PROFESSOR: Here is the game
plan for the next
two or three lectures.
I'm going to start by talking
about the chemical forces that
are important for the structure
and function of
these biomolecules.
And then I'm going to relate
them, as we go along, to how
these properties influence the
characteristics of these key
macromolecules.
And in particular we'll be
talking about covalent bonds,
hydrogen bonds, ionic bonds, a
force known as van der Waals
forces, and something that's not
really a force but it's a
characteristic that's very
important, particularly when
we think about proteins
and lipids, called
hydrophobicity--
literally "fear of water."
And then the order
of the molecules.
As we talk, I'll talk about
carbohydrates first.
I'll try and get
to that today.
Then we'll talk about proteins,
nucleic acids, and
lipids, in that order.
As you'll see, these two will
be sufficient to understand
most of the characteristics
of carbohydrates.
Whereas we're going to need all
five of these to be able
to get an intelligent
understanding of
how proteins work.
Now, I'll caution you.
It's going to seem--
God, he's going to talk about
covalent bonds, as everybody
is rolling their eyes.
I heard about covalent bonds
in grade one or something.
But the difference here is that
we're going to be looking
at some of these forces, some
of which you've been exposed
to already, but from a
biological perspective.
And I hope if you kind of watch
that, you'll begin to
see that you're looking at
something that may be sort of
familiar to you.
But you have to start thinking
about it in a different way
once you start thinking of what
are the implications of
the properties of these forces
and the way these molecules
behave for biology.
So, begin with the one that
everybody undoubtedly knows,
which are covalent bonds.
And this is the principal force
that holds atoms together.
And it's based on sharing
electrons.
And as I'll say, these are
very strong bonds.
And so in the simplest sort of
example, hydrogen atom has one
unpaired electron, a
carbon has four.
And so you can make
methane, CH4.
And commonly in chemistry and
biology we use a line to
represent a pair of electrons.
So there's methane.
As I said, apart from you know
it burns, if you go out in a
swamp or in a beach and you
see bubbles, muddy bottom
coming up, those are bubbles of
methane made by methanogens
that are living in the anaerobic
layer underneath.
A cow has a special
fermentation,
digestion cavity inside.
It's huge, called a rumen,
stuff sloshing around.
And it's full of archaea,
that are methanogens.
And a cow makes about 400
liters of methane a day.
And Penny will tell you, it's
a very bad greenhouse gas.
It's much more potent
than carbon dioxide.
And so the typical length of a
covalent bond is about 1.5 to
0.2 nanometers.
And I hope you'll try and begin
to get a sense of the
links of some of these
things, too.
But the key point about this
is to break a carbon-carbon
bond needs 83 kilocalories
per mole.
So that's a lot of energy.
At 25 degrees centigrade, if
you take, say, a typical
vibrational mode of a covalent
bond, the energy that it has
is about 0.6 kcals per mole.
So what that means is that
covalent bonds don't break on
their own under physiological
conditions.
They can bend, they can rotate,
and they can stretch.
So they're back and forth this
way, they can go this way,
this way, but they
don't break.
And so this sort of leads to
another topic that we'll talk
about, which is utterly key--
It's one of the secrets
of how life works--
are these protein molecules
that are known as enzymes.
And we'll also talk a little bit
about a similar thing made
of RNA called a ribozyme.
But what these are are
biological catalysts that
enable specific bonds-- and
this is important--
specific bonds to be broken or
formed under physiological
conditions.
And this part is so important.
If you're trying to work out
a chemical reaction, the
original process for taking
nitrogen gas and making
ammonia, the Haber process,
involved some very, very tough
molecule to break the bond of.
So just heat it up to 500
degrees and put in a catalyst.
But if you're a living
organism, you
don't have that option.
You have to continually make
and break bonds under the
conditions -- the very, very
narrow conditions where life
is possible.
If you go a little too high,
things like proteins unfold.
And then they don't work as
properly as machines anymore.
So we'll be talking more about
that as we go along.
