Lecture 27: Organic Chemistry

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OK, let's get started. One announcement: test three. It's a celebration of learning. It's going to be on Wednesday, which means no lecture. No lecture Wednesday. Instead, it's a celebration. Please go to your rooms as assigned here.

If you are writing at 11 o'clock, probably in this room right now, and those are the room assignments. They are the same as for test two, but not the same as for test one. We moved more people out of 10-250 to give a little more room here.

And, there's plenty of room in these others. Those that write at one o'clock do not go to the normal lecture room 6-120. We're writing at 26-100. And, the coverage is right here. So, that's up on the website.

So, you should know what we are going to be examining you on, and the same comments that I made before: I want to give you feedback, let you know how you're doing, whether your study methods are effective or not.

It's not an attempt to retest your admission to MIT, or anything like that. If you do your work, you should do very well. And if you have not done your work, you shouldn't do very well, and we'll be able to tell you so.

So, just what we said in the past. Please take the time to read the whole exam, do the easy questions for you; the ones that you find easiest, do those first. But let's be honest, this isn't high school anymore.

So, I mean, I had people tell me after the second test, I had my aid sheet, but I hardly used it. Well, think about it. Do you think I'm going to give you an aid sheet and then give you a question that requires you to take something off the aid sheet and transfer it to the answer paper? It's not pattern recognition.

I mean, this isn't medical school, for crying out loud. You've got to think here. You got to think. So, this is now the third test. There's going to be more and more thinking, and less and less rote, although we'll have some confidence builders on there.

I don't want to knock people off their balance. So, work parametrically. Try not to immediately start punching in numbers. And, if you don't know how to do a question in great detail, write me an outline.

Tell me what you would do. Either you equate energies, or minimize something, or maximize something. Give us a sense that you have a grasp of the material. And, keep your eyes on your own paper. Our effort is going to be to get those graded on Wednesday, and back to you in recitation on Thursday.

So, I think that's, and remember, please bring your five items. We are having more and more people showing up minus periodic table, table of constants, something to write with; I've had people ask me for pens, for calculators, no one has asked me yet for an aid sheet.

So, I haven't had to help anybody on that. OK? Good. Well, today what I want to do is start a new unit. I want to talk today about organic chemistry. Now, this is going to be the most minimal introduction to organic chemistry.

The reason we talk about organic chemistry at all is in order to prepare ourselves for more solid-state chemistry, specifically, I'm going to talk about polymers. And ultimately, we're going to talk about biochemistry because even though this is solid-state chemistry, we as living beings are solid-state devices.

We are made of soft matter. This is a polymer. This is a semiconductor, band gap of two to three electron volts. But it's not made out of three-five semiconductor. Nature has figured out a way of doing this without inorganic means.

So, we need to know a little bit. But if you really want organic chemistry, you're going to have to take 5.12. So, I can't teach you 50 minutes what people teach in one semester. So, let's get a few definitions under our belts.

Organic chemistry is the chemistry of compounds containing both carbon and hydrogen, not just carbon, carbon and hydrogen. So, diamond and graphite are not organic because they contain carbon but not hydrogen.

And, what makes these two elements special that they figure so prominently in living organisms? First of all, hydrogen is the element with the lowest atomic number. And, when it loses its electron, it changes character dramatically.

We talked about this in recent lectures where we have simply a proton. It's capable of forming covalent bonds, but when it finds itself in a compound forming a covalent bond with something that's strongly electronegative, it's so denuded of electrons even in the covalent bond that that protonic entity can start mischief and form hydrogen bonds.

Carbon is dead center: four valence electrons. So, it's nowhere near metallic. It's nowhere near nonmetallic. So, it's got an intermediate value of average valence electron energy. It's very unlikely that carbon is going to acquire four electrons and become C4 minus, or lose four electrons and become C4 plus.

