Introduction to Chemical Biology 128. Lecture 03. Reactivity and Arrow Pushing.

[ Silence ] >> Welcome to week
two of Chemistry 128: Introduction to Chemical
Biology. I'm Professor Weiss. I'll be talking to you
today about reactivity. OK. So last week we
talked about the molecules that compose your cells. And our goal this week is to understand how those
molecules interact with each other. There are two forms
of this interaction. The first kind is that
the molecules can decide to react with each other. They can start to
form covalent bonds. Bonds can break. Bonds can form. So, we want to understand
this property that we're going
to call reactivity. And to understand this,
we're going to look at arrows and the language of arrows which
organic chemists have developed as a way of communicating
this reactivity.

I have to tell you, I
think that this is one of the great achievements
of organic chemistry. This is one of those
accomplishments that all humans can be proud
of because it reduces something that otherwise seems mysterious
to a simple set of rules from which you can derive
many, many reactions, essentially all reactions
found on our planet. And to me, that's really
exciting because that means that this language is universal and it's one that's
very broadly applicable. And so, that's my
bias going in is that I think this
is really cool.

OK. So we're going to have
quick review of arrow pushing. And then I'm going
to show you examples of applying this language of
arrow pushing and this language of reactivity in chemistry
to the chemistry that's found on our planet before
life started. This is a type of chemistry
called pre-biotic chemistry. Now, obviously there
were no humans present to observe directly
what was going on. However, we can infer
what was going on in this pre-biotic period. And this is an artist conception of what the planet
might have looked like for both fossil record
and also from experiments that have tend to
recreate the conditions that were found during
that pre-biotic period, OK? So, we're going to be using what
we've learned to really to look at syntheses of the
molecules that compose a cell. And then the next
topic we'll talk about this week is
making molecules using a combinatorial approach. This is essential
in chemical biology. This combinatorial approach
takes place in your cells. It's one of the reasons why your
immune system can very rapidly respond to foreign invaders. And it also is used in many
chemical biology laboratories around the world.

And so, for this reason, I
have to introduce this concept of combinatorial chemistry and combinatorial
biology to you this week. And then finally, we'll
look at the second mode of molecules interacting
with each other. Recall at the start of this,
I said there were two modes. The first mode being reactivity
that forms covalent reactions that results in covalent
changes, bonds forming and breaking. The second mode is
non-covalent interactions. This is when two molecules slide
along side each other and decide to form a complex
with each other.

And the rules that
determine whether or not this complex
forms are also rules that we can understand. And importantly, this is
also a really tough frontier for chemical biology. So, while I'll be able to
tell you about the rules for reactivity and
covalent bond breaking and bond forming reactions, I
cannot reply with such certainty where we start talking about
non-covalent interactions. There's a lot less
that we understand and that makes it one
of the challenges. But at the same time, it
also makes it really exciting because that means there's
opportunities for people like yourself to get out
and do new experiments, to start to elucidate
those types of rules. OK. So, I have some
announcements before we go in that's kind of the overview. Let's zoom down and
look at the particulars. First, for this week, I'd
like you to read chapter two in the textbook,
that's this book here. Now, there's occasional
times where the treatment in the textbook is more advanced than what I'm talking
to you about. For example, there's
information about inversion of phosphate geometry–
phosphorus geometry.

I'm not going to discuss that. And if I don't discuss it in the
lecture, then don't get too hung up on it in the book, OK? So, simply skim the concepts that are not presented
in the lecture, OK? So if I don't talk about in
lecture, it's not important for the class in
terms of our exams and what I'll be testing on. Simply skim through it. Homework. I'd like you to
work chapter two problems, in particular, every
odd problem. And there will be a worksheet to
guide our discussions this week which will be posted
to the class website. In addition, there will
be one handout this week which will be posted
to the course website. Please download this on
Tuesday to skim through it. This handout is an example
of a journal article report which I used to call
a book report.

And then on Thursday,
I'll discuss it with you on further detail, OK? At this point, I would usually
ask if you have any questions. If you do have questions
then you can either e-mail me or the TAs. OK. So, let's review where we've
been and then we'll get started on what I told you about at the
beginning as our big picture.

So, we want to understand
the function of human cells at the level of atoms and bonds. This is the smallest unit that actually is meaningful
to ask as chemists. And as I described
to you last week, cells are bags of molecules. They are bags that are
chock-full of molecules. The molecules are
stuffed inside the cells. There is no elbow room. These things are
jam-packed into the cell. So, because of that, we expect
lots and lots of interactions which is our topic
for this week. But I'm getting ahead of myself. Let me continue reviewing
what we talked about in the previous week. First, we talked about
the composition of a gene as an on/off switch
with instructions.

We talked about how
molecules are synthesized in the cell using the template
of DNA to a messenger RNA which then is translated into
proteins, and then proteins and RNA carry out all
the various instructions that are articulated by the DNA. We also discussed six
types of organisms. But in this class, we're
going to be generally talking about either bacteria
or human cells, OK? And it turns out there's
a lot of chemistry in just bacteria or human cells.

So our goal in this week
is to reduce the complexity of diagrams like this down
to a few rules that chemists like ourselves can understand. OK. So, let's get
started with what is life? What is this stuff? What are the molecules
that compose cells? What are the rules
that govern them? In 1948, the physicist,
Erwin Schrodinger, wrote a very influential
book called "What is Life?" I highly recommend
this book to you. It's a slim little volume
and it's a fun read. It's not particularly
challenging. But the concepts
that he presents to me are really
earth-shattering. These are paradigm changing. What Schrodinger argued
is that the molecules that govern your cells, that
allow organisms like, you know, yeast and bacteria in humans
that allow organisms to live, those molecules are governed by
physical laws, by the same laws that we talked about in
chemistry and physics classes. There was nothing
special or unique about the molecules found
in living organisms. They are simply molecules that are governed
again by physical laws.

