PROFESSOR: So what I want to do today is--
I want to introduce this to you very quickly--
is-- and I was going to show you this at the end of the last
class--
if we simply go to the far end of the scale, the picometer
scale-- you see carbon.
I'm not going to start you with carbon, that is a little dull.
But over the next few weeks--
few classes, rather, because we have to do this in fast order--
we will cover details of carbohydrates,
amino acids nucleosides, and phospholipids
and how those building blocks are put together--
their properties, their ability to interact and engage
in non-covalent interactions with other molecules
and the ability to make polymers out
of some of these, such as the nucleosides and the amino acids
and the carbohydrates, which then start
to create the richness of life.
We will also discuss today the super molecular chemistry
of phospholipids as they make micelles and lipid
bilayers, which are the key boundary of cells.
So this is very important.
And then in the following week, we'll
go to some of the bigger things like proteins, nucleic acid,
polymers-- for example, here's RNA.
So the course will literally do this--
take you from one end of the scale to the other.
So I want you to get a sense of these dimensions.
I want to mention one sort of fairly stupid thing
with respect to how chemists and biochemists talk
about certain metrics, certain distances that
are pertinent to biology and biochemistry.
Engineers tend to talk about micrometers and nanometers.
There is one unit that chemists and biologists use quite a lot,
it's the Angstrom after a Finnish or Sw--
no, not Finnish.
I think it was a Norwegian.
And that is equivalent--
10 Angstrom equals 1 nanometer.
So when you're looking at scales,
we tend to talk about Angstrom because they're
a convenient number.
But don't get fooled by this.
It can be a little bit confusing because it's
10 to the negative 10.
So a nanometer is 10 to the negative 9,
you know that quite frequently.
Picometer-- 10 to the negative 12, micrometer-- negative 6.
But the Angstrom is just a funny unit we use a lot,
and it's 10 to the negative 10.
So just to make sure there's no ambiguity about that particular
detail, OK?
All right.
So today's lecture will focus on the molecules of life.
And in particular, I'm going to, through the next few classes,
introduce you to the various molecules of life.
But first of all, we have to do a little bit
to understand chemical bonding.
And in particular, we want to look
at both covalent and non-covalent bonding
because covalent bonding is important-- it's the structure,
it's the framework.
But non-covalent bonding is what gives us dynamics.
These are much weaker forces that
can be broken and remade very readily that
are essential for things like forming
the DNA duplex, folding your proteins,
associating the lipid bilayer.
All of those are non-covalent forces
and they are dynamic because they're weak,
you can break one relatively easily as long
as you're ready to make another one in its place.
So I will spend a little bit of time on that.
And then today, we'll talk about lipids and membranes.
But first of all, let me introduce you
to some of the molecules of life in this rendition that's done
by David Goodsell at Scripps.
So up in the top corner here, you look at 2.3
is the three dimensional structure of a protein.
It's folded into a globular state
through non-covalent forces.
I brought a little 3D model of a protein for you
to look at and take a look at later.
That was one of the suggestions I made.
You could coordinate printing a 3D model
as one of your later projects.
We will learn about the forces that hold the polymer
together-- the covalent forces.
But then the non-covalent forces that
make globular structures that are
very important for function.
They're not much use as unraveled spaghetti.
They're way more useful as their three dimensional structure.
Down here in the corner is a carbohydrate.
It really looks pretty pathetic in this rendition,
but carbohydrates have a lot of value,
particularly in energy storage but also in things
like the extracellular matrix and as entities
that signal information between cells.
There's a lot of communication done
by cell surface carbohydrates.
Over here you see the canonical structure
of double stranded DNA.
We'll look at the covalent structure
of those single strands, but then we'll
focus in on the non-covalent interactions that
make the double-stranded DNA and store genetic information which
is also central to life.
And then lastly on this, but we'll cover this today,
is a lipid bilayer.
It's a fascinating supramolecule structure
that really is at the heart of how all your cells are
held in a compartment surrounded by a lipid bilayer.