There are different types
of covalent bonds.
And again, the first
part of this isn't
going to surprise you.
There are single bonds,
like this.
There are double bonds,
and triple.
Excuse me.
I'll just stay with carbon
for the moment.
The more electrons that are
shared, the stronger the bond.
And these two are referred to,
if it's a carbon compound, as
being unsaturated bonds, the
same term you hear when you
hear about unsaturated fats.
And what that means is a fat
with an unsaturation, that's
unsaturated, will have somewhere
in it a double bond,
or in some cases, many
double bonds.
However, there's another aspect
of this which might not
have been relevant to you, but
you'll see it becomes relevant
for thinking about proteins as
soon as the next lecture.
And that is, a single bond is
able to rotate this way.
These guys can't rotate.
And that, as you'll see, becomes
important in quite a
variety of situations.
But we'll run into a very
important example of that when
we're thinking about the very
backbone of all proteins, the
peptide bond, which is at the
heart of being a protein.
There are other molecules that
have more than one bond that
are important.
Oxygen is one.
And nitrogen, as I said, is a
particularly hard nut to crack.
Most organisms, as I said the
other day, are unable
to break this bond.
The only organisms that
have learned how
to do it are bacteria.
The vast majority of them use
one, single enzyme called
nitrogenase that evolved that's
a very complicated
enzyme and has very, very
stringent requirements and
needs a huge energy input.
But it is able to crack
this bond and get
it made into ammonia.
But it's an example of another
molecule that has a triple
bond in it.
Let's see, how are
we doing here?
Okay.
So another aspect of these
covalent bonds that you need
to think about has to
do with when you're
thinking about carbon.
And it's a property
called chirality.
And it comes from the fact that
carbon has four bonds but
they come out as
a tetrahedron.
So that doesn't matter in
the case of methane.
But I'm going to depict the
tetrahedron in this way, so
that this bond is coming --
these two are in the plane of
the board, this one's coming
out, that one's going back.
And let's just put on four
different substituents.
Now if I get the mirror image
of that, we will have--
these two molecules are called
optical isomers.
And if you sit down and play
with this, you will find you
can't convert one to the other
without actually physically
breaking a bond.
And this is really important,
one of the central concepts
that I hope you might remember
from this course because it
cuts across a lot of the stuff
talk we'll be talking about.
At a molecular level, much
of biology relies on the
interaction of complementary
3D surfaces.
We're actually very familiar
with this at a macro level in
our own lives.
Imagine you've just come back
from the party late on
Saturday night, you're
crossing the Mass.
Ave. Bridge, the wind is
howling, you're freezing.
But no problem, you've
got your gloves.
And you reach in your pocket and
you have two left gloves.
No matter what you do, you can't
get that right hand to
fit properly into the
left-handed glove.
One's a mirror image
of the other.
But we run into this problem
even in our own lives.
When you saw how that DNA
had fit right into a
groove in the protein.
If we had a mirror image of
the DNA or we had a mirror
image of the protein,
it wouldn't work.
This principle goes all the
way through biology.
There is another characteristic
of covalent
bonds that becomes
important again.
And that is how equally the
electrons are sharing.
So again, it goes back
to the sharing of
electrons, but with a twist.
If we have a carbon-carbon or
a carbon-hydrogen bond, it's
pretty much equal sharing.
And this is known as
a nonpolar bond.
But if you have a nitrogen or an
oxygen bond, it's unequal.
And these are known
as polar bonds.
And the term that's used to
describe this unequal sharing
of electrons is known as the
electronegativity of the atom.
It's basically a word that
means the greediness of a
particular atom for electrons.
So if you have an oxygen and a
hydrogen bond, although we
write it like that on the board
and you've undoubtedly
seen this for many years in
chemistry, in fact, the
electrons spend more time down
here than they spend up there.
So there's a little bit of a
negative charge on the oxygen
and a little bit of a plus
charge on the hydrogen.
That's usually represented by
a little delta to indicate
that this has a wee bit of
negative charge, that has a
wee bit of positive charge.