It's a prominently covalent player, but it's very small. And, because it's small, it can form multiple bonds. It forms multiple bonds. And, we've already seen evidence where it forms sigma and pi bonds.

So, it can form double bonds and triple bonds not only with itself but with other elements, notably nitrogen, oxygen, phosphorus, and sulfur. And even though silicon and germanium hybridize, because they are not small, silicon does not form double bonds.

You will not see silicon analogues to carbon compounds throughout. Yeah, silicon sp3 hybridizes, and single crystal silicon has the same crystal structure as single crystal diamond, but that's pretty much where the similarity ends.

It's pretty much sp3 hybridization, and that's the end of the story. So, these two have very special qualities that make them so important. And, I mean, the other thing that carbon can do is it can continue to link with itself and form chains and various other structures.

And, we're going to meet a few of those today. There are millions of organic compounds; I'd say billions, billions, and billions. But we are not going to cover all that. We are going to cover the ones that are most germane to our subsequent discussions of polymers and biochemistry.

So, let's first of all go through the taxonomy. So, I mean, I've posted all this on the website. And it's in the reading. I have tried to characterize this in a way that, or categorize it so that it makes some simple sense as opposed to just standing here and barking out a litany of facts to you.

So, the first thing I want to do is look at the left-hand column called the alkanes. And, these are hydrocarbons. The whole thing we're going to look at is hydrocarbons for starters. And, these, as the name implies, are compounds that consist of only carbon and hydrogen.

And the first group that we're going to look at is the alkanes, which you can see in the slide. And, these are characterized by sp3 hybridization on the carbon. All the carbons are sp3 hybridized. And so, that gives us the maximum possible number of bonds, maximum number of carbon-hydrogen linkages.

And so, the chemical term for such compounds is saturated. These are saturated hydrocarbons. And, all of these bonds are sigma bonds. All the bonds are sigma bonds because they're all single bonds.

So, these are the characteristics. In a general formula, if you go through the math, it's going to be C some sub n number, whatever the number of carbons in the molecule, and then if we go through sp3 hybridization and put hydrogens at all the noncarbon linkages, it'll be H two times the C number plus two.

And so, these are called alkanes. And, this is the way you build the name. The nomenclature is to, it's going to have an -ane ending, which is going to indicate that we're talking about sp3 hybridized hydrocarbons.

And then, in order to indicate the number of carbons, we call out the carbon number by the prefix. So, this prefix identifies the carbon number, and they are shown here actually. Let's take a look at a set of those.

This is right out of your reading. And there they are: one, two, three, and so on. And, it's basically meth-, eth-, prop- is three, but- is four, and these are all historical. And, you can go into the reading and figure out butyric acid is one of the elements in rancid butter and so on.

So, there is some historical reasons. But, after you get to four, it's strictly the Latin ordinals. From here on, this is pent, and what do we have? We have hex-, hept-, non-, dec-, and so on. So, just go back to your Latin, and you're going to be fine.

I'm not going to expect you to pull these out of memory. I would tell you C5, well, what, you've forgotten your Latin? Was it an admission requirement? How many semesters of Latin did you need in high school to get into MIT? None, OK, so I think I'd better keep that in mind.

So I will tell you pentane, and I'll give you the formula: C5H12. All right, so these are also called, they are called straight chain molecules because we you look at them a little bit more carefully, we will see the following.

Let's have a look at what they look like. Oh, before I go, I just wanted to hearken back and show that there is some overarching theme here. So, these are three hydrocarbons that I've chosen: propane, which we know is used as a fuel in, among other things, transportation and even domestic barbecue grills.

And, it's a gas at room temperature. These are all the same. They are long chains with hydrogen around them. OK, these are all symmetric molecules. Octane, which is the principal constituent of gasoline, is a liquid at room temperature.

And, icosane is a solid at room temperature. So, all of these have, they are symmetric molecules, and the only thing that's happening as you go to larger and larger molecules is your increasing polarizability, your van der Waals bonds between molecules of the same identity increases, and you can see that as the molecular length increases, the molecule converts from gas, to liquid, to solid.