So by the time we start talking about those,
you'll understand the forces that put in place that lipid
bilayer that arguably-- and I've read articles that say this--
that the evolution of lipid bilayers
is as important as the genetic code.
Because if cells did not have a surrounding,
did not have an inside where you could concentrate reagents
and macromolecules and do biochemistry,
life wouldn't exist in the same way.
OK, so let's take a look at the composition of living systems.
And remarkably, we are about 75% water.
So most proteins are very hydrated.
There's a lot of water in cells.
There's a lot of water outside of cells in the matrix.
And really, we sort of survive weird.
We survive in an aqueous environment.
And the thing that you also want to think about
is when we think about non-covalent forces,
these are forces put in place in water.
We don't live on a far distant planet
where we're in sort of liquid methane or anything like that.
So water is critical to life.
The establishment of the hydrosphere when Earth first
formed, the evolutionary events that
happen after that were really hand in hand with the fact
that it was an aqueous environment.
Because forces are different whether they
are in hydrophobic environments or hydrophilic environments.
And really, you'll start to get appreciation
for that as we move forward.
So this basically suggests that if I put one of you in a giant
desiccator and pumped out all the water I could possibly pull
out, there'd be about sort of-- depending on your weight--
40 pounds of things left behind.
Of what's left behind, the majority of it
is going to be biological macromolecules-- whoops.
And then the rest of it, that little sliver,
are things like ions and small molecules-- calcium, magnesium,
iron, manganese, those small inorganic ions
as well as small molecule metabolites that are
involved in central metabolism.
Let's now look at the macromolecules
and their sort of proportions relative to each other.
The smallest sliver of the lipids,
which we'll talk about today.
Then you have the nucleic acids that
are critical for information storage.
You have proteins, which make the largest piece of the pie.
And carbohydrates, which are is the 25%.
So you can see how important carbohydrates
are because of their proportion being relatively large.
The lipid proportion, though, is small but absolutely critical,
harking back to the membrane bilayer.
Because if we didn't have the membrane bilayer,
once again, we wouldn't have life in the same way
that we have it now.
So that gives you a sense of the relative proportions of things.
And frankly, when I discuss the macromolecules,
I really like to start with lipids because
of the membrane bilayer, but because their structures are
comparatively simple relative to amino acids and nucleic acids.
So we can get a few of the basics
of the chemical structures down and how we render them
on paper so that we can do that with lipids,
which are a little simpler.
Now life, to a chemist, they have
to sort of worry about this entire mess
of the periodic table.
But the good news for you is for biological systems,
we deal with very focused components
of the periodic table.
So those biological macromolecules
are made up largely of only six elements--
hydrogen, carbon, nitrogen, and oxygen, phosphorus, and sulfur.
So that makes the amount of stuff
you need to know about basic covalent structures way
more simple than it is for the average chemist who
has to worry about everything down here in the nether regions
and-- whoops, what are you doing?
And the things that are radioactive,
all kinds of other things.
You don't have to worry about any of that.
So the covalent bonding we will talk about
is amongst those six different elements.
And they make up 98% of the cellular mass.
And then the other components that are important in cells
are some metal ions--
the alkali and [? alkalia ?] elements.
So sodium, magnesium, potassium, calcium-- those are all
quite important in life.
And then these transition metal ions
that are really important in enzyme catalysis, for example.
But we will not cover very much of that.
But those are what are known as trace elements that are very--
transition metal elements that are very
important for biochemistry.
And then last of all, there are some rogue ones
that there's even smaller amounts in physiologic systems.
These are things like chromium, molybdenum, and tungsten,
selenium, and iodine.
And of those, certain of these elements only
are found in totally bizarre organisms.
So for example, you and I don't have much molybdenum
and tungsten, I don't think, unless it
slipped in there by accident.
But you and I definitely need selenium and iodine
as trace elements.
Does anyone know where iodine comes
and figures most prominently?
Yeah?
AUDIENCE: Thyroid.
PROFESSOR: Thyroid, absolutely.
So the thyroid hormone is a small organic molecule
with several iodines in it.
