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## UV/Vis Spectroscopy

Current time:0:00Total duration:11:12

# UV/Vis spectroscopy

## Video transcript

- [Voiceover] Different
molecules can absorb different wavelengths of light and
if a molecule happens to absorb light in the
ultraviolet or the visible region of the electromagnetic
spectrum we can find the wavelength or
wavelengths of light that are absorbed by that compound by using a UV/Vis spectrophotometer. Now essentially what
that does is it shines light with a range of wavelengths. The wavelengths range
from approximately 200 nanometers all the way
up to 800 nanometers. We shine that range of wavelengths
of light through a sample of the compound and you
get an absorption spectrum. Here is an absorption spectrum for this molecule, for 1,3-Butadiene. Now if we look over here
we can see that this molecule absorbs most
strongly right about here and if we drop down we
can see what wavelength of light is absorbed most
strongly by the compound. And we see that's just
under 220 nanometers. It turns out out to be 217 nanometers. We call this lambda max. The wavelength of light absorbed by this molecule is about 217 nanometers. It absorbs in the UV
region therefore Butadiene does not have any color, it's colorless. Let's look at the dot structure a little bit more carefully here. We have four carbons and all four of these carbons, each one is sp2 hybridized. Which means each one of those
carbons has a p orbital. So we're talking about four p orbitals here or four atomic orbitals. And when you're dealing with molecular orbital theory, four atomic orbitals recombine to form four molecular orbitals. Two bonding molecular orbitals and two antibonding molecular orbitals. Let's go over here and let's
look at the four molecular orbitals and we're going to
focus in on the left side first. The bonding molecular orbitals are lower in energy than the antibonding ones. So this orbital and this
orbital, these are our bonding molecular orbitals here
and this one and this one are the antibonding molecular orbitals. And you can see energy, right? So energy is increasing
and so the antibonding molecular orbitals are higher in energy. Let's look at the dot
structure again for Butadiene and let's see how many
pi electrons we have. So here are two pi electrons
and here are two pi electrons. So a total of four pi electrons. When you're thinking
about molecular orbitals, you can think about
electron configurations. So we have four electrons and where do we put those electrons? We're going to put them in the lowest energy orbitals first. And we're also going to pair our spins. So four electrons, we're
going to put two into this bonding molecular orbital
and we paired our spins. And then two into this
bonding molecular orbital. So the four pi electrons go into the bonding molecular orbitals when you're talking about the ground state. So here's the ground state of Butadiene. So next we shine light
on Butadiene and the molecule's going to absorb
energy from the light. Let's look at that here,
so there's a difference in energy between the orbitals
and in particular we're concerned about these two
orbitals right here so there's a difference in energy
between these two orbitals. This orbital down here, this
is occupied by electrons and it's higher in
energy than this orbital. So this is the highest
occupied molecular orbital. So highest occupied
molecular orbital or HOMO. This orbital right here is unoccupied. The antibonding molecular
orbital right now is unoccupied and it's
lower in energy than this antibonding molecular orbital. So this is the lowest
unoccupied molecular orbital. When you're talking about a molecule absorbing energy, we're
considered about the HOMO, the highest
occupied molecular orbital and the LUMO, the lowest
unoccupied molecular orbital. The energy difference between those two orbitals is what we're thinking about. So the molecule absorbs energy and a pi electron absorbs energy from the light and is promoted to a higher energy level. Let me go ahead and write over here. Now we're talking about the excited state so we shine light on the molecule. This is the excited state of Butadiene and these two pi electrons stay there. One of these pi electrons
stays here and one of the pi electrons absorbs
the energy from the light and is promoted to a higher energy level. So I'm saying this one right here was promoted to a higher energy level. It goes from the HOMO to the LUMO and it had to absorb a specific amount of energy in order to do that. So it had to absorb the right amount of energy in order to
make that transition. We know that energy came from the light and we also know the energy of a photon of light is equal to h,
where h is Planck's constant, times the frequency of light which is new. Over here for the absorption spectrum, we have everything in
wavelengths so we need to write the energy in
terms of a wavelength. We know that the frequency of light and the wavelength of light are related by the speed of light is equal to the wavelength times the frequency. The frequency is equal
to the speed of light over the wavelength and we can take that, frequency is equal to c over
lambda, and plug it into here. Now we have the energy, the energy is equal to h times c over lambda. This is really important;
energy and wavelength are inversely proportional to each other. You can think about one wavelength giving you a specific amount of energy. This energy difference
between the HOMO and the LUMO corresponds to a wavelength
and if we go over here to the absorption spectrum
for Butadiene we're talking about a wavelength
of 217 nanometers. At first it might be
a little bit confusing because it looks like we
have a very broad range of wavelengths that are absorbed here. Don't worry about that too
much, this just results from the different vibrations and
rotations of the molecule which can change the
energy differences slightly and so we don't see one exact wavelength, we end up seeing this broad band of wavelengths being absorbed here. So what you do is, you just
look for the one that's absorbed most strongly
and think about that as being the wavelength
that corresponds to the energy difference between
these two orbitals here. So that's how to think about it. Let's look at another molecule here, instead of Butadiene let's look at this molecule, so we have ethanal. Here is our dot structure and if we look at this molecule we know we have two pi electrons here for ethanal. So two pi electrons. We know that those
electrons are going to go into the bonding molecular orbital. So let me draw a line
right here on this diagram. This is our bonding
molecular orbital down here. We're talking about two pi electrons. Let's put in our two
pi electrons into here. Let me just go ahead and
change colors up here. Up here is our antibonding molecular orbital which we call pi star. So there is an energy difference between the bonding molecular orbital and the antibonding molecular orbital. This is delta E and we talked about the fact that this corresponds to a certain wavelength of light. Ethanal can have, when
it promotes one of these pi electrons up, it can have
a pi to pi star transition. So the molecule is going to
absorb energy and the energy-- Let me use a different color here. The energy corresponds
to a wavelength of light so this energy difference
between our two orbitals. It turns out that this
pi to pi star transition is approximately 180
nanometers which is below the range of what you're usually measuring when you're using a
UV/Vis spectrophotometer. But we have another possibility here too. Let me go ahead and highlight a lone pair of electrons here on the oxygen. We have a lone pair so we
have non-bonding electrons. Non-bonding electrons occupy a non-bonding orbital which is actually
a little bit higher in energy than our
bonding molecular orbital. So another possibility,
we call this n right here. This is a non-bonding orbital
so non-bonding orbital here. And we can put some
electrons into that orbital. So we put those two electrons
into the non-bonding orbital. And we can have a different
type of transition. We're still talking about a pi star, an antibonding molecular
orbital right here. We can have a n to pi star transition. We can have a n to pi
star transition as well since we have a carbonyl compound. We're not just talking
about pi electrons here. We can think about a
non-bonding electron here. And let's think about
this energy difference. This energy difference
is smaller than before. So this energy difference is smaller than this energy difference. What would happen to the wavelength of light that's absorbed? If we have a smaller energy difference, energy and wavelength are
inversely proportional so this must be a longer wavelength. So this absorbs light at
a different wavelength, a higher wavelength,
and it turns out to be-- Let me go ahead and change colors here. So this energy transition
corresponds to a wavelength of light that's
approximately 290 nanometers. This n to pi star transition,
a smaller difference in energy corresponding
to a higher wavelength. This is an important concept. As you decrease the
energy difference between your orbitals, you're
going to increase the wavelength of light that's absorbed. We'll talk much more about
that in the next few videos because that's where the
idea of color comes in.