So, this persuaded– this
book persuaded a generation of physicists to explore
biology after World War II. This was an amazingly
influential book. And this persuaded
this generation that include great
scientists like Francis Crick and Jim Watson and many
others to explore biology and to do this by applying
concepts from physics and concepts from chemistry. And the results are more
or less what I presented to you last week when
we talked a little bit about the structure
of molecules. So, this is good news for us. The good news is everything
that you've been learning about in chemistry classes
before now applies to biology. There's nothing special
about biology. There's no sort of life force that animates molecules
found inside the cell. No. Rather, the same
rules that you learned about in general
chemistry, that you learned in organic chemistry, those apply to the molecules
found inside your cell. OK. So, let's talk
a little a bit about those molecules
found inside your cells.

So, our goal is to
understand first the reactivity of those molecules. And then second, we'll talk about their non-covalent
interactions. So, covalent interactions,
reactivity first. In organic chemistry, you
learned the powerful language of arrows which are a way
of depicting the overlap of molecular orbitals. Let me remind you of some
conventions of those arrows and the conventions of this
language of organic chemistry. So, the first to this is that
these arrows depict the overlap of molecular orbitals so that
they show for example electrons in a highest occupied
molecular orbital overlapping with the unoccupied lowest
energy, molecular orbital of the second reactant
of the reaction. OK. So, in this basic
reaction we have an amine, we have a ketone and the
two of these are going to be reacting with each other. So, if you take amine
and you take ketone and you mix them together,
we can predict in advance that our reaction
will take place.

And here's why. What we can predict
is that the lone pair on the nitrogen is going
to be highly reactive. Why is that? What's special about
that lone pair? It happens to be
very high in energy. It is a highest occupied
molecular orbital. And it's going to want to
react with the pi bond, the– this carbonyl functionality
of the ketone.

What's so special about
the carbonyl functionality of ketone? Well, it happens to be– it happens to have a low energy
unoccupied molecular orbital, OK? Now, let's break down what these
molecular orbitals actually look like. What this looks like is the
lone pair of this nitrogen on this nitrogen is
found in an N orbital, so it's in a high energy state. It's the highest occupied
molecular orbital, the HOMO. And it's going to be overlapping with the lowest energy molecular
orbital of the carbonyl of the ketone which happens to be the antibonding
orbital of the pi bond. OK. This is what it looks like
in terms of molecular orbitals, and this is what it looks
like on top in terms of organic chemistry and
organic chemists speak. Good news, we organic chemists
have agreed to the convention that we will depict
complicated reactions like this one using the
simplified descriptor, OK? And this is good news.

I don't think anyone wants
to spend lots of time on the test deriving what these
molecular orbitals look like and trying to describe an
antibonding orbital in terms of the lobes and
so on and so forth, it would be just
way too complicated. So, we're going to be using
this description here. Now, the real challenge
for us comes from the fact that the molecules
that we talked about in biology often times
have multiple functional groups. It's not a typical for a
biomolecule to have say hundreds if not thousands of carbonyls or
to have thousands upon thousands of different lone pairs. So the real challenge is
for us to figure out which of those lone pairs and which of these carbonyls is actually
going to engage in a reaction. And when that happens, we're
going to fall back on orbitals to decide which of these is
going to be most reactive.

OK. So again, what we're
going to be talking about is this overlap
of molecular orbitals. That overlap of molecular
orbitals, the filled-unfilled overlap
leads to the formation of new bonds and consequence
breakage of others, OK? So, when this lone
pair overlaps, so if the antibonding
orbital of the carbonyl, the pi-star orbital
of the carbonyl, the result is a new
covalent bond directed by this first arrow. Now, on the other hand, we know
that this carbon can't have more than five bonds to it or can't
have more than four bonds to it and so, five bonds
would be disallowed. And so, for this reasons in
consort with this formation of a new bond, there's breakage
of the pi bond between carbon and oxygen of this ketone. This is good news, right? This totally makes sense because
what we're doing is we're populating this antibonding
orbital. And in doing so, we're
making the orbit– we're making that pi
bond break, right? If you put electrons into
an antibonding orbital, what does it do? The bond breaks hence
the name antibonding, OK? So, this overlap, to me, it's
kind of like the peanut butter and jelly of organic chemistry.

We're always going to
be talking about a HOMO, a highest occupied
molecular orbital overlapping with the lowest unoccupied
molecular orbital. And in the same way that
peanut butter and jelly taste so good together, orbital
overlap works so well. It is so complimentary
in terms of reactivity. OK. So, let's get back
to our challenge again. The challenge is in biology, we often times have
many different puzzle of reactivities, we often times
had many different possible reaction mechanisms
that we can draw. Despite that plethora
of possibilities, what we will see is that
there's often times one and only one true
mechanism, dominant mechanism for a particular
set of molecules. And again, this is
good news, OK? So, for example, let me show
you sort of an easy case where we're going to be
looking at reactions, two possible mechanisms. One that makes chemical
sense and one that does not. And so, by dong this, we
can start to eliminate a lot of different examples. OK. So, here is a clash
of two possible wills. In this reaction, one possible
mechanism has the lone pair attacking the antibonding
orbital of the carbonyl and going through the transition
state that's depicted down here.

This reaction is an addition
elimination reaction, it goes through this
transition state in addition, and then in the elimination
reaction, the chloride is eliminated
giving us a substitution of nucleophile in
place of chlorine, OK? Makes sense, fundamental
reaction. A different type of reaction
mechanism might look like this, where the nucleophile directly
displaces the chloride. In doing so, the nucleophile, lone pair on the nucleophile
is populating the sigma-star antibonding orbital of the bond between the carbon
and the chlorine, OK? So, two possible mechanisms. One involves the
pi-star orbitals, this one involves the
sigma-star orbitals. And I guess I first blushed
these two reaction mechanisms might both look totally
legitimate and both equally valid. The problem is they aren't, that we can actually readily
eliminate the reaction mechanism on the right that– that is
governed by the SN2 reaction. Instead, what we can do is
actually very quickly decide that only the addition
elimination reaction will work. So, returning to this possible–
this clash of two wills, when we look at a
transition state or reaction coordinate diagram
for the two possibilities which I think tells us
which possibility is correct and which one is wrong.