And we need-- absolutely need-- iodine in our diet
in order to build the thyroxine hormone that
deals with a lot of aspects of metabolism.
So we don't need a lot.
And if we get too much, it's bad for you.
But we definitely need traces of these elements.
Now I will spend a very small amount of time
just laying down the basics of organic chemistry bonding.
Now who have you either taken the chemistry GIR
or had high school chemistry quite recently?
Is that pretty much all of you?
And now if you didn't put your hand up, don't worry.
We're here to bring you up to speed if you need it.
Frankly, if you just know what's on the next two or three
slides, you're in great shape.
All the information that you need has been condensed .
But if it's a little bit out of nowhere,
you could come see me in office hours
and I can just run through things for you
and we can just get you up to speed.
There is no need for pre-knowledge,
I just need an idea of how much pre-knowledge you have.
So when we talk about covalent bonding
and start to think about the elements that
are critical for life, it's important to consider
the electronic structures of these elements
and why they happen to be the chosen elements, OK?
The most important thing about hydrogen, carbon, nitrogen,
oxygen, phosphorus, and sulfur is they
love to make covalent bonds.
A lot of metal ions form salts, you
know-- sodium chloride or many other different salts.
But covalent bonds are the main structure
of all macromolecules.
Strong bonds between elements, such as these six
in particular-- these six--
where they share electrons in covalent bonds
rather than form ionic interactions
where somebody gives an electron to someone else
and you have a plus-minus type interaction.
So these shared bonds are important for life.
So it's good to understand why hydrogen, carbon, nitrogen,
and oxygen, and then phosphorous and sulfur are so important.
In order to understand the covalent bonding
of these elements, it's useful to know
the electronic configuration, but you
could live without that.
The most important thing is that covalent bonds,
such as the one between carbon and hydrogen here,
reflect a shared pair of electrons--
one from the hydrogen, one from the carbon--
to make a stable covalent bond.
Because of its electronic configuration,
carbon is neutral when it has four covalent bonds.
Nitrogen is neutral when it has three covalent bonds.
But there's an extra lone pair of electrons that are not
forming bonds in neutral nitrogen.
And oxygen is neutral when it has two covalent bonds.
These could be with hydrogen, they could be with carbon,
they could be with several of the other elements.
For carbon, we don't deal with charged states of carbon
because they're pretty high energy.
They may be high energy intermediates
in an enzyme catalyzed reaction, but they're not
sitting there as high energy intermediates
in your macromolecules.
The key thing you want to notice is for all of these elements,
the valence shell is complete with eight electrons.
But these lone pairs--
and I-- or bunny ears, as people like to call them--
really feature very prominently in biochemistry and biology
because they are places for hydrogen bonding interactions.
So we run a lot on electrostatic hydrogen bonding
and hydrophobic interactions.
If we know where the lone pair electrons are,
we know one part of a hydrogen bonding interaction.
It turns out that in biology, we're
mostly at pH 7 or in that range except for a few sub
cellular compartments.
But at pH 8, nitrogen lone pair of electrons
will pick up a proton to become a positively charged nitrogen.
And you'll mostly see that as a positively charged.
So the side chain of lysine, which has an NH 3, an NH 2
at the very end of a carbon chain,
is most commonly protonated and positively charged.
So it could be involved in an interaction.
So we can consider both the neutral and the positively
charged state of nitrogen.
For oxygen, that oxygen lone pair
can pick up a proton to form the hydronium ion.
So that's a positively charged OH group.
So it would have an extra proton, using up
a lone pair and three hydrogens, or it could give up a proton
to form the hydroxide ion.
And those are the states of oxygen that are most common.
So in that, we've kind of dispatched those first four
of the six elements.
Phosphorus and sulfur are a little tricky,
but there is some good news.
Sulfur copies oxygen, so you don't really
have to worry too much about sulfur.
You'll just consider it to really
be sort of an older sibling of oxygen
where all the chemistry is very, very similar.
Sulfur, or the negatively charged sulfur anion,
are both important.
Phosphorus is different.
Phosphorus does not tend to show up
in the version that copies nitrogen.