So, it's something to keep in mind. So, these are the ball and stick images. Methane we've met before. There's the sp3 hybridization, four bonds, 109° apart. In the case of ethane, this is C2H6. One of the four bonds is a carbon-carbon bond.

It's a sigma bond. OK, and then we go to propane. So, now we've got C3. So, that's what's cueing us in on which of these to choose. But, I want you to note that at propane, this sp3 hybridization requires that as we add more and more carbons, if we're going to put one more carbon, we're going to put it up here.

You can see that even though this thing is straight, it, in fact, at the local level is zigzagging. At the local level, there is some zigzagging. So, we want to keep that in mind. The sp3 hybridization gives bonding along the chain that in fact zigzags.

And, this is a 109° angle. And, so, if you are up really close, you say, well, this is clearly bent. But if you make this long enough, and you get far enough back, in a general sense, this is termed a straight chain.

It's a straight chain hydrocarbon. There is one other thing that is similar to what we've seen in the past in the case of silicates. And, that's shown here. As long as we've got 109° along this axis between the other bonds, there's no specification what happens down here.

So, this should remind you of what happens in a silicate when the oxygens don't all line up. And so, it's possible to get disorder. And so, here you see one case where you have the hydrogens on adjacent carbons not facing one another, whereas here they are lined up.

So, if I were to look on end from the left down the chain, I would only see one, that all the hydrogens are lined up. And we call this configuration an eclipsed configuration. This is slightly high-energy, whereas this is a little bit lower energy, staggered.

And, what's the effect of that? The effect is not to give us a straight chain that goes in a beeline. For that, you would have to have the eclipsed configuration. Here's two examples of C17H36. They're both straight chain, but because one's twisting a little bit more than the other, and here I'm talking about just the carbon-carbon bond twisting.

These are both termed straight chain. But, they have different conformations. So, this is starting to make you think about the plurality of possibilities even within the same chemical composition that would be present in polymers.

Here, we've only got 17. Imagine if instead of 17 it became 17,000, what this can do. So, this is an example of what we mean by straight chain. But it doesn't mean that it's a rod-like entity. OK, but I can now show you something that's not a straight chain.

So, let's also look at something called a branched chain. Now, that's different. And, I'm just going to look at one. We're going to look at butane. That's a gas at room temperature. It used to be used as fuel for cigarette lighters.

But I guess that's not PC anymore. Now you can use it for lighting candles. I can say that, can't I? Let's look at butane. So that's number four. So, all right, so just to smooth things along, it's possible to just write them in a straight line even though we know these are 109°.

And, we know that if there's nothing put at the end of the stick, we assume it's hydrogen. So, all of these are hydrogen, carbon hydrogen-linkages. So, I've got one, two, three, four carbons. And, I've got one, two, three, four, five, six, seven, eight, nine, ten.

So, this is as it should be C4H10. And certainly this is identical to something that could be represented as follows. All right, so now I'm going to have the four outlining a tetrahedron. So, there is one, two, three, four, one, two, three, four, one, two, three, four.

Remember, when you're in carbon chemistry, four sticks off of every carbon: that's what you're checking. OK, so this is the linear. But there's another way. There's another way we can do this. We can do the following.

We can put - and off one of the carbons. I can put a branch, and then I've got to use the four stick rule, so, one, two, three, one, two, so, one, two, three, four. This has one, two, three, four.

So, now, let's count: one, two, three, four, four carbons, and three, six, nine, ten hydrogens. So, on a formula basis, this is also a butane. But, it's a different. As solid-state chemists, we are very much attuned to molecular structure.

And this molecular structure, the linear one is different from this one. This is a branched one. So, this one's called iso-butane. And, the other way to look at it is, well, the backbone is only three carbons.