OK. So this transition– this reaction coordinate
diagram is depicted over here. So, in one reaction I showed
on the previous slide, the mechanism is an SN2
reaction and on the right, this is the addition
elimination reaction, OK? So, in this I acknowledge, this is a complicated
diagram, bear with me. So, over here, these are
the starting materials. This is the acid chloride. This is the nucleophile. And again, if this reacts
through an SN2 reaction, you will get this left
reaction coordinate. And if for reacts through an
addition elimination reaction, you get the right coordinate. Now, two possibilities,
small little hill, big hill. Which of these two is preferred? Small hill, big bill? All right. Now, let's just imagine, you're
in electron, you have to decide which one would you prefer. Would you prefer tramping up the
very, you know, steep key slope, or would you prefer
the much shorter hill? OK. I will tell you also
that electrons are lazy that they do not expend any
extra energy than they need. And in doing so, they are going to prefer very strongly
the tiny little hill or the much smaller hill of the
additional elimination reaction to the SN2 reaction, OK? This makes sense.

That's the way electrons
live their lives. So, what this tells
us is that yes, there are two possible reaction
mechanisms for this reaction, yet only one is actually
correct. The only one that's correct
is this one on the right. The addition elimination
reaction, one on the left has to go through a much
higher energy SN2 reaction. OK. Now, I'm going to
explain in greater detail in a moment why it is that the
one on the right is preferred than the one on the left, OK? To understand that,
I need to tell you about three possible
components of orbital overlap.

So, the energy in this
interaction is proportional to three components, OK? And let me go back. Recall over here that on
reaction coordinate diagrams, the Y axis depicts energy
where a higher number up here indicates
higher in energy. And again, electrons being
lazy prefer lower energy, OK? So, that again is why the
smaller hill is preferred to the bigger hill in terms
of which side to go on, left side or right side. OK. Now, this energy
is proportional to three components. Component number one are
charge-charge interactions, OK? So, if these molecules
happen to have plus charges and minus charges then, that
will have some interactions, some Coulombic interaction. In addition, if the molecules
have a repulsive interaction with each other, that will also
contribute energy as well, OK? So, charge-charge
interactions, these are govern by the social convention
like opposites attract, OK? So, in social circles,
opposites attract, I think is commonly accepted.

It works as a formula
for dating websites. It also works recently well as
a formula for molecules as well. So happily social conventions
mirror atomic formulas, OK? So, charge interactions
are one possibility. If I go back, you could see that we don't really have any
charge interactions operative in this mechanism
as depicted here. Nucleophile is neutral,
acid chloride also neutral, charge interactions
off the table. Second term, repulsive
interactions. So, this would be if the
molecules have some sort of steric hindrance
that prevents them from overlapping
with each other. And this is a really
important component in terms of preventing molecules
from interacting. It's used extensively
in biology. It's used extensively
in enzymatic catalysis. Again, over here, that doesn't
seem to be a possibility, right? The nucleophile has a
wide open lone pair. The acid chloride similarly wide
open, it has just a methyl group on this attached to it. So, there's really no
repulsive interactions that are operative here. And by the way, just to
remind you, the repulsive term of this equation is the
term that allows this hammer to get pound– to pound in
the nail in this wall, OK? So, these repulsive
interactions, that's basically the
Pauli Exclusion Principle, that means that electrons
cannot occupy that more than two electrons cannot occupy
the same molecular orbital, OK? And so, for this reason hammer
starts pounding on nail, nail goes into the wall to
get away from hammer, OK? They don't, you know,
suddenly merge with each other and magically start
to create some sort of hybrid material, OK? Things don't happen that way.

OK. So repulsive interaction
is clearly important, not so operative in
this reaction, right? These two can snuggle up
as close as they want. There's no, you know,
prevention of that by, you know, steric shrubbery. Last one, attractive
interactions. This third term, I would
describe as mysterious, right? This is not the term that
we used to talking about. This attractive interaction
is nothing more than the filled-unfilled
overlap that I've been talking to you about today, OK? So, here reduce down to
its terms is a different representation of the same
equation one from up above. In other words, their reaction
energy for a particular set of interactions is
proportional to Coulomb's law which governs charge-charge
interactions plus the steric terms, minus the filled-unfilled
orbital overlap, OK? And it's this third term over
here that governs whether or not the molecules actually
get to form and to break bonds.

OK. Now, here is the
deal, the problem is that these three terms interact
in a complicated way, OK? That if we go out
and just, you know, start applying this equation to every possible social
situation we find ourselves in, we're going to have trouble, OK? And I guess the most
obvious thing is, you know, the opposites attract rule
only carries you so far, OK? Before you get married
to your, you know, snugly significant someone, it
might be a good idea to find out whether that opposites
attracting carries over to, you know, I don't know,
temperature of the bedroom or something like that, OK? So, for this reason,
this equation over here is a good
deal more complicated. Why don't we take a look? OK. So, opposites attract. Here's an example, we have
hydroxide, we have nitrogen. If they attract so much
negatively charge hydroxide, positively charge nitrogen. Our first instinct might be to
try to attempt to draw a bond, an arrow between the lone
pair on this hydroxide and the positive
charge on the nitrogen.

That would be wrong,
wrong and wrong. It will be totally wrong. And the problem is that this
is wrong at every level. The results here would be
a fifth bond to nitrogen. And nitrogen being
in the first row in the periodic table
cannot possibly handle such a large number of bonds. Remember, first row at the
periodic table, carbon, nitrogen, oxygen
cannot handle more than eight electrons
around the atom, OK? That's four bonds.

Five bonds totally wrong, OK? Another big problem with this that infuriates me is
notice the arrow starting on the negative charge and
moving to the positive, OK, that's wrong too because
again, arrow is supposed to depict overlap of orbitals. I'm getting a little
ahead of myself, OK? Here's the correct way to do it. The correct way to this is to show hydroxide
attacking the carbon in and displacing the
positively charge nitrogen in an SN2 reaction, OK? So these opposites attract
business only carries us so far, OK? So that's our first problem. Is that this is–
that this really, this charge-charge
interactions is very rare to provide an operative
mechanism in organic chemistry and for that matter in
bio organic chemistry.

Really charge-charge
interactions are very important for non-covalent bonding. Not so important for
covalent bonding and in fact, potentially very, very
misleading, so cautionary note. Instead, we need to turn to
molecular orbital theory. Molecular orbital
theory can explain the otherwise unexplained. And I'll give you one example
of this before we go back to our canonical example that
I showed you earlier, OK? So, for example, this methyl
ester has a preference for the syn conformation
versus the anti conformation.

