UV/Vis spectroscopy | Spectroscopy | Organic chemistry | Khan Academy

Summary

This script explores the principles of UV/Vis spectrophotometry, demonstrating how molecules like 1,3-Butadiene and ethanal absorb different wavelengths of light. It explains the concept of molecular orbitals, including bonding and antibonding orbitals, and how the energy difference between the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO) determines the absorbed wavelength. The script also discusses electron transitions, such as pi to pi* and n to pi*, and how they relate to a molecule's color and light absorption characteristics.

Takeaways

• 🌈 Molecules can absorb different wavelengths of light, and this can be measured using a UV/Vis spectrophotometer.
• 🔍 The spectrophotometer shines light across a wide range of wavelengths (200-800 nm) through a compound to produce an absorption spectrum.
• 🧪 For 1,3-Butadiene, the most strongly absorbed wavelength is just under 220 nm, specifically at 217 nm, which is referred to as lambda max.
• 🏞️ Butadiene is colorless because it absorbs in the UV region, which is not visible to the human eye.
• 🔬 The molecule's structure, specifically the sp2 hybridization of carbons leading to p orbitals, is crucial for understanding its molecular orbitals.
• ⚛️ Butadiene has four pi electrons that occupy the bonding molecular orbitals in its ground state.
• 🌟 When light is shone on Butadiene, a pi electron can absorb energy and be promoted from the HOMO (highest occupied molecular orbital) to the LUMO (lowest unoccupied molecular orbital).
• ⚡ The energy absorbed by the molecule corresponds to a specific wavelength of light, which can be calculated using Planck's constant and the speed of light.
• 🌟 The energy difference between molecular orbitals is inversely proportional to the wavelength of light absorbed, leading to different colors for different molecules.
• 🔬 Ethanal, another molecule discussed, has two pi electrons and can undergo both pi to pi* and n to pi* transitions, with the latter corresponding to a longer wavelength of light due to a smaller energy difference.

Q & A

What is the purpose of a UV/Vis spectrophotometer?

A UV/Vis spectrophotometer is used to determine the wavelengths of light absorbed by a compound in the ultraviolet or visible region of the electromagnetic spectrum by shining light with a range of wavelengths through a sample of the compound and analyzing the resulting absorption spectrum.

What is the wavelength range of light that a UV/Vis spectrophotometer typically covers?

The wavelength range of light covered by a UV/Vis spectrophotometer typically extends from approximately 200 nanometers to 800 nanometers.

What is the term for the wavelength of light absorbed most strongly by a molecule?

The wavelength of light absorbed most strongly by a molecule is referred to as lambda max (λmax).

Why is 1,3-Butadiene colorless?

1,3-Butadiene is colorless because it absorbs light in the UV region, specifically at a wavelength of 217 nanometers, which is not visible to the human eye.

What type of hybridization do the carbons in 1,3-Butadiene have, and what does this imply for their orbitals?

The carbons in 1,3-Butadiene are sp2 hybridized, which means each carbon has a p orbital. This results in four atomic orbitals that combine to form four molecular orbitals, including two bonding and two antibonding orbitals.

How many pi electrons are present in 1,3-Butadiene, and where are they located in the molecular orbitals?

1,3-Butadiene has four pi electrons, which in the ground state are located in the two bonding molecular orbitals, with two electrons in each.

What are the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO), and why are they significant in the context of light absorption?

The HOMO is the highest energy orbital that is occupied by electrons, while the LUMO is the lowest energy orbital that is unoccupied. They are significant in light absorption because the energy difference between these two orbitals determines the specific wavelength of light that a molecule can absorb, leading to an electron transition from HOMO to LUMO.

How does the energy of a photon of light relate to its wavelength?

The energy of a photon of light is inversely proportional to its wavelength. This relationship is described by the equation E = h * c / λ, where E is the energy, h is Planck's constant, c is the speed of light, and λ is the wavelength.

What is the significance of the energy difference between the HOMO and LUMO in terms of the molecule's absorption spectrum?