So, if the backbone is only three carbons, according to this rule, this should be some kind of a propane. Another way to name this is to call it a propane. But, it's got instead of just hydrogens, one of the propane side groups instead of being a simple hydrogen has this CH3, which is a methyl group.

So we could call this one a methyl propane. And, we could further number the carbons from left to right. So, the first one is the number one carbon. The second is number two. The third is number three, and the methyl group is attached to the one in the second, so this is two-methyl propane.

And, what we have here is two identical chemical formulas, but two different structures. And so, we all have sigma bonds everywhere. So, we call these isomers. These are called isomers. But, they are an isomer of a particular kind.

They both have the same chemical formula. But, what we can do is we can look at the constitutive groups here. So, what we can do is say, here we have, here's a CH3. Here's a CH3, and here's a CH3.

And, that leaves this one, which is a CH, whereas over here, I've got a CH3, a CH3, and in the middle, two CH2's. These are units. Some people call them constitutive groups. So, here, the constitutive groups are different from the constitutive groups in the linear configuration.

And so, these are called constitutional isomers because they have identical chemistry but different constitutive groups or constituent groups. And so, we can see this. The other thing that I need to introduce you to, OK, so this is the same.

This is just the methyl propane and the straight chain butane. Down towards the bottom, there is something called ethyl. And, that's a radical. So, we need to know the radical terminology. The radical is a species with one or more unpaired electrons, with one, in some cases more, unpaired electrons.

So, unpaired electrons living in an orbital, that's a broken bond. So, this is very highly reactive. It's highly reactive. OK, so this is, if you like, think of it as a broken bond, highly reactive.

So, we need to know these. And the ones that we need to know, are these little units that have just been drawing. These little units, so if I take methane, with four hydrogens, then I break one of the hydrogens off, and now I have a single electron sitting here, this is capable of being attached to some other species.

In this case, it was attached to the carbon backbone. So, this radical that comes from methane is called the methyl radical. Another way to denote it is with the dot up here indicating that there's an unpaired electron.

Obviously, the unpaired electron is on the carbon, but in the nomenclature of organic chemistry, people are fairly quick to make the necessary change. So, if you see that, no one is suggesting that the electrons attach the hydrogen.

It's just shorthand for it. This one here, the CH2, the CH2 that we see over in the straight chain butane, this one is called, in this case there is hydrogens above and below. So, there's two unpaired electrons here.

So, they can then link up with carbons on either side. So that would be designated as follows. Or, if you want, CH2 with two dots. And, this is called methylene. This is the methylene radical. And then, the only other one that I really care about in the alkanes is the one that comes from ethane.

C2H6 is ethane, and so the radical from this would be C2H5. And, this is called the ethyl radical. And, we'll come across some compounds, one that we meet socially is ethyl alcohol where we put the alcohol functional group on the end of the ethyl.

So, that gives you an introduction to alkanes. So now, let's go back and look at the next column. That's the alkenes. So now are going to look at unsaturated hydrocarbons. And, the first example is the alkenes.

And these are characterized by sp2 hybridization. So, that means at least in one place it only takes one. At least in one place in the molecule, there's a carbon-carbon double bond. And that will give you, if all we have is the one carbon-carbon double bond, it will give us the chemical formula CNH2N.

And clearly, N must be greater than or equal to two. So, the simplest one is C2H4, which we've seen already when we talked about sigma and pi bonding. This is ethylene, and hydrogens here. We have sp2 hybridization.

So, we have one sigma bond, and one pi bond. So, the carbon and the hydrogens lie in a plane. This is 120°. We've seen all of that before. The ethylene is the common name for it, but the name that's regulated by the International Union of Pure and Applied Chemistry is following this nomenclature.

We take the Eth because C number is two. And, we add -ene. So, the formal name for this is ethene, not ethylene but ethene. But, you can use either one. No one's going to get very agitated about it.