And to the first approximation, this should strike you
as rather odd, right? Because in this case over here,
the methyl group is as far as can be away from
the lone pairs that populate the oxygen, right? Those lone pairs that stick up like Mickey Mouse
ears above the oxygen. And so, this anti
conformation showed to a first approximation appear
to be the preferred orientation. But, you know, when we
look closely at this and we can using various
spectroscopic techniques, what we find is actually the
dominant conformation is the syn conformation.

And you can start to understand
this if you think about overlap of molecular orbitals. OK. Here is– you know, here's– again, that's the syn
conformation should appear to have some steric clash. But again, molecular
orbitals explain why it is that it doesn't prefer that. OK. So, I keep talking
about molecular orbitals. I think it's time for us
to dive right in and start to dissect them and look
at them in greater detail and let's get started.

So, in molecular orbital theory,
we're going to be talking about atomic orbitals. So, the atoms of a molecule each
have atomic orbital associated with it, OK? So the nitrogen has
some atomic orbital. The oxygen, the carbon, even the
hydrogen has some little tiny molecular orbi– or some little
tiny atomic orbital associated with it. Those atomic orbitals are found
in S, P, D and F orbitals, OK? That's where the
electrons hang out. They hang out in
shells or orbitals. I prefer the word orbital
which describe their orbit as they orbit around the
nucleus of the atom, OK? And remember those electrons, that's the business
end of the atom. That's what endows it
with functionality. That's what makes molecules
the way they are, OK? Now, here's the thing. Often times, these
electrons are not simply in either an S orbital or P orbital instead
they typically hybridize into hybrids of S,
P orbitals, OK? And we're used to this concept. These hybrid atomic orbitals are
given the names SP3, SP2 and SP. Here's the important part, OK? So, this is review, I know that you've seen these hybrid
atomic orbitals before.

This is the part that matters
to us as chemical biologist and bio organic chemist. The S-character of these hybrid
atomic orbitals determines its stability. This totally makes sense, OK? So, an S orbital is a sphere
where in the very center of the sphere is the
nucleus of the atom. Nucleus is positively charge. The sphere defines the
orbit of the electrons. And in the sphere, those
electrons can cozy up as close as possible to the
positively charged nucleus, OK? So, this is an– a great
example of opposites attract and the attraction
equal stability.

On the other hand, a
P orbital as depicted up here has a nucleus, a node between the two
lobes of the orbital, OK? So, the nucleus is right
here in the center again, but that happens to
be a zone of exclusion where the electrons
are not allowed to exist rather the electrons
in this orbital are hanging around either in
this node up here or this other node down here. They're not in– oh sorry, lobe. This lobe up here and/or
this other lobe up here, they are not allowed
to get up too close to the positively
charge nucleus.

And so, for this
reason, the S-character of a hybrid atomic orbital
determines the stability of that orbital, OK, of the
electrons in that orbital. Conversely, the P-character
defines the instability. It defines how reactive and how
nucleophilic those electrons in that hybrid atomic
orbital really are, OK? That's kind of like, you know, defining how unhappy
the electrons are, OK? Happy electrons are found
in the spherical S orbitals, unhappy electrons are
found in P orbitals. And what happens when electrons
are in unhappy situations? Well, they will move. They will do everything they can to find more stable
orbitals for themselves. OK. So, these are
the atomic orbitals, specifically the hybrid
atomic orbitals over here, and P-character conveys–
confers reactivity and basicity.

So, for example, if we look at
a series of lone pairs found on carbon, what we find is that the higher the P-character
the more reactive that result in lone pair will be, OK? And this could be
dramatically illustrated in terms of basicity, OK? So, here's a lone pair in
an SP3 hybridized orbital. Its pK is 50. Compare that against
a lone pair in an SP or an SP2 hybridized orbital. The difference here
is truly dramatic, OK? So, the pKa is only
41 in the case of the SP2 hybridized orbital
and then, it's way down at 24 down in the SP hybridized
orbital. This is an enormous
difference, OK? Remember, pKas are a log scale. So in other words,
this guy up here is 10 to the 26th times more reactive
than this guy down here.

And by more reactive, I
mean, how avidly it's going to be reaching out
and ripping hydrogens, ripping protons off
of its neighbors, OK? And this tells us almost
immediately that for example, you know, organic metallic
compounds are going to be extremely avid at
grabbing protons to the point where they're nearly–
they're incredibly flammable and nearly explosive. OK. Now, this 10 to the 26th
times again is huge, right? That's a 1 followed by 26 zeros. It's such large number. It's hard to actually for
us to even imagine it. OK. So, enormous
difference is determined by this P-character,
S-character. I hope by now, everyone who's
listening to us and everyone in my class can explain why
it is that these guys are so much more reactive
than these guys. And it should make sense just from geometric considerations
as depicted here. Now, these hybrid
atomic orbitals recombine into molecular orbitals
in molecules, OK? So, the hybrid atomic
orbitals only carry us so far. More often, these hybrid
atomic orbitals are shared between atoms and that sharing
is what gives us bonds, OK? Now, these molecular orbitals
are given the names sigma, pi and N, OK? So, this hybrid atomic orbitals
form bonds with other atoms and that yields molecular

The energy of these
molecular orbitals is defined very specifically. And there's no way around this. I basically just have to
tell you, I'd like you to memorize this chart
on this slide, OK? So, please memorize the
order of this reactivity where a sigma molecular
orbitals are lowest in energy, pi are higher in energy,
N are even higher, OK? So these are the fields
on molecular orbitals. These are molecular orbitals
that have electrons in them. And these electrons are
depicted by the up arrows and the down arrows, OK? That's a convention
that you've seen before.

