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,
-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.
7 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.
9 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.
