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.