The energy difference between the HOMO and LUMO corresponds to a specific wavelength of light that the molecule can absorb. This is observed as a peak in the absorption spectrum, which can be used to identify the molecule.

What type of electron transition occurs in ethanal, and what is its approximate wavelength?

In ethanal, a pi to pi star transition occurs, which has an approximate wavelength of 180 nanometers. Additionally, an n to pi star transition is possible, corresponding to a longer wavelength of approximately 290 nanometers.

How does the energy difference between molecular orbitals affect the wavelength of light absorbed by a molecule?

A smaller energy difference between molecular orbitals corresponds to a longer wavelength of light absorbed. As the energy difference decreases, the wavelength of absorbed light increases, which is related to the color that the molecule can exhibit.

Outlines

00:00

🌈 UV/Vis Spectrophotometry and Molecular Orbitals

This paragraph introduces the concept of UV/Vis spectrophotometry, a technique used to determine the wavelengths of light absorbed by a molecule in the ultraviolet and visible spectrum. The absorption spectrum of 1,3-Butadiene is discussed, highlighting its strongest absorption at 217 nanometers, which indicates the presence of a lambda max in the UV region. The paragraph also delves into the molecular structure of 1,3-Butadiene, explaining its sp2 hybridization and the formation of molecular orbitals from p orbitals. The electron configuration in the ground state is described, with all four pi electrons occupying the bonding molecular orbitals. The process of light absorption leading to an excited state is also explained, emphasizing the transition of a pi electron from the highest occupied molecular orbital (HOMO) to the lowest unoccupied molecular orbital (LUMO).

05:03

🔬 Energy Transitions and the Absorption of Light

The second paragraph focuses on the energy transitions within molecules and how they relate to the absorption of light. It explains the quantification of energy absorbed by a molecule using Planck's constant and the frequency of light, and then relates this to the wavelength by using the speed of light. The inverse relationship between energy and wavelength is highlighted, with a specific example of how the absorption spectrum for Butadiene shows a broad range of wavelengths due to molecular vibrations and rotations. The paragraph also introduces the concept of different types of electronic transitions, such as pi to pi star, in the context of another molecule, ethanal, and how these transitions correspond to different wavelengths of light.

10:04

🌟 Impact of Energy Differences on Absorbed Wavelengths

The final paragraph explores how the energy difference between molecular orbitals affects the wavelength of light absorbed by a molecule. It contrasts the energy differences and corresponding wavelengths for two types of transitions in ethanal: the pi to pi star transition and the n to pi star transition. The paragraph explains that a smaller energy difference results in the absorption of light at longer wavelengths, using the example of the n to pi star transition in ethanal, which absorbs light at approximately 290 nanometers. The concept is linked to the idea of color, suggesting that as the energy difference decreases, the absorbed wavelength increases, which will be further discussed in subsequent videos.

Keywords

💡UV/Vis spectrophotometer

A UV/Vis spectrophotometer is an analytical instrument used to measure the absorption and transmittance of light by a sample in the ultraviolet and visible light regions of the electromagnetic spectrum. In the video, it is used to identify the wavelengths of light absorbed by a compound, which helps in determining the compound's structure and properties. The script describes how shining light through a sample of 1,3-Butadiene results in an absorption spectrum, with the most strongly absorbed wavelength being identified as 217 nanometers.

💡Absorption spectrum

An absorption spectrum is a graphical representation that shows how much light at different wavelengths is absorbed by a substance. It is fundamental in spectroscopy and is used to identify the presence of specific molecules. In the script, the absorption spectrum of 1,3-Butadiene is discussed, highlighting the peak absorption at around 220 nanometers, which indicates the compound's lambda max.

💡Lambda max (λmax)

Lambda max, often denoted as λmax, is the wavelength of light that a molecule absorbs most strongly. It is a characteristic property of a molecule and can be used to identify it. In the context of the video, the λmax for 1,3-Butadiene is identified as 217 nanometers, which is in the UV region, contributing to its colorless nature.