And then for n greater than two, the position of the double bond is not fixed. The position of the double bond: not fixed. So, we can show examples of that. Let's look at, here's one, butene. So, C4H8, so, one is to put the double bond at the very end.

So, I have just got one double bond in it. There's the 120° hydrogen, hydrogen. Now, one, two, three, four, one, two, three, four, one, two, three, four. So, this is C4H8. And this is called 1-butene because the double bond is off of the first carbon.

Or, we can put the double bond somewhere along the line. So, we can do this. So, the double bond is not at the very end. So, we have a methyl group at the end: one, two, three, four, one, two, three, four, so this would then be called 2-butene, indicating that the double bond comes off of the number two carbon.

So, 1-butene and 2-butene turn out to be constitutional isomers because they've got the same chemical formula. But they have a different mix of constituent groups. So, let's label this as constitutional.

These are both constitutional isomers. But, we can zoom in a little bit more on 2-butene and be introduced to, yet, another type of isomer. So, I'm going to redraw this because this has been drawn at right angles.

It's colloquial. We know this thing is zigzagging and whatnot. So, I'm going to redraw the 2-butene. And, so I'll begin by putting the double bond. And then, that means that I have to have 120° angles.

And, what turns out here is that the double bond goes to carbon, which has a hydrogen on one side. And on the other side it has a methyl group. And, the same thing here. So, now I have a choice. What I can do is I can put the methyl group up here in the hydrogen.

All I've done is redraw this. All I've done is redraw the 2-butene. And, I think you can see that I can put a symmetry plane here. And, what do I have? I've got both of my methyl groups above the carbon-carbon double bond.

And, both of the hydrogens below the carbon-carbon double bond. But another way I could have set up the structure, again, there's the carbon-carbon double bond. Let's put the 120° angles, will put the methyl group above the carbon-carbon bond on the left side, but below the carbon-carbon bond on the right side.

And then, we'll do the complementary positioning with the hydrogens. So, in the case on the right, which is also a 2-butene, it's also a 2-butene. But, we can see as 3.091'ers, we look at structure.

Structure is really important to us, and we say, OK, same constituent groups, but you can tell what the electron distribution is going to be different in the one on the left from the one on the right.

So, you start thinking, well, what about their properties? Guess what, they have different melting points. They have different boiling points. They have different density. You can see they're going to pack differently.

So, all of this, and yet they have the same chemical formula. So, we have to distinguish these two. They are not differed by their constitution, but they are different by their spatial layout. And, the term that we used to talk about isomers that are different in their spatial layout, spatial arrangement, is stereoisomers.

They are constitutionally identical, but structurally different. Identical constitution, different spatial arrangement. And so, we've got labels on this. The one on the right, to indicate that the various groups are on opposite sides of the double bond, is called trans.

So, the structure on the right is trans-2-butene, where the one on the left is cis-2-butene. So, now, we've met stereo isomers and we know what, oh, there's one other thing. We can have more than one double bond as well.

We can have more than one double bond in an alkene, more than one double bond in an alkene. And, we'll meet a few of these as well when we talk about polymers. If we have two, this is called a diene, and if we have three it's called a triene.

And, we'll look at one of these. So, for example, we could do something like this. Suppose I gave you CH2, you don't mind that I'm going to put the H on the inside. You'll forgive me? I mean, this is OK, right? So, you know the H's are outboard, but we're just going to go with it, C, all right, what else? CH, yeah, OK, so what are we going to do with this one? I'm not going to give you this one to name.

I just what you see how it works, and I'll make sure that we focus on the chemistry. But just for the record: one, two, three, four, five, so it's got to be something pent-. And, it's got to be an -ene because there's at least one double bond here.

And, there are two. So, this is going to be a pentadiene. And furthermore, we can say, this is carbon number one, carbon number two, number three, number four, number five, and the double bonds issue from carbon number one and carbon number three.

So, we could call this 1,3-pentadiene. So, now you see how all of this works. And, it's not so bad. Radicals: what happens if we want to use just a piece of one of these? So, the only one that comes up, well, there's two actually.

