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Hi, in this video, we will discuss some of the ways that atoms and
molecules bond and bind to each other.
These topics will be review to many of you, it'll be new to some, but it's
important that we're all get on the same page, so we can have more
discussions.
We'll talk about covalent and ionic bonds, and we'll talk about hydrogen
bonding, which is a particular type of electrostatic interaction.
We'll talk about hydrophobic effects, which is one way that molecules will
separate and organize.
And we'll sort of drive it home with two rules of thumb, which we'll talk
about over and over in this course.
One is about electrostatics, opposite charges attract and
like charges repel.
And the hydrophobic effect, which is basically saying that oil
and water don't mix.
To start the discussion, we want to talk about how atoms can be held
together, either by covalent bonds, or ionic bonds.
In the first case they share electrons, in the second case they
will transfer electrons to each other.
In this slide we have very simple schematics of covalent bonding on the
left, and ionic bonding on the right.
The trick is, how do you get to positively charged nuclei to be close
to each other?
Again, because like charges repel.
In the sharing of electrons, you have an excess of electron density in
between the two positive charges.
So whereas before, you have a spherical distribution, now you will
have a distribution of the electrons which are concentrated between the two
positively charged nuclei.
And that's a stable electrostatic configuration.
And that is what we think of as a covalent bond.
In the other case, you have the transfer of an electron from one atom
to another.
Now you have a positively charged atom interacting with the negatively
charged atom.
And again, this is an ionic bond.
This is an attraction between the two, and they will be in close proximity.
Now a prototypical ionic bond is one seen between sodium and chlorine.
Sodium atoms will donate one of their electrons to chlorine.
This gives both ions complete outer valence shells-- eight electrons in
the valence shell-- which is a stable atomic configuration.
And now sodium and chloride ions can interact with each other.
They can, in the absence of solvent, they can form regular grids, which are
visible here as salt crystals, when being evaporated.
On this slide, we'll talk about some of the strengths of interactions of
molecules and energy content in general.
We'll start by pointing out that the carbon-carbon bond, a single bond, is
about 83 kilocalories at room temperature.
And compare that to the average thermal energy available at room
temperature, kT, which is about 0.6 kcals.
So you can see that at room temperature it's extremely unlikely
that a carbon-carbon bond would ever break, due to just
average thermal motion.
Non-covalent bonds in water, like hydrogen bonds, can be from one to a
few-- two or three-- kilocalories.
And average thermal motions are large enough to disrupt those on occasion.
When we hydrolyze a molecule of ATP in the cell, we will get about 11 to 12
kilocalories per mole from that, verses approximately 673 kilocalories
for the complete oxidation of glucose.
On this slide we see the elemental composition of human bodies in red,
and that we are comprised primarily of four atoms, hydrogen carbon, oxygen,
and nitrogen, with many other atoms contributing small amounts.
With covalent bonding there's some different properties between single
and double bonds.
Single bonds, you're free to rotate around.
Double bonds are stronger, but they are constrained.
You will not get rotation around this double bond, and these atoms will all
be confined in a plane.
We are primarily carbon-based creatures.
As you can see, our hydrocarbon, carbon, is tetravalent.
It's, in this case, bonded to two other carbons and two hydrogens.
This can sometimes be drawn like this, without explicitly
drawing in the hydrogens.
Or like this, where every vertex is a carbon atom, and it's always implied
that there are as many hydrogens as there need to be to make each atom
tetravalent.
Aromatic compounds, such as benzene, it's more than double bonds when you
have four N plus 2 electrons in a cyclic structure that's aromatic.
There's extra resonant, extra thermodynamic energy there.
And the true benzene structure is a resonance between these two
structures.
And it's often written as a circle.
We have certain words for different groups.
As you can see here, it's showing a methyl group.
Here is a methyl group.
If we talked about these two carbons, this would be an ethyl group.
If we talked about these three carbons and all the hydrogens, that would be a
propyl group.
So the different nomenclature is saying something about the number of
carbons in this case.
In addition to hydrocarbons, you can also have oxygen.
In this case, the hydroxyl group, which is an alcohol.
A slightly more oxidized version of this carbon is a carboxylic acid,
which in solution is a negatively charged molecule, as it will lose a
hydrogen ion to water.
Some covalent bonds can have partial ionic character.
That is the case with water.
There's covalent bonds here between oxygen and hydrogen, but since oxygen
is more electronegative than hydrogen, it will attract more of the electron
density closer to its nucleus.