OK. Now, sigma makes sense. Sigma are the molecular orbitals
that define a single bonds, S for single, S for sigma. Pi, this defines double bonds
and that's convenient, right? Pi looks kind of
like a double bond. And the electrons in N
orbitals are the lone pairs that are hanging out
around the atoms, OK? So, when the N orbitals are
present, those are going to be the highest occupied
molecular orbitals. So, almost immediately, that
clues us in that we need to pay attention to
those lone pairs, OK? What about the unfilled
molecular orbitals that we're going to
encounter in chemical biology? These will be found in
three molecular orbitals. And again, I need to ask
you to memorize these– the order of the energies, OK? The lowest in energy
are P orbitals. P orbitals are exactly
what I showed you a couple of slides ago, OK? That's them.

These over here. This is what a P orbital
looks like, it has a lobe and another lobe down here. P orbitals we find when we
look at carbocations, OK? The empty hole that is the
carbocation is a P orbital, OK? The other electrons that
surrounds the carbon, that surround the carbon
or the carbocation, those other electrons are in an
SP2 hybridized atomic orbital. So, the remaining empty atomic
orbital is a P orbital, OK? So, most of the time, we don't
really have carbocations. The reason for this is that they
are extremely reactive being so low in energy and so, for
this reason in biology we very, very rarely find carbocations. In week eight, I'm going to
show you an exception to this.

But for now, let's keep in
mind that we're just not going to see this very much. And again, the reason is
biology takes place in water and carbocations react
avidly with water. Pi-star, this is the
antibonding complement partner to the pi orbital. And sigma-star is the
antibonding complement to the sigma orbital. And again, we're
seeing this relationship where pi-star is lower in
energy than sigma-star. OK. So, here's what
I need to tell you. Good news. You don't have to worry about where all those
electrons are in a molecule. And it's really fabulous
news, OK? If you just stop, take a
moment, take a deep breath, pause and appreciate this. Because the molecules we
talked about when we talk about biology are fiendishly,
fiendishly complicated, OK? This goes back to the business
that I talked about earlier of the hundreds, if not
thousands of pi bonds, the thousands upon
thousands of lone pairs.

The good news is we get to
simplify all of that complexity down to just worrying about
the frontier orbitals, OK? So in other words,
we only have to worry about the frontier highest
occupied molecular orbital and the frontier lowest
unoccupied molecular orbital, OK? In other words, all
we have to do is focus in on the highest
occupied lone pair or highest occupied molecular
orbital that has a lone pair in this N orbital or/and also
the lowest unoccupied molecular orbital over here. So in other words, if there's
an available P orbital, it's going to react first. OK. If there's a carbocation,
everything else will come to a halt and carbocation
gets to stay in the sun, it gets to dance around, OK? If there is a lone pair, lone pair will be the
dominant reactivity, OK? This is good news.

OK. It simplifies everything. We just have to look for
the highest energy, HOMO, and the lowest energy, LUMO, OK? What does this mean? What this means is that
this highest occupied, a HOMO is the filled
frontier orbital. And this is the orbital from
whence all nucleophilicity, all basicity springs forth, OK? And I apologize for the
kind of antiquated English but really that's how
we think about this, OK? This is the orbital. That is the business end of
this complicated molecule. It doesn't matter how many
possible lone pairs it has.

It doesn't matter how
many different possible configurations it has. All that really matters
is its highest high and its lowest low, OK? Again, this is majorly important because it simplifies
things for us. OK. So this HOMO, highest
occupied molecular orbital is the filled frontier orbital and it's a nucleophile
in reactivity. OK. Now, the intrinsic
nucleophilicity is governed by the energies of these
molecular orbitals where again, the highest in energy
is the N bonding, the non-bonding molecular
orbital that has the lone pair and the lowest in
energy are the electrons of the sigma or single
bonds, OK? To reduce it down to simplest
terms, we're never going to be really seeing reactions
that start with sigma bonds, OK? It just doesn't happen
in chemical biology.

Most of our reactions are
going to spring forth, are going to spring
from lone pairs that are in non-bonding orbitals,
occasionally electrons and pi bonds, but really–
we don't really have to worry about electrons and sigma bonds. We know they're there, you know
they're there, they're there, but we don't have to
get wrapped up in them. And this again is good news
because there's a huge number of electrons in this complicated
molecules that have, you know, thousands upon thousands
of atoms. OK. What about the lowest energy
unoccupied molecular orbital or the LUMO? This is the unfilled
frontier orbital and the lowest energy
unoccupied is the most available molecular orbital. This is the molecular orbital
that's going to be the center of attention for reactivity, OK? And again, where you have these
complex molecules, this is kind of like the funnel to which all
reactivity zooms towards, OK? Now again, we need to know this
order over here where P is lower in energy than pi-star
which in turn is lower in energy in sigma-star.

So, if we're given a choice of
different sites for nucleophile to attack, nucleophile
will choose every time to attack the P orbital
because it's lowest in energy. And again, we see P orbitals
when we look at carbocations. If there are no carbocations
present which again, I had said earlier, is
exceptionally common because carbocations are
very, very rare in biology where biology takes
place in water. So, if there are no
carbocations present, we can eliminate this one and we start focusing
on pi-star orbitals. If there are pi-star orbitals
that are available for reaction, then it's likely that
this will be dominant, the dominant reaction. Occasionally, you
come across a molecule that doesn't have a
pi-star in which case, then you might have an
attack on a sigma-star.

This is rare, OK? Especially it's depicted here. This is utterly wrong. It's depicted in the slide. I find it offensive,
but I'm stuck with it. This might happen for example
if there was a sulfur here, then you might have the
sort reaction taking place. For now, let's keep in
mind that we're going to probably be having
reactions or electrophiles in our reactions are going to
be molecular orbitals consisting of antibonding pi bonds, OK? So, it's the pi-star or antibonding pi
molecular orbital. OK. I want to switch gears. If you have any questions
about molecular orbitals or hybrid atomic orbitals, don't
hesitate to shoot me an e-mail or talk to the TAs, come to
my office hours, et cetera. We now have to talk– I
told you about what you do to decide what the
reaction mechanism is.

We now have to talk about
how to actually tell me what that reaction mechanism is, OK? So, often times at
chemistry we have some notion that molecules are reacting
but we need a clear way of communicating
that reactivity. So, organic chemists have
developed this wonderful vocabulary using arrows. And so, let's take a closer
look what those arrows are. The arrows are going
to be starting from highest energy occupied
molecular orbitals, the HOMOs, and they're going to be ending on the lowest energy unoccupied
molecular orbital, the LUMO, OK? This is a golden rule.