And a molecule that's
very important with
respect to this is water.
Because water, as you
know, is H2O.
But it's not symmetrical.
The angle here is
104.5 degrees.
And so the oxygen has a little
bit of a negative charge but
each of these has a little
bit of a plus charge.
Actually, water is 55 molar.
So it's a little dipole.
You've got 55 molar, these
little dipoles going on.
This property of
electronegativity and nonpolar
bonds then leads to the second
of the forces that we're going
to be talking about.
That's force number two.
And that's a hydrogen,
or H bond.
And this is a bond that's made
possible by a little bit of a
negative charge that's on
oxygen, or nitrogen, or a few
other molecules and a little
bit of a positive charge
that's due to the hydrogen
that's in a polar bond.
This is very important, as
you'll see, for proteins,
nucleic acids, and for
carbohydrates.
And it has a huge amount
to do with the
way that water behaves.
Because in that 55-molar water,
you'll have one water
molecule that's going
to be like this.
And there will be another water
molecule down here with
a little bit of a
negative charge.
And this a little bit of a plus
charge on this hydrogen
and a little bit of a negative
charge can form what's known
as a hydrogen bond
between them.
And what's especially important
about these hydrogen
bonds is they're about
1/20 the strength
of a covalent bond.
And that means that in a
distribution of molecules at
physiological temperatures,
there will be some guys up in
the -- the most energetic
molecules within the bunch
will have enough energy to
break hydrogen bonds.
But they're much easier to do.
And just to peer ahead, when we
talk about replicating DNA,
those two strands are held
together by hydrogen bonds.
So the backbones are really
solid, just like two strips of
Velcro or something.
But the hydrogen bonds hold the
two strands together, but
1/20 the strength.
So it's basically like molecular
Velcro between the
two strands of DNA.
And we'll see some more
examples of this.
Let's see if I can go
back to this and get
this thing to play.
This is static representation
just illustrating this.
But in fact, what happens,
water molecules are
continually changing partners.
So they're constantly making
shells, and cages, and so on.
And the next little movie is a
picosecond simulation of water
just at zero degrees.
And you can see how the
molecules are changing
partners, making little
shells and things.
And here's a picosecond
simulation of water at the
boiling temperature.
And what you can see from this
is every now and then, a
molecule like this one will get
enough energy to break out
of this constant sharing
of little
hydrogen bonds and escape.
And another thing, when we talk
about getting something
dissolved in water, this
is something we'll
have to think about.
Because if you try and dissolve
something in water,
like stir a lot of oil into
it, you know what happens.
You can stir like mad
and doesn't go in.
Part of the problem is if you
put something in the water,
it's going to have to break
these existing hydrogen bonds.
And that's an energy cost.
So in order to get something to
dissolve, you're going to
have to get the energy back.
And we'll be talking
about that.
But it's one of the fundamental
parts of water.
You're familiar with the
characteristics of water.
There's surface tension.
It's why trees can grow 300 feet
tall, because they've got
basically little nanotubes
and little capillaries.
And with this surface tension,
water, due to hydrogen bonds,
can go 300 feet up.
The water can go right up.
You've seen water bugs
walk around on water.
There's a particularly
interesting lizard in South
America, Central America called
the basilisk lizard
that's about 2 and
1/2 feet long.
It's able to run across
the top of the water.
It's actually called the
Jesus Christ lizard.
And it's able to do that
because of this surface
tension in the water.
In fact, when I finished my
Ph.D. Thesis, I went in a
competition for the theses.
And mine was something like, a
chemical enzymatic synthesis
of oligoribonucleotides.
And I was competing against a
guy who said why do lizards
run on water, and his entire
talk consisted of movies of
this thing running
across the water.
I thought I was toast, but I
actually won that prize.
But anyway, every time I
see this I remember it.
For example, when they go and
explore Mars or think about
planets, they're always looking
for water because it
has this very, very special set
of properties that are so
important for life.
