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

Video transcript

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 essentially what that does is it shines light with a range of wavelengths so the wavelengths range from approximately 200 nanometers all the way up to 800 nanometers and so we shine that range of wavelengths of light through a sample of the compound and you get an absorption spectrum and so here is an absorption spectrum for this molecule for 1 3 butadiene if you 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 so we see that's just under two hundred and twenty nanometers and it turns out to be 217 nanometers so we call this lambda max all right so wavelength 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 alright let's look at the dot structure a little bit more carefully here so we have four carbons and all four of these carbons each one is sp2 hybridized which means each one of those carbons has ap 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 to bonding molecular orbitals and to anti-bonding molecular orbitals so let's go over here and let's look at let's look at the four molecular orbitals and we're going to focus in on the left side first the bonding molecular bonding molecular orbitals are lower in energy than the anti-bonding ones so this orbital and this orbital these are our our bonding molecular orbitals here and this one and this one are the anti bonding molecular orbitals and you can see energy right so energy is increasing and so the anti bonding 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 right so here are two pi electrons and here are two pi electrons so a total of four pi electrons and when you're thinking about molecular orbitals you can think about electron configurations right 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 right 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 orbitals 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 the ground state of butadiene so next we shine we shine light on butadiene and the molecule is going to absorb energy from the light and let's look at that here so there's a difference in energy right 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 right 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 right is unoccupied the anti-bonding molecular orbital right now is unoccupied and it's lower in energy than this anti-bonding molecular orbital so this is the lowest unoccupied molecular orbital so when you're talking about when you're talking about a molecule absorbing energy we're concerned 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 all right so let's 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 goes from the homo to the LUMO and it had to absorb a specific amount of energy in order to do that right so it had to absorb the right amount of energy in order to make that transition and we know 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 nu and over here for our for the absorption spectrum right we have everything in wavelengths so we need to write the energy in terms of wave in terms of a wavelength and 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 and so the frequency is equal to the speed of light over the wavelength and we can take that so frequency is equal to C over lambda and plug it into here so now we have the energy the energy is equal to H times C over lambda and this is really important right so energy and wavelength are inversely proportional to each other and you can think about one wavelength right giving you a specific amount of energy and so this energy difference right this energy difference between the homo and the LUMO right this 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 and so at first it might be a little bit confusing because it looks like we have a very broad it looks like we have a broad range of wavelengths that are absorbed here and 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 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 between these two orbitals here so that's how to think about it all right let's uh let's look at another molecule here so instead of butadiene let's look at this molecule so we have ethanol all right so here's our dot structure and if we look at this molecule alright we know we have we have two pi electrons here for ethanol so two pi electrons we know that those electrons are going to go into the bonding molecular orbital so let me draw a line let me draw a line right here on this diagram so this is our this is our bonding our bonding molecular orbital down here and this is our so we're talking about two pi electrons so let's put in our two pi electrons into here and then let me just go ahead and change colors up here so this appears our anti-bonding molecular orbital which we call pi star alright so there's an energy difference between the bonding molecular orbital and the anti-bonding molecular orbital all right so this is uh this is Delta e and we talked about the fact that this corresponds to a certain wavelength of light and so ethanol ethanol can have when it promotes one of these pi electrons up right it can have a PI we call this a pi to pi star transition right so the molecule is going to absorb energy right and the energy the energy will use a different color here the energy corresponds to a wavelength of light so this energy difference between our two orbitals and it turns out that this pi to pi star transition is approximately a hundred 80 nanometers which is is below the range of what you're usually measuring right when you're doing when using a uv-vis spectrophotometer all right but we have another possibility here - let me go ahead and highlight a lone pair of electrons here on the oxygen alright so we have a lone pair right so we have non bonding non bonding electrons all right and nonbonding electrons occupy a nonbonding orbital which is actually a little bit higher in energy than our bonding molecular orbitals so another possibility we call this n right here so this is a non bonding orbital so non bonding orbital here and we can put some electrons into that orbital all right so we put that put those two electrons into the nonbonding orbital and we can have a different type of transition so we're still talking about we're still talking about a PI star alright an anti-bonding molecular orbital right here we can have a n to pi star transition right so we can have an n to pi star transition as well since we have a carbonyl compound so we're not just talking about pi electrons here right the carbonyl we can think about a nonbonding electron here and let's think about this energy difference so this energy difference is smaller than before alright so this energy difference is smaller than this energy difference and so what would happen to the wavelength of light that's absorbed all right so if we if we have a smaller energy difference energy and wavelength are inversely proportional so this must be a longer wavelength all right so this absorbs light at a different wavelength right so 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 and so this this n to PI star transition a lower difference at a smaller difference in energy I should say corresponding to a higher wavelength this is an important concept so as you decrease the energy difference between between your orbitals you're going to increase the wavelength of light that's absorbed and we'll talk much more about that in the next few videos is that that's where the idea of color comes in