This is a rule that
always applies, OK? Your arrows, start on
orbitals, they end on orbitals. They start on HOMOs,
they end on LUMOs. And again, they're always going
to start on the highest HOMO and the lowest LUMO and
they'll end on the lowest LUMO. OK. So again, that lowest LUMO
is the lowest energy unoccupied molecular orbital,
that's the most available and in turn that's
the most reactive. Now, the problem is again, we
often times have many HOMOs, we have many LUMOs, what's
an organic chemist to do? What a student supposed to do? So, when in doubt, refer
to this idea of looking for the highest HOMO
and the lowest LUMO. I can simplify it, cut it down
to make it even easier for you. Most of the time, just start,
put your pen on a lone pair, and start pushing
electrons to end on the best electrophile, OK? It's not easy.

If you're in doubt,
you're stuck there at your desk during an exam,
you don't know where to start, put the pen on the lone pair
and just start drawing, OK? End the arrow on the
best electrophile, nine times out of 10, 99 times
out of 100, maybe even more, you'll get the answer
right just by doing that. OK. So, I need to talk
to you about some rules. We have rules because
it's a language. And in order for us to
be clear in what it is that we're communicating, we
need to have some conventions. OK. The conventions
we're going to follow in this course are
the following. And by the way, before I
present these conventions, I should tell you, I'm a
stickler for these rules, OK? If you give me something
that doesn't have– doesn't follow these three
rules, chances are even if it's correct conceptually,
it won't get full credit, OK? And the reason for this is it's
kind of like turning in an essay that has incorrect grammar
to your English class or something like that.

What's your English
professor going to do? You know, give you
an A for great ideas and a C for bad English? No. Your professor is probably
going to give you a C overall because the goal was to
communicate effectively. OK. So in the same way, when
we speak using the language of arrows, we have to
follow these conventions because this is what
convinces us that we know what we're
talking about, OK? So the conventions are arrows
never indicate the motion of atoms.

And this is one that if
we simply think about it, actually it's kind of profound. I think that all of us are
used to having arrows showing, you know, football
player who's over here, let's say the quarterback
moves back here and then gets behind this guy and then another arrow shows
this guy moving forward. Those were the kind of arrows
that you've been drawing, you know, I guess since you
were able to draw arrows, OK, which is to show motion, to show
the fourth intervention really, to show some element to time.

Organic chemistry, we don't
use arrows in that way. Rather, we're using arrows to
depict overlap of orbitals. We're not depicting
it in terms of time. We're depicting it in terms of
thermodynamics, not kinetics. In other words, we're
depicting an overlap of orbitals that's allowed, OK? So, arrows do not indicate
the motion of atoms. Yes, it's true. The atoms must cozy
up to each other. That's kind of understood. That's lurking in
the background. But that's not really
what arrows show you. Arrows never start
or end on charge. Since these arrows are depicting
the interaction of filled and unfilled molecular orbitals,
charges are relevant, OK? Charge, formal charge is one
of those nice conventions that makes the lowest structures
so much easier to understand, yet the charge itself
does not show you where the electrons are. It doesn't show you anything
about the molecular orbital. And so, drawing an arrow
from one formal charge to another is worthless, OK? So again, arrows never
start or end on charge.

Arrow instead– here's the one
that really embodies everything. Arrows begin with lone pairs,
with pi bonds or sigma bonds and end on unfilled orbitals. OK. And I want you
to be really precise about how you draw these things. That precision indicates that you understand
what is going on, OK? And drawing your arrows to
precisely end where it is that that pi-star orbital should
be, you're telling me something, you're telling me a story,
you're telling me where it is that those electrons
are going to appear.

And in doing that,
you're describing to me the reaction
that's talking place, OK? So I need to have all of
these things taken care of when it comes time
to, you know, for exams and things like that, OK? Make sense? OK. So, why don't we take
a look at an example, OK? So an example is this
very simple problem, OK? And the problem is we're
going to have a lone pair on the nitrogen over
here and we're going to do a simple substituted
nucleophilic attack, substituted nucleophilic
reaction that substitutes for chlorine this
lone pair on nitrogen or this nitrogen, this amine.

Everything looks good. This is a reaction very
similar to the one I showed you at the very beginning
of the class. Avoid the poisoned candied apple of simplicity rather a fall
back on HOMOs and LUMOs. Let me show you what I mean. OK. And to do that, I need
to roll up the screen. Here's what I mean by avoiding that tempting poisoned
apple of simplicity, OK. So, in this example– [ Pause ] In this example, there's a
lone pair on the nitrogen. All right. OK. So here's our
reaction again. And the simple– and I would
call it even simple minded possibility is for the nitrogen to simply displace the
chloride as an SN2 reaction. All right. So the simple case, we have
this reaction mechanism here. [ Pause ] OK. OK. This case, I would
call an SN2 reaction, right? Substituted nucleophilic two
reaction and this will give us– [ Pause ] — this guy over here. And then, this can
lose a proton. OK. So I'll show this base. The base for example
could be chloride. [ Pause ] OK. And this base
can deprotonate. This nitrogen giving
us the product. [ Pause ] OK. Now, what's wrong with this? This is totally, totally wrong
and completely unacceptable.

It appalls me. It's upsetting to me. What's so wrong about this? It's so appalling, is the fact that we're attacking a
sigma-star orbital, OK? This is an attack on
this sigma-star orbital over here, right? When we have several
perfectly good pi-star orbitals that are available, right? So sigma-star is not the
lowest energy unoccupied molecular orbital. Far from it. There's plenty of pi
star molecular orbitals that are going to
be lower in energy. Why don't we explore those as a
possible reaction mechanism, OK? So, a different mechanism
would start. OK. So I'll draw some lines
through here, a different and more correct mechanism. Well, this time have the
lone pair in that end, non-bonding orbital of nitrogen
attacking the pi-star orbital of this alpha beta
unsaturated carbonyl.