One of them comes from ethylene or ethene. So, I'm going to put the hydrogens indicate when I don't have one. So, I'm going to break this bond, throw away the hydrogen, and use this. So, this is the radical that comes from ethylene.

Well, you might say, well, why don't we call this, since it's the formal name, I'm going to use the IUPAC name, ethene. Well, we can't say ethyl because ethyl is already taken. Ethyl is the one that's used over here for C2H5.

So, we need to distinguish this. And, the name for this radical when we take ethylene or ethene is called vinyl. And, we'll meet this one because if we want, we could then stick onto there reacted with chlorine and then make, this is vinyl chloride.

And then, later on, we'll polymerize this and we'll make polyvinyl chloride. So, vinyl is the radical that comes from the compound ethylene or ethene. And then, there's one other one that comes up in the life sciences a fair bit.

If we look at propene or propylene, so it's got to have a three, so, one, two, three, all right, so this is, if I lose one of the hydrogens here and make this into a radical, the radical that comes from propene can't be propyl because that's going to be C3H7.

So, here, this will be called allyl. And, as I say, this comes up sometimes in the life sciences. So, we may find ourselves referring to that. So, I'm giving you the direct path, the short course in organic nomenclature for our future work on polymers, and on biochemistry.

The last one, let's look at the right-hand column, which is also unsaturated. That's the alkynes. And, these are characterized by sp hybridization. So these, then, are going to be hydrocarbons containing at least one carbon-carbon triple bond.

This is the capability of carbon-carbon triple bonding. And, it has the general formula CNH2N minus two, **CnH(2n-2)** for N greater than or equal to two. And, the main one that comes up that you are apt to meet is simply C2H2, which has the triple bond between the carbons, and the sp gives it the 180°.

It's a linear molecule. And, this, according to the nomenclature, should be ethyne because it's one of the alkynes. But you know this molecule as acetylene, which is used as a fuel in such things as welding torches.

And obviously, the carbon-carbon triple bond has enormous energy. So, that's one example. OK, there's a few others. One is the aromatic hydrocarbons. We need to know those because we're going to meet those again, aromatic hydrocarbons, and the IUPAC name is arenes.

These are arenes. And, the prototypical one is benzene, benzene is the main one that we need to know and its chemical formula is C6H6, so it qualifies as a hydrocarbon. And, people were mystified by it chemical structure.

And it was Kekule who proposed the following structure. He proposed a hexagon of carbons with hydrogens off the corners, and then alternating double and single bonds. So, double, single, double, single, double, single.

And, this is the way things lay until the 20th century, when in the light of data it became known from spectral evidence that first of all, all carbons lie in the same plane. All carbons lie in the same plane.

You cannot have that if you've got alternating double and single bonds because the single bonds are going to be coming out at something other than 120°. And then, the second thing that mystified people was the finding that all carbon-carbon bonds in benzene are the same length.

All carbon-carbon bonds are the same length. So, if the Kekule structure is correct, you have this situation where the double bond is the same length as a single bond, and that doesn't sit well. And so, we had to wait until Linus Pauling, who in 1931, proposed that in fact there are two structures.

There are two structures that involve alternating between the structure that I've drawn here, and the complementary structure where the double bonds move to where the single bonds are in the existing structure.

So now, we have two of these. And he said each of these is a hybrid. Remember, Pauling was the one who described hybridization in carbon in the first place. So, these are hybrids. But they are hybrids of a different type.

It's a mixing, and in fact, he proposed that the structure resonates between the two. So, these are resonant hybrids. And, sure enough, it's been found that the carbon-carbon bond length is on the order of about 1.47 angstroms.

The carbon-carbon double bond is 1.33 angstroms, generally. If you look in some of these alkanes, but in benzene, the carbon-carbon in benzene was found to be on the order of about 1.39 angstroms, which puts it in between.