It is capable of adopting higher oxidation states.
And all of the phosphorus you meet in biochemistry
for the most part-- there's a few odd things in weird
organisms--
is going to be in the form of an oxidized form of phosphorus,
which generally has one, two, three, four, five
bonds to phosphorus.
It can take on a higher oxidation state.
And you will see phosphorus.
Phosphorus in the phosphate form is absolutely essential to life
because it's the place where we store a ton of reactivity
for the reactions of nucleotides,
adenosine triphosphate, adenosine diphosphate,
the phosphodiester backbone in nucleic acids, phosphorylation
of amino acids to form phosphoproteins.
It's always in this state with all the extra oxygens
and that configuration of bonds, OK?
If you know this, you've got a lot of the covalent bonds
under control.
So any questions about this?
Is everyone all right?
I know it might be-- it's probably
a refresher for most of you.
The next thing I just briefly want to mention
is the most typical functional groups that
occur in biological molecules.
And you may, say, well, what does it mean, functional group?
Usually it's a place where the action happens.
If you have a large molecule that's
a bunch of carbon-carbon and carbon-hydrogen covalent bonds,
there's not a lot going on unless you can really
rip those bonds apart, but they're high energy.
But functional groups are oftentimes
where chemistry happens or biochemistry happens.
So there's the OH hydroxyl.
We, as chemists and biochemists, will
tend to use an R where we mean something else.
So we don't write out a whole structure,
we would just put R OH equals--
I'm going to just say anything.
So for example, if R was CH 3, CH 2, you would have ethanol.
But I'm keeping it more generic.
The next functional group that is important
is the carboxylate group, or the carboxyl group.
Looks like this.
Now when we look at these molecules,
you always want to sort of think where the lone pair
electrons are.
There's two on oxygen, two on oxygen, two on oxygen.
So that actually shows you where the rest of the electrons are.
This is the carboxyl group.
But in nature, in physiologic systems,
this shows up most commonly in its anionic form.
That's important because when we start
to think of interactions between enzymes and their substrates,
or the folding of proteins, we're
thinking of something with a negative charge, not a neutral.
So this group loses a proton to form the carboxylate group.
And if you want to know where the lone pairs are now,
that's what they look like.
So those are two of the key ones.
Let's now go to nitrogen. That is the neutral amine.
But as I just mentioned to you, that
will very commonly pick up a proton
and be in the positively charged state.
Now when I show you both of those guys
in the positively charged state, what you could immediately
tell me is that if I have an amino acid with one
of these groups and a nearby amino acid with one
of these groups, they could form an electrostatic interaction
between themselves-- plus and minus complementing each other.
So if you know the charge states,
you're much better off because you
can tell where non-covalent types of ionic or electrostatic
interactions occur.
So these are very important.
Then there's the phosphate group--
it's often ionized-- and the sulfhydryl group.
So phosphate-- the sulfhydryl group
is also called the thiol group.
And I'm sure I've spelt that wrong because hydryl--
they look like that.
And the most common state of the sulfhydryl--
well, not the most common--
can also appear as the anionic structure.
So that's the basic functional groups.
Now there are two more functional group assemblies
that you will see a lot in physiologic systems
that are basically composites of some of these structures.
Because when we have single building blocks,
we need to join them to each other
through different types of chemistries.
So I want to show you the types of chemistry
that you get by forming a composite of hydroxyl
and a carboxyl group and a composite of a carboxyl group
and an amide.
Because the polymer that's the protein
polymer has building blocks that have a means and carboxyls,
but they're all put together into what polymeric structure
where those groups have been joined in a condensation
polymer.
So let me just show you what those look like.
And then we'll be done with the functional groups.
So there are-- the first one--
because I've drawn them in this order, OK--
is the amide.
And the other one is the ester.
When you do these two reactions, if you do them in the lab,
they're called condensation reactions
because as you form that bond, you
kick out a molecule of water.
These are really important new functional groups to you
because your proteins are held together by amide groups.
In fact, they're so important in proteins,
we often call them peptide groups.
You'll see more about that on Monday.
And the esters are really important.