Let me show you. OK. So here's the lone
pair, it's now going to attack the pi-star molecular
orbital, electrons bounce, bounce all the way to the
electronegative oxygen, OK? So again, this is attack not
on a sigma-star but it's attack on a pi-star molecular
orbital, an antibonding orbital. OK. Now, why is this
so much better? This is better because
the pi-star is lower in energy than sigma-star. And so, for this reason, this addition elimination
reaction is greatly preferred, OK? This is actually the operative
mechanism for this reaction. OK. You could continue on,
I encourage you to do so, and in end you get
this product over here. OK. There will be many
times in this class. I'll kind of step up for you. I'm going to then let you
finish it off on your own. I apologize for that. This is upper division
organic chemistry, upper division chemical
biology where at that point where I don't have to
show you every step.

There are fundamental steps
that I want you to know, there are fundamental steps
that I expect you to know, but I'm not going to show them
to you during every lecture, OK? Rather, I want you to go home, I want you to fill
them in in your notes. I want to make sure
that you know them because on exam I will ask
you to show me those steps. But on the other
hand, I'm not going to dwell on them today, OK? We just don't have enough
time to talk about them in the class of this length.

OK. So, the lesson
from this is clear. The lesson is don't be tempted
by simplicity, instead look at the overlap of orbitals. Which one is a better overlap? Overlap of the sigma-star
or overlap with the pi-star? Pi-star is lower in energy and it's therefore
greatly preferred. OK. Let's move on. I have to lower this again. All right. Now, the other thing to make
this work is that you also have to make sure that you're
drawing the correct Lewis acid structure.

For the most part, I don't think
this is going to be a problem in Chem 128, but it is important that you set things
up correctly, OK? If you are drawing for
example five bonds to nitrogen, a lot of the reactivity of
this in oxide is not going to be apparent because
this is totally wrong. OK. Similarly, you know, in
terms of the number of bonds that you draw, this
helps you in terms of keeping track of things.

For that matter, it's
also essential for you to depict correctly
the formal charge. Oh thanks, I'm sorry. This formal charge
helps to guide us. For example, the negative
charge on this carbon over here, that should look kind
of funny to you, right? Covalence, that should
look funny, that should be extremely
reactive. So, formal charge helps
to guide us in terms of drawing these
correct mechanisms. I'll have a lot more to
say about hydrogen bonds. I don't really care
about dative bonds. We won't see them in this class. Don't– let's not
get into it today. I'll talk to you more about
hydrogen bonds in a moment. OK. So, arrows start
with bonds or lone pairs.

And here are some correct
depictions of arrows, OK? So in this case, where we're
showing a bromide leaving for an elimination reaction. The bromide takes off. And notice that the
arrow is starting at the carbon-bromine bond, OK? In other words, the electrons in
that carbon-bromine bond decide to step out the door and leave
with their friend, the bromine, giving us a bromide ion, OK? So, here's electrons that are
starting with another bond. In this case, a pi bond. Here they are starting
with a sigma bond, here they are starting with a pi
bond, and here they are starting with a non-bonding
lone pair, OK? All three of these
cases are correct. Contrast that with
these cases over here where I'm showing you
arrows starting on charges. This again is deeply
appalling and totally wrong. So, arrows do not
start on atoms. So for example like this or
like that, instead we want to draw them starting
on the bonds themselves. This should make sense, right? Arrows are trying to depict
the overlap of orbitals. They need to start
where the electrons are. The electrons are
found in these bonds.

Electrons are not found
in this negative charge. They're not really found
around this bromide, instead we're talking about
the electrons that are shared between bromine and carbon. Those are the electrons
that matter. OK. Now, that's where
they should start. Let's talk about
where they should end. So, arrows need to end
on atoms and bonds, OK? So, here's a lone pair
attacking a proton. It's ending directly
on that proton, OK? So, here is it ending on
an atoms, the protons, here it is ending on the
carbon of a carbocation, here it is ending on the
hydrogen or the proton during in a beta elimination step.

OK. So, atoms never
terminate in empty space. So for example, when bromide
is stepping out the door, the electrons don't simply
hop out and then the door over opens into empty space. For that matter, this arrow
would be wrong if it started at this carbon-bromine and then
had the electrons just going off into empty space. That's not correct. The electrons don't get to
walk off into empty space. That would be extremely high in
energy and extremely repellant, rather the electrons get to end on this bromine atom
giving us bromide, OK? So, arrows need to end on atoms. They will depict again
this overlap of filled and unfilled orbitals. OK. Hydrogen for that matter is
always attached to something. OK. I'm starting to get down to
my pet peeves but this is one of those pet peeves
that doesn't matter, OK? Hydrogen is not some atom that
kind of like it is floating around next to the molecule, rather hydrogen is
directly attached to some particular atom. And this matters a great deal because where it's
attached will determine to large extent whether or
not it's going to be acting as an acidic proton or
perhaps not acidic at all, OK? So, these terms,
proton, hydride, and hydrogen atoms are
three different depictions of the hydrogen atom, and they
have three different meanings.

They have totally
different meanings, OK? So, H plus is the proton,
H minus is the hydride, and H radical is the hydrogen. They really– we don't find
them just kind of floating around like this
in the chemistry that takes place
inside cells, OK? Protons aren't just floating
around inside the cell, rather they are always
attached to something. Maybe they're attached
to a water molecule to give you a hydronium ion. But they're not just
kind of hanging out, they're doing something, OK? Hydrogens do not like
being by themselves, OK? So, in other words,
what you want to avoid is showing proton just
kind of hanging out in space, waiting around for some lone
pair electrons to attack it. That's not what happens.

Hydrogen doesn't
get to do that, OK? Furthermore, hydrogen radical
also doesn't really occur nor does hydride really occur. OK. Rather in solution
chemistry, we find species that could either
donate a proton, donate a hydrogen radical,
or donate a hydride, OK? So, what I propose you do is
instead of depicting H plus as a reagent, instead depict
H plus as catalytic "H plus." Those quotes, unquotes
are going to tell us that yes, we mean H plus. But what we really mean is we
mean H plus that's been picked up and delivered by
some other species. In this case, that might
mean attached to this methyl, this methanol molecule
that's going to be its delivery character, or
you can even write catalytic HA where in this case,
it's HA that's attached to the conjugate
base that's going to be delivering the proton, OK? Anyone of those is fine.

