This introductory physical chemistry course examines the connections between molecular properties and the behavior of macroscopic chemical systems.

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From the course by University of Minnesota

Statistical Molecular Thermodynamics

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This introductory physical chemistry course examines the connections between molecular properties and the behavior of macroscopic chemical systems.

From the lesson

Module 1

This module includes philosophical observations on why it's valuable to have a broadly disseminated appreciation of thermodynamics, as well as some drive-by examples of thermodynamics in action, with the intent being to illustrate up front the practical utility of the science, and to provide students with an idea of precisely what they will indeed be able to do themselves upon completion of the course materials (e.g., predictions of pressure changes, temperature changes, and directions of spontaneous reactions). The other primary goal for this week is to summarize the quantized levels available to atoms and molecules in which energy can be stored. For those who have previously taken a course in elementary quantum mechanics, this will be a review. For others, there will be no requirement to follow precisely how the energy levels are derived--simply learning the final results that derive from quantum mechanics will inform our progress moving forward. Homework problems will provide you the opportunity to demonstrate mastery in the application of the above concepts.

- Dr. Christopher J. CramerDistinguished McKnight and University Teaching Professor of Chemistry and Chemical Physics

Chemistry

Now that we know everything there is to know about the allowed energy levels for

atoms, let's move on to molecules and start with the simplest kinds of

molecules, diatomic molecules. So just as we thought about how energy is

stored in an atom, let's think about how you can store energy in a molecule.

Exactly as was the case in an atom, we can still have electronic energy.

So this is kinetic and potential energy associated with one or more electrons,

usually more than one in a molecule. and it's their attraction to the two

nuclei in a diatomic, as well as their repulsion from one another that dictates

their potential energy, and then the kinetic energy of the electrons moving

about the nuclei. Similarly, there is a translational

energy that the molecule, as a whole, this diatomic molecule, is moving in

space. And there's kinetic energy associated

with that. We get this, both of these energies, in

the same way as we do for atoms. So there is no exact solution of the

Schrodinger equation for a many electron diatomic, just as there wasn't for a many

electron atom. But we can, do calculations or we can

look up energetics to get electronic energies.

The translational energy doesn't depend on whether something's an atom or a

molecule or a baseball for that matter. Once you know the mass and you know the

size of a box, you've got a particle-in-a-box and you can get the

solutions for the translational energy levels.

However, translational energy is kinetic energy.

There are two other kinds of kinetic energy available to a molecule that are

not available to an atom. The first is rotational energy.

So the entire molecule can rotate in space about various axes.

And the relevant Schrodinger equation that we need to solve in order to get the

allowed energy levels is called the rigid-rotator equation.

In addition, if I have two atoms connected by a bond, their motion

relative to one another, a vibration is a place where energy can be stored.

And the relevant Schrodinger equation to understand the energy levels available

for vibration is called the Quantum-mechanical harmonic oscillator

equation. So each of those we'll take a look at to

see what their, what the available energy levels are.

Let's start with rotational energy levels.

So here's a little diagram showing a diatomic molecule, which can be thought

of as two masses and they have some distance between them.

There is a center of mass, which need not be the midpoint if they are not equal

masses. And so we can define certain constants

within the system. We can find a center of mass.

We could define the distance of each of our masses, or I'll call them atoms

because we're doing chemistry, from that center of mass.

The sum of those two distances is the bond distance.

And from those data, we can compute the moment of inertia.

So moment of inertia is the, the equation is simply the mass times the square of

the distance of that mass from the center of mass.

And so that has appropriate units of mass times distance squared.

And when we solve the Schrodinger equation associated with rotational

motion, we discover that the allowed energy levels have associated with them a

quantum number, usually written as a capital J.

This is the first time we've seen a quantum number series start with zero.

Zero is an okay quantum number. Any integer in principal is okay.

Zero or one is typically where quantum numbers start from.

And the energy levels are given by h bar squared.

So h, remember is Planck's constant, and h bar means divided by 2 pi.

So we see that quantization constant creeping in.

The moment of inertia appears in the denominator of the energy value.

So the larger and larger the, the moment of inertia, the smaller and smaller the

energy levels. And they are spaced by J times J plus 1.

So if we just think about J equals zero for a moment, we would multiply zero

times 1 is zero, times a bunch of constants is zero.

So the ground state energy shown here is zero.

And if you like, that's not rotating. So there is no energy associated with

rotation. And it's, if you're not rotating.

So that's why zero is a quantum number there.

The next level up, when J is equal to 1 would give 1 times 1 plus 1 is 2.

So 1 times 2 is 2. That cancels this two in the denominator.

I'll get h bar squared over I. And so the next level up, shown here for

J equals 1, has that energy. And then I'll get 2 times 3 is 6 divided

by 2. I get a factor of 3 and then a factor of

6. I'll let you do the simple algebra if you

like. so we see this spacing, not unlike in the

hydrogen atom. The spacing gets sorry, exactly unlike in

the hydrogen atom, the spacing gets bigger and bigger as you go up.

Remember, in the hydrogen atom for electronic energy, the spacing got

smaller and smaller. So that's an interesting difference.

And the degeneracy of the levels, g sub J, is given by 2J plus 1.

So the ground state where J is 0 has no degeneracy.

If you like, there's only one way not to rotate.

when you go to the first level, it has degeneracy 2 times 1 is 2 plus 1 is 3.