💡sp2 hybridization

sp2 hybridization is a type of atomic orbital hybridization that occurs in molecules with three regions of electron density around the central atom, such as in trigonal planar molecules. In the script, it is mentioned that all four carbons in 1,3-Butadiene are sp2 hybridized, meaning each carbon has one unhybridized p orbital available for pi bonding, which is crucial for the molecule's electronic structure and its absorption spectrum.

💡Molecular orbitals

Molecular orbitals are the atomic orbitals combined to form new hybrid orbitals that result from the interaction of atomic orbitals in a molecule. They can be bonding or antibonding. The video script explains that four atomic orbitals of the carbons in 1,3-Butadiene recombine to form four molecular orbitals, two of which are bonding and two are antibonding, which are essential in understanding the molecule's electronic transitions.

💡pi electrons

Pi electrons are a type of electron found in pi bonds, which are formed from the overlap of parallel p orbitals in molecules. The script refers to the four pi electrons in 1,3-Butadiene, which occupy the bonding molecular orbitals in the ground state and are involved in the electronic transitions when the molecule absorbs light.

💡Highest Occupied Molecular Orbital (HOMO)

HOMO stands for the Highest Occupied Molecular Orbital, which is the highest energy orbital that contains electrons in the ground state of a molecule. In the video, the HOMO of 1,3-Butadiene is the orbital from which a pi electron is promoted to a higher energy level upon light absorption.

💡Lowest Unoccupied Molecular Orbital (LUMO)

LUMO refers to the Lowest Unoccupied Molecular Orbital, which is the lowest energy orbital that does not contain electrons in the ground state. The video script explains that when a molecule absorbs energy from light, an electron from the HOMO is promoted to the LUMO, which is a key concept in understanding electronic transitions and the molecule's absorption spectrum.

💡Excited state

The excited state of a molecule is a higher energy state achieved when one or more electrons are promoted to higher energy orbitals, typically by absorbing a photon of light. In the script, the excited state of Butadiene is described as the state where a pi electron has absorbed energy and moved from the HOMO to the LUMO.

💡Planck's constant (h)

Planck's constant is a fundamental physical constant that relates the energy of a photon to its frequency. In the video, it is used in the formula for the energy of a photon, which is crucial for understanding how the energy from absorbed light is related to the electronic transitions in molecules.

💡Energy difference and wavelength

The energy difference between molecular orbitals determines the wavelength of light a molecule can absorb. As explained in the script, a smaller energy difference corresponds to a longer wavelength of light absorbed. This concept is central to understanding how different molecular transitions, such as pi to pi* and n to pi* in ethanal, result in the absorption of light at different wavelengths.

Highlights

Different molecules absorb different wavelengths of light, particularly in the ultraviolet or visible regions.

A UV/Vis spectrophotometer measures the wavelengths of light absorbed by a compound, typically ranging from 200 to 800 nanometers.

1,3-Butadiene absorbs light most strongly at approximately 217 nanometers, indicating it is colorless and absorbs in the UV region.

1,3-Butadiene has four sp2 hybridized carbons, each contributing a p orbital, forming four molecular orbitals.

In molecular orbital theory, four atomic orbitals recombine to form four molecular orbitals: two bonding and two antibonding.

Bonding molecular orbitals are lower in energy compared to antibonding molecular orbitals.

The highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO) are key in understanding energy absorption.

The energy difference between the HOMO and LUMO is crucial for determining the wavelength of light absorbed by a molecule.

Energy and wavelength are inversely proportional; as the energy difference between orbitals decreases, the absorbed wavelength increases.

In Butadiene, upon absorbing light, a pi electron is promoted from the HOMO to the LUMO, resulting in an excited state.

The absorbed energy is determined by Planck's constant and the frequency or wavelength of light.

Ethanal, another molecule discussed, has two pi electrons that can transition from a bonding to an antibonding molecular orbital.

Ethanal can undergo a pi to pi star transition, absorbing light at around 180 nanometers.

A non-bonding to pi star (n to pi*) transition is also possible in carbonyl compounds like ethanal, occurring at a longer wavelength of around 290 nanometers.

Understanding the relationship between energy differences and absorbed wavelengths is essential in the study of molecular color.

Transcripts

- [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.