It is important for you however
to follow these conventions because they communicate to me that you know what molecular
orbitals are being overlapped and it tells me whether
or not, excuse me, you understand the
chemistry that's involved with these reactions. OK. One second. OK. I want to conclude
today's lecture by discussing with you one other
element of hydrogen atoms and that's the hydrogen bond. Everyone needs a favorite bond. My favorite bond is of course
the great Sean Connery. But today, I'm going
to be talking to you about a second favorite
which is the hydrogen bond. OK. The hydrogen bond
govern so much a biology that it is essential
for us to really get to understand it correctly, OK? So, hydrogen bonds are actually
largely a Coulombic interaction.

They describe the sharing of a hydrogen atom
between two partners. One partner is going to
be our hydrogen bond donor and our second partner will
be a hydrogen bond acceptor. And this hydrogen
bond will be depicted by this dashed line, OK? So this dash line is/are going to be our convention
for hydrogen bond. We're going to use
this a lot, OK? Hydrogen bonds for example hold
together the two strands of DNA. They make molecular
recognition possible, the non-covalent
interactions between molecules. So, hydrogen bonds absolutely
essential to chemical biology. However, it turns
out that the energy of the hydrogen bond is very
sensitive to the environment and the geometry that's involved with the sharing
of that hydrogen. The geometry in this case that
I'm showing you over here is of a perfectly linear
hydrogen bond which is the best
possible example, OK? So, in this case, the lone pair
on this water on this oxygen of water down here is
perfectly positioned to share this hydrogen of this
water up here and the oxygen, hydrogen and oxygen are
lined up as a straight line.

Often times that
isn't the case, OK? So, for example, we can look at
hydrogen bonds that are found in the active sites of
enzymes and we find instead of having this neat
straight line, we get a bendy line instead. That bendy line, that
bent hydrogen bond, much, much weaker, OK? So, this is kind of
the optimal geometry, optimal hydrogen bond
acceptor which is a lone pair. Optimal hydrogen
bond donor over here. And this is kind– this will
be our canonical hydrogen bond. Now, here's one of the problems. One of the problems amongst
others is curved arrows. Curved arrows confound
us when it comes time to talk about hydrogen bonds. The reason is curved arrows and
hydrogen bonds simply don't mix. Curved arrows depict
the overlap of filled and unfilled molecular orbitals, whereas hydrogen bonds are
showing us a partnership of sorts between the
donor and an acceptor. And there really isn't this
sort of overlap that leads to covalent bond in the
case of hydrogen bond.

And this becomes
tremendously confounding. So, for example, if you want to
just show transfer of the proton on this nitrogen atom to
the lone pair of the oxygen, you might be tempted to simply
draw a hydrogen bond in here. And that would be
utterly incorrect because this hydrogen bond is
basically saying the hydrogen is somewhere between here and
there, somewhere in the middle, somewhere in the sides,
whereas over here in that case of the curly arrows,
you're saying, no, it's going to pick this
up, it's going to pick up the proton wholesale,
hang on to it for a while and give you a positive
charge on oxygen.

These are two very
different depictions. So, what we've– what we
were going to be doing in this class is showing
those hydrogen transfers as an explicit step, OK? So, hydrogen bonds are going
to be useful for us for talking about non-covalent interactions,
but not useful at all for talking about covalent
interactions, the reactivity that I've been talking
to you about today. It turns out that hydrogen
bonds that proton transfers of the sort that I showed
in the previous slide, these sorts of proton transfers over here are extraordinarily
fast, OK? There are often times
diffusion controlled. In other words, they hit the
speed limit of reactivity for reactions that
take place in solution.

That kind of speed
limit and that kind of proton transfer ability is
actually tremendously useful, OK? So, this is a diffusion
controlled reactions. So, proton transfers to and from
heteroatoms very, very fast. Proton transfers for that
matter in the same way that hydrogen bonds
require a linear geometry. Proton transfers also
require linear geometries. And I can tell you that almost
immediately this is going to annoy you. This takes away one
of the conventions that you mislearned back in
sophomore organic chemistry.

I know it was cool back then
to show a proton transfer as a neighboring oxygen, say
picking up a proton over here on the nitrogen and you have
this completely ridiculous and totally crazy four-membered
ring transition state. It galls me to even say this. Can you imagine four atoms
getting together to form, you know, some sort of very
strained four-atom ring transition state? It's totally– it's
total insanity. Even more insane, notice that
the geometry between oxygen, hydrogen and nitrogen
is not perfectly linear, instead it's bent at
a 90-degree angle. And this kind of proton
transfers do not happen this way, instead they exclusively
prefer a linear geometry. So, only linear geometries are
going to count when we talk about proton transfers. And so, for this reason, I need
to take this particular step out of your vocabulary, OK? And now, it was acceptable back
in sophomore organic chemistry, it's no longer acceptable.

So, acids and bases are required to catalyze proton transfers
and tautomerizations. So instead of showing
it like this, a much better alternative
would not even alternative, the correct way to
depict this would be to show the oxygen picking up
a proton from a catalytic acid. And then, in turn the conjugate
base of this acid acts as a base to deprotonate the neighboring
positively charge nitrogen, the ammonium ion. OK. So, in all cases, we're
going to see that acids and bases are required to
catalyze these proton transfers and that turns out
to be a general rule. And the good news for us in
terms of chemical biology is that often times or all times,
we can find abundant numbers of different molecules that are
all too willing to volunteer to be those catalytic
acids and bases and obvious examples would
be for example water. Water can be a hydronium ion
to act as a proton donor. Water can also act as a
base to accept protons and become a hydronium ion. And since all biology
takes place in water, feel free to use water
as the catalytic acid and that catalytic base.

OK. We've come quite a ways. I've shown you proton
transfers, I've shown you how to draw arrows, we've
talked about the rules that govern these
electron transfers in terms of filled-unfilled overlap
of molecular orbitals. We're now going to
transition to looking at some examples of this. And I'm going to show you
this on Thursday when we talk about the molecules
found on earth that composed all living things. So, why don't we stop here. When we come back next time,
I'll be showing you examples that apply the principles
that we've talked about today. Thank you very much. [ Silence ]

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