So, if you take an average, we've got three double bonds and three single bonds. This is consistent with the notion of bond order 1.5. And so, today, what people do to represent benzene is, rather than drawing these two structures, or one of them, and, say, figure out the rest for you, it's to use a molecular orbital representation.

Molecular orbital representation is as follows. We draw the hexagon and a circle inside. So, this indicates the resonant hybrid that is present in benzene. And, it's also present in a few other aromatics.

And, the other thing is that what happens in terms of electronic structure. I'm going to try to draw this. My drawing isn't the best, but you're stuck with me. What I want to do is to try to look at what happens in terms of the single and double bonds from the standpoint of formation with the various orbitals involved.

And in particular, I want to look at what happens with the pi orbitals, so, one, two, three, four, five, six. Well, let's say for argument's sake that the double bond is where I've drawn it. So, this means this is formed by the combination of a sigma bond and a pi bond.

So, the pi bond involves the smearing of these two pi orbitals. So, let's indicate that with a little bit of fuchsia chalk. There is a single bond next door. Now, there's a double bond. And, this double bond involves the smearing, again, of two pi orbitals, P orbitals, excuse me, to form a pi bond.

And then lastly, there's a double bond across the lower left here. And so, the pi orbital is formed. But when we realize that all six of these bonds are found to be identical in length, then everything is brought together so that we have equal spacing, at which point the electrons, all six of these end up smearing, hence this concept of electron delocalization.

In other words, the electron is no longer confined to alternating pairs of P orbitals. But rather, it's moving amongst all six. So, this can move through the entire structure. And, you can imagine, if we, then, go back to graphite, go back to graphite, then the electrons can move through the entire solid, which explains the observed electronic conductivity in graphite.

And, I think I've got a sketch right here to indicate, oh, that's just a stick-ball model of the stereoisomers. OK, so there's a diagram from a book. OK, so the artist that worked for the publisher did a better drawing than I did.

All right, I accept that. But you can see the delocalization of the electrons here. And, it doesn't just apply to something like benzene. It can apply in a straight line. And, the recipe is for a conjugated electron system.

And, what they mean by that is alternating between double or triple and a single. So, whenever you have multiple bond, single bond, you have the possibility of pulling everything in tight enough that the electrons can move from one multiple bond to the next multiple bond.

So, this is a conjugated structure. And here's one that's shown, this is a straight-line molecule. And this one is a, I just showed you a multiple bond and an alkene. So, this is a 1,3-butadiene where you have double bond, single bond, double bond, butadiene, 1, 3, and what can happen is this can pull in so that we have a double bond here.

So, this two. This is a double bond. This is one. Two times two, two times two plus one, we have this is essentially five thirds average bond order through this linear molecule. And, that's what you see up here.

So, you can imagine that happening in other systems where we have alternating single bonds, double bonds, single bonds, double bonds, or it could be even single bonds and triple bonds, multiple, single, multiple, single, pulls everything in.

And all of these will be called resonance hybrids when they allow the electrons to move as they do. Last thing is the radical. Last thing is the radical. The radical here, if we take benzene, and now I'm going to indicate that this is the one that's got the missing electron, OK? So, there's only one electron in the orbital.

This radical, which is C6H5 is called phenyl. So, we could take something like, let's build two of these. Let's take, here's ethene. And, I'm going to break one of these off, and I'm going to attach it to phenyl.

So, I can either call this vinyl benzene, or we could call it phenyl ethene, right? I mean, are you benzo-centric? So, you could say that this is vinyl attached onto benzene, or maybe you are etheno-centric and so you say this is a phenyl attached onto the thing, but neither one of these; these are both correct, by the way, but we don't use this.

This compound is called styrene. And, what we're going to do later on his break this double bond. And, we're going to make polystyrene. So, that's styrene. And, two others that I'd like you to be familiar with are we can take and just simply add methyl groups.