For example, in derivatives of glycerol
that make fatty acids or phospholipids,
you'll see esters occurring again and again.
The other composite group that you can also see
is with the phosphate plus an alcohol.
And what that group looks like is as follows.
And you're going to see this sort of endlessly
in nucleic acids.
Let's keep the charges all even here.
And this is what's known as the phosphate ester.
OK, and that is yet another condensation
where you kick out water.
All right, so let's just run back to this image.
And we can sum it all up.
Those are all the groups that I just described to you.
And if you want, you can go back and put lone pairs
of electrons on everything.
And then the composite groups that I
want to mention to you in particular
are the amide and the ester.
And they're very important in physiologic systems.
They are the bond that holds together
the biopolymer in many cases.
Not shown on this picture is the phosphate ester--
I've added that this year because it's
kind of important--
is a similar condensation reaction between phosphorus
and an alcohol, and that in particular
is the bond you'll see that holds together nucleic acids.
And now one sort of thing that we
won't go into a lot of detail--
I want you to notice that this nitrogen here
has a lone pair of electrons.
It picks up a proton very readily.
The amide nitrogen is not so willing to pick up a proton
because it messes up the rest of its chemistry.
So that nitrogen in an amide tends
to be observed as a neutral.
However, that hydrogen can be involved in hydrogen bonds.
OK, any questions about that before we
move on to non-covalent bonds?
Is everything clear?
Now I try to put everything in one place
so you have it in front of you.
What I've put on those two slides
is what you need to know about organic covalent bonding.
It doesn't go beyond it.
I will say there's a tiny bit of memorization,
but once you commit that stuff to memory,
you're in a good place with respect
to understanding how the molecules of life
are put together.
OK.
Now what is more important to me once we've
put those structures in place is non-covalent bonding.
Because to me, non-covalent bonding is synonymous with
dynamics--
forces that can be readily broken and reassembled, broken
and reassembled.
The energy, the strength of a typical bond between two
carbons or a carbon and a hydrogen
is on the order of 90 to 100 kilocalories per mole.
It takes a lot to break those bonds.
We can't break them at will to go and do
some biological activity.
But the range of energies of the non-covalent bonds
are far more modest.
They range from-- so this is covalent.
But the non-covalent range from 1, maybe
to about 10 kilocalories per mole.
So when you think about those forces,
they're readily broken and made, broken and made.
And what's so amazing about protein and nucleic acid
structure is that you can gradually
break a bond while you're making another non-covalent bond
so you can have the dynamics of the structure that
define a lot of its functional properties.
Because structures are dynamic, an enzyme
that's a composite of a lot of non-covalent interaction
combined a substrate can gradually
form a set of covalent bonds with that substrate
but then can start changing the shape of that structure
and that shape in order to go through a catalytic cycle
to do chemistry and then to liberate products.
That is all driven by changes in non-covalent bonding.
Subtle changes that occur without big energy
barriers that would be necessary to break the covalent bonds.
So shown at the top here, you see the average bond energy
of covalent bonds.
This small number is something, for example,
between two chlorines.
That's a pretty weak bond.
But of course, we don't have a lot of them running around.
So really, carbon-hydrogen, carbon-carbon,
they're at the higher end--
about 100 kilocalories, 80 kilocalories per mole.
The other important interactions, though, that
make up the non-covalent interactions are as follows.
So the first important one is the ionic bond.
It is also called a salt bridge or
an electrostatic interaction.
Why we give three names for this probably
comes from which type of chemist decided to define them.
They are all the same things.
They are basically interactions between a positively charged
entity, a protonated amine; and a negatively charged entity,
a deprotonated carboxylate.
Those are about the strongest of the non-covalent bonds,
but it's very variable because it depends
a lot on their environment.
If those two entities are in a hydrophobic environment,
they're going to charge right for each other
to form a strong electrostatic interaction.
But if those are out in water, each of those groups
could be solvated by water and they'd
have to give up solvation in order
to form a good electrostatic interaction.
When we talk about protein folding,
we'll go into that in a little bit more detail.