And as a handy pneumonic to sort of remember this, a handy way to think about

that is if you think of there being three axes in space, a molecule can sort of

rotate about each of those axes, and that would give rise to a degeneracy of three.

That doesn't actually explain the degeneracy here, but it's a great thing

to keep in your mind to remember this rule for degeneracies, because you'll

immediately come up with 2J plus 1. If you remember, the ground state is, is

one, and the first state is three. So let's take a moment then, and you can

play a bit with computation of a moment of inertia.

[SOUND] Now let's move on, to look at the last place where energy can be stored in

a diatomic molecule, and that is, in the vibrations.

So, we'll look at the vibrational energy levels.

The vibration is associated with the two atoms moving in and out relative to one

another's positions. And one can think of the physical system

that decribes that as being two masses connected by a spring.

So this is a nice classic physical system.

the masses correspond to the atomic masses.

And the distance between them, when the spring is at its equilibrium length, that

is it's neither stretched so that there's a restoring force, nor is it compressed

so there's a different restoring force. It's just at rest.

That defines a distance R called the equilibrium distance.

So if we solve the quantum mechanical Schrodinger equation, then what we

discover is that the allowed energy levels for the harmonic oscillator are

given by this expression here, E the energy level indexed by a quantum number

v. So v for vibration, if you like.

And the quantum number, just like in rotation, it starts at 0.

Goes 0, 1, 2, and so on. The energy is v plus a half times

Planck's constant, there's Planck's constant again, times nu, the frequency.

That is, the frequency of time associated with a vibration.

So this equation, it's not my fault that physicists over many, many years decided,

and chemists, too, to use v and nu in the same equation, because it takes a really

good font to distinguish those in a substantial way, and you just have to

know from context which one is which. V is the quantum number.

Nu is the frequency. So, we'll try to be very careful and

always keep track, but it's just something to be aware of.

So if I go and look at these energy levels, and, what's shown here is the

potential energy curve, it's a parabola. And that's what the potential energy

looks like for a spring. As you stretch a spring, or as you

compress a spring, the energy goes up as the square of the stretch or the

compression. And how quickly it does so is given by

the force constant of the spring. In a molecular system, that'd be like the

strength of the bond. But in any case and that strength

dictates the frequency. So the stronger the bond, the higher the

frequency of the vibration. The really weak bond, big amplitude, slow

vibration, low frequency. But irrespective of that, the spacing

between the levels, because it's linear in the quantum number v, is constant.

So this is yet another way that levels can be spaced, that we haven't seen yet.

Constant spacing between all the levels. So here's the ground state, the first

excited state, the second, and so on. So, the degeneracy of all these levels is

one. There are no degeneracies for the quantum

mechanical harmonic oscillator. The spacing in between them, the

variation, if I start at 0 and I get a half, and I go to 1 and so I'll have 3

halves, the difference between 3 half and 1 half is 1.

And so the spacing between all of them is h nu.

So Planck's constant times the frequency of the vibration.

Notice that the ground state energy is above the bottom of the potential energy

curve. So that's an interesting phenomenon.

That phenomenon is known as zero-point energy.

So it's energy that the system is storing, even in the ground state.

Were I to take the atoms completely apart, I could recapture that energy in

principle. But they are they're not allowed to go

down to the bottom. And that's associated with certain

features of the quantum mechanical nature of this particular problem.

So knowing what we know about the quantum mechanical harmonic oscillator, let's

move from the, the model system of two masses on a spring and talk about a real

chemical bond. So a bond does not go up a potential

curve indefinitely. Instead, the atoms cease to interact with

one another at a certain point and the bond breaks.

So a more accurate curve for chemical bonding is shown here.

When I bring, so this axis is distance. When I bring two atoms very close

together, they repel each other strongly and the potential goes up, up, up, up,

up. There is some equilibrium distance where

the potential is at its lowest. And there will be a ground state

vibrational energy level with a average distance somewhere near that lowest

energy distance. And then as I pull the atoms apart, they

go up for a while, looking like a parabola, but ultimately they just stop

interacting. And that defines the zero of energy.

They don't interact with each other anymore.

And so, two important quantities on this curve.

DE, so that is the separation between the bottom of the potential and zero.

And D0, the separation between the ground state vibrational level, and zero.

Those are different, and they differ by the zero-point energy, which is one half

h nu. So the relationship I, I just described

here, and I'll just show it in an equation.

Namely, the negative of the ground state electronic energy.

So here is an electronic energy, De, defined for a diatomic, by calling zero

separated. And De the bottom of the potential is

equal to D0, so that's dropping down to the ground-state vibrational level plus

the zero-point vibration. So this is referred to as the

dissociation energy. It's how much energy we really have to

provide to rip the, the diatomic molecule apart.

But the electronic energy has this additional h nu over 2 component.

And so just as an example, hydrogen gas has a, a bond dissociation energy of 432

kilojoules per mole. If we measure the frequency of vibration

of hydrogen with an infra-red spectrum, for instance, it's 4, 401 wave numbers.

And if you plug that in to the relevant expression that we saw some time ago, to

relate wave numbers to energy, that would be 52 kilojoules per mole.

One half of that would be 26. And the difference between these two, the

electronic energy and the bond association energy, is that 26.

All right, so that completes the necessary quantum mechanical background

to understand the energy levels of diatomic molecules.

All that's really left now that we've done atoms and diatomic molecules, is

polyatomic molecules. So that'll be the subject of the next

video, and then we'll have wrapped up the necessary underpinnings of quantum

mechanics to advance on to considering thermodynamics.

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