And therefore, there'll be a buildup of negative charge.
You'll have a partial negative charge of oxygen on the oxygen, and
corresponding partial positive charges on the hydrogen.
This will result in a net dipole moment for this molecule.
Contrast that to a case of oxygen, where the atoms are identical and
therefore completely equivalent electron sharing.
There'll be no partial positive and negative charges on this molecule.
Now we'd like to discuss a type of electrostatic interaction that occurs
so frequently in biology that it's referred to as a hydrogen bond.
It's really nothing other than straight electrostatics.
You have electrostatic interaction between partial negative charges and
partial positive charges, and those attract.
They want to be as close to each other as possible.
You have repulsion of partially negative charges with partial negative
charges, which is what gives this stable configuration of atoms sort of
a linear geometry.
If you were to move this at an angle, so you make the hydrogen oxygen bond
here, now you'd be bringing these partial negative charges closer to
this partial negative charge, and that is energetically unfavorable.
So you will get a linear confirmation.
The distances between the hydrogen, one Angstrom roughly, for the
covalently bonded portion, and 1.7 angstroms between the oxygen and
hydrogen that are participating in the covalent bond.
Now the oxygen has two lone pairs, which can participate in hydrogen
bonding, and two hydrogens.
So each water molecule has the ability to take part in four
hydrogen bonding events.
And that is indeed what happens.
Most the time it's transient in liquid water because of thermal motions being
of comparable energy, and able to disrupt it.
But when you freeze it down into ice, thermal energy is down, and you sort
of freeze-in this very organized structure between them.
And this is a contributing factor to why ice is less dense, and rises to
the top of lakes for instance, than liquid water.
Now when you dissolve a hydrophilic atom or molecule into water, it has
the ability to be solvated.
A positively charged sodium ion will attract the partially negative charges
on the oxygen, forming a configuration.
If you put in a negatively charged chloride atom, this attracts the
partial positive hydrogens, and so you can say that these ions are solvated.
They're stabilized.
It's a fine situation.
In this case, you have an organic molecule with the ability to form lots
of hydrogen bonds.
And these water molecules can exchange with one another quickly, and so this
is an energetically favorable situation.
Now in the case of hydrophobic molecules, hydrophobic molecules are
not able to engage in hydrogen bonds.
And I should have mentioned, when I introduced hydrogen bonds, that only
hydrogens that are connected to electronegative atoms--
like oxygen or nitrogen--
can participate in hydrogen bonding.
Hydrogens bonded to carbons cannot participate in hydrogen bonding,
because there is no great electronegativity difference between
the hydrogen atom and the carbon atom.
And so you won't have the partial positive and negative
charges on the two atoms.
And you won't have then, the ability to form these electrostatic
interactions.
Since hydrophobic molecules are not able to participate in hydrogen bonds,
what you have is water molecules nearby which are not making bonds,
which they want.
And you're also not able to rotate around freely.
So this is an unfavorable energetic interface, and one that
thermodynamically will be minimized.
And we see an example of that here.
We have two hydrophobic methyl groups, which have energetically unfavorable
interfaces with a large amount of solvent.
By clustering these two methyl groups together, you're minimizing, in
essence, the surface area of these molecules.
And so you're minimizing the energy of the system, although it's primarily an
entropic consideration, because here you're
restraining many water molecules.
Whereas in here, there are fewer water molecules that have their
confirmations restrained.
Hydrogen bonding plays huge roles in the shape of
proteins and nucleic acids.
In this case, we have hydrogen bonding between backbone atoms of perhaps an
alpha helix.
We'll see examples of this in future lectures.
And of two base pairs, nucleic acid base pairs, a C and a G, interacting
with each other through a number of hydrogen bonds.
So when proteins bind to each other with their convoluted shapes, what
they're doing are, they have complementary
shapes and charge patterns.
Here you can see electrostatic interactions between the two.
There are also regions of the protein that are neither positively nor
negatively charged, complementary on the other binding partner, and these
can engage in hydrophobic patches or interactions.
And those will contribute to the interactions of
biomolecules with each other.
So this was a very quick run through different types of bonding
considerations.
But again, I want to highlight the two rules of thumb which we'll use over
and over in this course.
And that is that opposite charges attract, like charges repel, and this
is electrostatic interaction.
And that's an enthalpic contribution to the free energy.
The other rule being, that we can summarize that oil
and water don't mix.
This is what gives rise to hydrophobic forces.
And these are primarily entropic contributions to energy stability.