So the reason why this says very variable
is not to drive you crazy.
It's just they're very variable.
But they will still range, I would say,
from 2 to 10 kilocalories.
Come on.
So those are important-- easy to pick out.
The strongest of the set.
If Dr. Ray gives you a problem set
and starts asking you to pick out non-covalent interactions,
that's the one you take care of straight
away because it is the most obvious.
The next most important, though, is the hydrogen bond.
Now hydrogen bonds have been known
to mystify people for years because people
are like, how do I pick these things out,
how do I pick these things out?
I'm going to give you a foolproof way of picking out
hydrogen bonds so you will never be at a loss for hydrogen
bonds, OK.
Well, how do we recognize them?
They are between hydrogens that are on electronegative elements
such as oxygen--
of course, there's other things attached here--
or on nitrogen, or on sulfur.
So all of those three functional groups
can serve as hydrogen bond donors.
They can give a proton in a hydrogen bond and share
that proton between a hydrogen bond acceptor, OK.
So these are all going to be known as donors.
So you can recognize them.
This-- carbon is not a hydrogen bond donor.
Carbon's got his hydrogen and he's not
giving it away to anybody for love or money.
Its holding on tight.
So this is not a hydrogen bond donor.
OK?
Now what are the hydrogen bond acceptors
are places where that hydrogen would want to sit-- yes.
AUDIENCE: There's the two lines next to it--
PROFESSOR: Actually, they just read-- they could be double
or they could be single, but I was just putting them
so that you see that the nitrogen has one, two, three
bonds to it.
OK, yeah.
It could alternatively also be the form of nitrogen--
just to confuse you--
that has an extra proton that could be the protonated version
because that can still be a hydrogen bond donor.
OK.
Now what are the hydrogen bond acceptors?
They are any place where you have a lone pair.
So let's just think of a carbonyl group--
two lone pairs.
A hydroxyl group-- two lone pairs.
A nitrogen that is not protonated--
one lone pair.
Those are the hydrogen bond acceptors.
So as long as you know your structures
in the functional groups and you know where the lone pairs are,
you can figure out where there could be a hydrogen bond.
So all of these types are acceptors.
OK.
So in protein biochemistry, for example,
those kinds of hydrogen bonding is very, very important
to form the three-dimensional structures of proteins.
And the reason why is because in a protein,
proteins are made up of amide bonds
where this Hn can be a donor, this O can be an acceptor,
and you can get networks of hydrogen bonding interactions
to establish structures of proteins.
When a small molecule binds to a protein,
it may look to fit in a place where
it can maximize electrostatic interactions and the hydrogen
bonding interactions.
So we'll ask you to start to be able to pick out
hydrogen bonding.
So here you saw the electrostatic interaction.
Here is a typical hydrogen bonding interaction
between a hydroxyl and a carbonyl group.
I couldn't spot that very readily
unless I remembered that there were lone pairs of electrons
there, OK.
The other two ty-- any questions about that?
Any questions about hydrogen bonding?
Are you comfortable with thinking
you could derive your way to figuring out where they are?
You'll see them used a lot, so they'll
become more and more familiar to you as you move forward.
OK, good.
The last two types of interactions
are the hydrophobic interactions and van der Waals forces.
I never get the spelling right, but I'll
get the concepts over you.
Now hydrophobic interactions are incredibly important.
So when you think of folding a protein driven solely
by electrostatic interactions and hydrogen bonding,
you have a bit of a problem because all of those groups
are hydrogen bonded to water.
So you'd have to get rid of the water
before they could make interactions with each other.
Does that makes sense?
Because we are in water.
We're folding in water.
Hydrophobic interactions are really great because they
want to form in water.
If you're making, you know, a batch of salad dressing,
oil and vinegar, and you shake it up, what happens?
It separates.
The oil goes to the top, the vinegar goes to the bottom.
Why?
Because of hydrophobic interactions in the oil phase.
So if you have a large protein that
has a bunch of hydrophobic groups,
they will want to collapse out of the water
to interact with each other.
And then hydrogen bonding and electrostatic
will fall into place.
