity at constant pressure?
Now we're going to look at this side.
Again, we're going to allow this gas to have heat
enter the cylinder, but we're going to allow this
piston to move up while it does that so that
the pressure inside of here remains constant
and this is going to be the heat capacity
at constant pressure.
Well, again, we're going to get that it's
Q over Delta T, and just like the first law said,
Q has to equal Delta U minus W.
So we get Delta U minus W, the work done,
over Delta T, this time W is not zero.
What's W going to be?
Remember W is P times Delta V.
So this is a way we can find the work done by the gas,
P times Delta V, so this is going to equal
Delta U, we know what that is.
If this is again, a monatomic ideal gas,
this is going to equal three halves nR Delta T
plus this is P times Delta V,
but we have to be careful,
in this formula this work is referring to
work done on the gas,
but in this case, work is being done by the gas,
so I need another negative.
Technically the work done on the gas would be a negative
amount of this since energy is leaving the system.
So that negative cancels this negative
and I get plus P times Delta V,
all of that over Delta T, phew.
So what do we get?
Three halves nR Delta T
plus I want to rewrite P times Delta V,
but I know how to do that.
The Ideal Gas Law says PV equals nRT,
well if that's true, then P times Delta V
is going to equal nR Delta T.
So can rewrite this as nR Delta T divided by Delta T,
almost there, all of the Delta T's go away.
Look what I'm left with.
I'm left with C.
Heat capacity at constant pressure is going to be
equal to three halves nR plus nR,
that's just five halves nR,
and if I wanted the molar heat capacity again
I could divide everything, everything around here
by little n, and that would just give me
the molar heat capacity constant pressure
would be five halves R.
And notice they're almost the same.
The heat capacity at constant volume is three halves nR,
and the heat capacity at constant pressure
is five halves nR.
They just differ by nR.
So the difference between the heat capacity
at constant volume which is three halves nR,
and the heat capacity at constant pressure
which is five halves nR,
is just Cp minus Cv which is nR,
just nR, and if you wanted to take the difference
between the molar heat capacities at constant volume
and pressure, it would just be R.
The difference would just be R
because everything would get divided by the
number of moles.
So there's a relationship, an important relationship.
It tells you the difference between
the heat capacity at constant pressure
and the heat capacity at constant volume.