So, if you take benzene and add one methyl group, this will be methyl benzene or we could call this toluene. And, there's the last one I'll introduce you to, and that's two methyl groups. So, this is dimethyl benzene, or it's called xylene.

And, this whole sequence of benzene, let's just put benzene up here for completeness, so, benzene, toluene, xylene. This BTX is used as additives to jet engine fuel to improve the octane number and improve performance.

So, let's move to that. So, first of all, a little bit about Kekule. Kekule is an interesting person. He entered the University of Giessen to study architecture but then he switched over to chemistry.

And after his Ph.D. he moved to Britain. He took a job at St. Bartholomew's Hospital in London. And, he would fall asleep on the bus going back to his flat. And the story goes that one day he fell asleep on the bus, woke up far beyond where he was supposed to get off.

And during the course of that bus ride, it occurred to him in a dream that the way to describe the structures of some of the carbon compounds that were known at that time, people could characterize them by their molecular weight.

So, they knew what the structure was. But nobody, until that time, had the vision to suggest that carbon could link to itself and form a chain. So, this is where he first proposes his chains in 1855.

Then he got a job as a faculty member at the University of Kent, went back to continental Europe, and one night he fell asleep by the fireplace. And he was dreaming about the benzene molecule. This guy really took his work home with him.

So, the story goes that he had this vision of a snake biting its tail and spitting in the dream. And that's where he came up with the idea of this Kekule structure of - remember, he already had the courage to put carbons in a line.

So now, he's going to fold that line over onto itself. And so, he's really been considered the father of structural chemistry. And so, I try to abstract, just keep the noise down. Three more minutes and then you're gone.

So, what's his formula for success? Well, he moved into chemistry from another field. So, sometimes this cross fertilization is good. And the other thing is he is a dreamer. And, I always tell people, you got to keep dreaming.

When you stop dreaming, you stop thinking big. But if you're going to dream, you have to get some sleep. I know people in this room don't sleep enough. So, I'm going to challenge you to get some sleep.

Especially get a little bit of sleep on Tuesday night. I think you will perform better on Wednesday if you've had some sleep on Tuesday night. The last thing I want to talk about is anti-knocking agents in automobiles, and also in aircraft.

If we took only straight chain alkanes for gasoline, they would burn very unevenly. Remember, in an internal combustion engine, you admit fuel as a vapor. You compress, and then you ignite. But after a number of firings, the engine chamber is hot.

And, if the fuel is too reactive, it will ignite by itself on the compression cycle. And, you want staged compression. It has to be done in concert on cue, not before its time. So, this is an unprovoked firing.

It's called knocking. And you'll hear that sometimes when you accelerate up a hill. You'll hear this pinging sound. Well, this is knocking. So, the figure of merit was introduced to 1927 is called the octane number.

So, it has an iso-octane. See, it's trimethyl pentane. Well, you know, that's five carbons, and the other three carbons to make octane are sidegroups. And, they're off of the two, two, and a four.

And then, compare that to heptane, which is absolutely abysmal. So, this is 100. This is zero. And then, you'll take whatever your gasoline is, and you compare it against this standard solution. In this case, it's 90% of the trimethyl pentane, 10% heptane.

It ends up with octane number of 90. Anything else that performs in the same manner, so when you go to the gas pump and it says octane 89, or whatever, that's where it's coming from. And, you can use additives to increase octane.

In fact, what you are trying to do is to repress ignition. High-octane fuel burns less well, but it burns on cue. That's the thing. So, you could add tetraethyl lead, which was added for years. It's now banned, or ethyl alcohol.

Ethyl alcohol will raise the octane rating. And, the way to make a mix of branched and cyclic alkanes is catalysis. You use a catalyst of alumina, silica, 450, 550°C. It's called catalytic cracking.

And so, by catalysis, we can direct the synthesis to get the right mix of the right structures, and thereby improve performance. Good luck on Wednesday.

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