So hydrophobic interactions are a very important and vital
force in nature in the non-covalent bonding.
And those are literally interactions
amongst molecules that have a lot of CH and CC bonds.
The final force that's shown up there
is the van der Waals force.
And we don't worry too much about that,
but it is simply the interaction between very weakly polarized
carbon-hydrogen or other types of bonds
where there's a little bit of a dipole between the bond
and they form little dipolar interactions.
But mostly, I think you really want
to focus on the electrostatic, the hydrogen
bond, and the hydrophobic.
These are more minor and it's a little bit of a subtlety.
So let's focus on those three.
All right, so with that said, the key thing for you--
what do you need to be able to do
is understand them and recognize them in complex systems.
Lastly I'm just going to leave this.
It's going to stay in your notes.
We in biochemistry tend to use line angle drawings.
It's kind of complicated to draw these sort of great big things
with all the hydrogens and oxygen and stuff spelled out,
so we use the line angled drawing.
There's some shown here for different molecules.
And the rules are laid out so that you can go and just
figure out, do a bit of practicing,
and just figure out the line angle drawing
and what it means.
Basically, every line represents a bond,
every vertex represents a carbon atom.
But what you do show on the drawings
are the non-carbon atoms.
So for example, oxygen, or nitrogen. And when you show,
you imply the hydrogens that are bonded to carbon
but you have to show the hydrogens that
are on nitrogen or oxygen, for example,
and you have to figure out what your charged state might be.
So I'm going to leave you with that.
All right.
OK.
So what we've learned so far is these basic forces in biology
are critical for the assembly of the building blocks
of biological macromolecules.
What I want to talk to you about now-- and we'll probably,
because I've spent a little bit of time on that,
spill over a little more to next week--
but I'm going to talk to you about the first class
of macromolecules, which are the lipids.
So what makes something a lipid?
These are the most sort of complicated mixture
of biological molecules.
And formally, they're not really macromolecules.
They're small molecules.
But what's common to all of them is that they are very rich
in carbon-carbon and carbon-hydrogen bonds
because all of these--
the line angle drawings of all of these
would suggest to you that the dominant feature of all
these molecules is a bunch of CC and CH
ions, which makes the molecules quite hydrophobic.
There are no functional groups there.
And they behave very differently.
For example, they would have a tough time dissolving
in water in some cases.
And so this complicated looking set of molecules
can be distilled out as being very rich
in carbon-hydrogen and carbon-carbon bonds.
And we call those collectively lipids.
And they have a lot of different functions.
So for example, triglycerides, such as shown here,
with three ester bonds are storage for energy--
things like estradiol, things like steroids.
They have this 6-6-6-5 arrangement of rings.
All your steroid hormones kind of look like that.
A lot of CH bonds.
There are some vitamins.
So for example, retinol is a vitamin.
It's also a lipid.
And then there are the phospholipids shown down here.
I just briefly want to mention a little bit about retinal
and retinol, which are crucial.
Retinol is a critical vitamin.
It comes actually from carotene, which
is a molecule that you find in a lot
of orange and yellow fruits, such as carrots.
But the oxidized product of retinol
is this lipid called retinal, which
is central to the process of vision.
So retinal binds to proteins that sit in the membrane.
When light shines on them, the shape of the retinal changes.
It goes from a particular configuration
of the double bond to a different one.
The shape just changes, and that sends a signal to your brain.
So lipids are important, absolutely essential,
in vision and sight because they are involved
in the signaling process because their shapes change and send
signals.
Other types of lipids--
so these things-- and we call them fatty acids mostly
because they are greasy long-chained acids
with a long hydrophobic tail and a hydrophilic end group here.
These molecules are also what are
known as amphipathic because they have
a sort of split personality.
They have a hydrophobic moiety and a hydrophilic moiety.
Whenever you see amphi at the beginning of a word,
it means in both.
So both hydrophilic and lipophilic.
So these are important.
And these are very important components.
You probably heard a lot of press
about some of the fatty acids and how bad trans fats
are for you and how you should be careful to make sure
your diet is rich in cis fats rather than trans fats
because the trans fats are contributors to coronary heart
disease.
So you may wonder, what's the relationship between heart
disease and these two types of lipophilic components
which are in the body?
So let me describe to you that relationship.
Remember that cis fats are rich in things like the nut oils
and fruit oils, such as olive oil.
So coronary heart disease is associated with trans fats.
What's the linkage, what's the biology in that?
So the story is related to cholesterol.
Cholesterol is a critical component in our membranes.
The trouble is we have to be able to move cholesterol
around.
But it's so hydrophobic it doesn't dissolve in water, OK?
So in the body, your cholesterol is moved around
in the form of lipoproteins that bind to the cholesterol
and take it to the different organs
where it is needed, all right?
And so the lipoproteins can either
be low density and associate with cholesterol,
or they can be high density, and those also
associate with cholesterol.
The high density lipoproteins are kind of large.
In fact, they're fairly agile.
They don't stick to arteries and vessels,
and they can be excreted in the liver
or move around the bloodstream without any problem.
It's the low density ones that are problems because they're
low density and they kind of stick
to the walls of your arteries and start
making buildups and then plaques, which
contribute to heart disease.
So the low density ones have cholesterol,
but they're very small, sticky, and it's a physical interaction
with your blood vessels and they start to clog your arteries.
What's the relationship to saturated and trans fats?
It's that they increase the low density
lipoprotein in preference to the high density.
So if you have a lot of trans fats,
you make a lot of low density lipoproteins,
it's trying to carry cholesterol around,
but it gets stuck to your blood vessels
and you start to clog your blood vessels.
That contributes to heart disease.
So these lipophilic molecules are important.
They are places to store energy.
They are critical to hormones and signaling, for example.
But there are some complications with disease
because certain types of fatty acids
contribute to heart disease.
Yeah.
AUDIENCE: Is it a lower density if it
doesn't have a bend in it?
PROFESSOR: No, no.
the density is of the entire physical particle.
It's a nanoparticle that would show a different density
respective to how it floats in water.
So the density is really the physical metric
of the entire particle as opposed to just the molecule.
It might be different because of the way it compacts,
but the important thing about the trans fats
is that they really contribute to making
the protein that forms the low density particles.
OK, all right.
So I'm just going to introduce these--
not quickly, but I'll show you some cool images
at the beginning of the next class.
This is the last group of lipidic molecules,
and they are actually--
whoops-- esters and phosphoesters
of fatty acids with glycerol.
This is a small molecule that forms
esters through its oxygen to these long chains
and also to phosphate.
And these contribute to really important functions
in the body.
They are also amphipathic because they
have a hydrophobic component and a hydrophilic component.
And we often draw them in a shorthand form like this
to represent this head group and these tails.
And I want to just leave you with this wonderful image
of the sorts of supramolecular structures
that these kinds of phospholipids can form.
So supramolecular is a very important term in biology
as it is in engineering-- supramolecular.
It means it's a structure that's above the molecular level.
It's an aggregation of different molecules
to make a super molecule with different properties
from the individual components.
Phospholipids self-assemble-- and that's another important
term--
into supramolecular structures that are very, very important
in living systems.
Some of them just are useful in other sorts of engineering
approaches, such as liposomes and micelles,
but the most important supramolecular structure
of a phospholipid is the lipid bilayer
that surrounds your cells.
And what happens is you simply put those molecules--
the phospholipids in water and they
will self-assemble on their own into these supramolecular
structures.
Whether they form micelles or liposomes or bilayers
is dependent very much on the tails of the lipids--
what sorts of shapes and structures you get.
But in physiology-- in human physiology--
the phospholipids that we have want
to form these bilayer structures that have incredibly
important properties.
Most importantly that they are semi-permeable
and they can wrap, form the boundary to certain cells.
So I will continue with the final discussion
of this on Monday before we move forward
to the amino acids, peptides, and proteins.
And I just quickly want to move you to ask you
for Monday to try to catch a read of the section 3.2
in the text if you have a chance.
It'll give you a nice preview.