- [Instructor] Let's talk
about hydrogen bonds.
Depicted here, I have three
different types of molecules.
On the left, I have ammonia.
Each ammonia molecule
has one nitrogen bonded
to three hydrogens.
In the middle,
I have something you're
probably very familiar with,
in fact, you're made up
of it, which is water.
Each oxygen is bonded to two hydrogens.
And then here on the right,
I have hydrogen fluoride.
Each fluorine is bonded to one hydrogen.
Now, why are these types
of molecules interesting?
And what does that have
to do with hydrogen bonds?
And the simple answer is,
in each of these cases,
you have hydrogen bonded
to a much more electronegative atom.
Even though these are covalent bonds,
they're going to be polar covalent bonds.
You are going to have a
bond dipole moment that goes
from the hydrogen to the
more electronegative atom,
from the hydrogen to the
more electronegative atom,
from the hydrogen to the
more electronegative atom.
The more electronegative atom
is going to hog the electrons.
The electrons are gonna
spend more time around that.
So that end of the molecule is going
to have a partial negative charge.
And then the ends with the hydrogens,
those are gonna have
partial positive charges.
Another way to think about it is,
if you added these dipole moments,
you would have a net dipole
for the entire molecule
that would look something like that.
So we are dealing with polar molecules.
And the polarity comes
from both the asymmetry,
and you have a very electronegative
atom bonded to hydrogen,
oxygen, very electronegative
atom, bonded to hydrogen.
So this end of the molecule
is partially negative.
This end of the molecule
or these ends of the molecule
are partially positive.
For hydrogen fluoride, this
end is partially positive.
This end is partially negative.
And so what do you think could happen
when these molecules
interact with each other?
The nitrogen end right
over here, of this ammonia,
could be attracted to
one of these hydrogens
that has a partially positive
charge right over there.
Or this hydrogen, the
partial positive charge,
might be attracted to that nitrogen
that has a partial negative charge.
And this attraction
between the partial positive hydrogen end
and the partially negative
end of another molecule,
those are hydrogen bonds.
And they are an intermolecular
force that will be additive
to the total intermolecular force
from, say, things like
London dispersion forces,
which makes you have
a higher boiling point
than you would have if you just thought
about London dispersion forces.
And to make that clear,
you can look at this chart.
You can see all of these
molecules are formed
between period two elements and hydrogen.
In fact, all of these molecules
have similar molar masses,
methane, ammonia, hydrogen
fluoride, and water.
If we were just thinking about
London dispersion forces,
London dispersion forces are proportional
to the polarizability of a molecule,
which is proportional to
the electron cloud size,
which is proportional to the molar mass.
And generally speaking, as
you go from molecules formed
with period two elements
to period three elements
to period four elements
to period five elements,
you do see that as the molar mass
of those molecules increase,
there is that general upward
trend of the boiling point,
and that's due to the
London dispersion forces.
But for any given period,
you do see the separation.
And in particular,
you see a lot of separation
for the molecules formed
with oxygen, fluorine, and nitrogen.
These molecules, despite
having similar molar masses,
have very different boiling points.
So there must be some other
type of intermolecular forces
at play above and beyond
London dispersion forces.
And the simple answer is yes.
What you have at play
are the hydrogen bonds.
Now, some of you might be wondering,
well, look at these molecules formed
with period three elements and hydrogen
or period four elements and hydrogen,
they also don't have
the same boiling point,
even though you would expect
similar London dispersion forces
because they have similar molar masses.
And the separation that you
see here in boiling points,
this, too, would be due to other things,
other than London dispersion forces.
In particular, dipole-dipole
forces would be at play.
But what you can see is
the spread is much higher
for these molecules formed
with nitrogen and hydrogen,
fluorine and hydrogen,
and oxygen and hydrogen.
And that's because hydrogen
bonds can be viewed
as the strongest form
of dipole-dipole forces.
Hydrogen bonds are a special
case of dipole-dipole forces.
When we're talking about hydrogen bonds,
we're usually talking about
a specific bond dipole,
the bond between hydrogen and
a more electronegative atom
like nitrogen, oxygen, and fluorine.
And so we're specifically talking
about that part of the molecule,
that hydrogen part that has
a partially positive charge
being attracted to the
partially negative end
of another molecule.
So it's really about a bond
dipole with hydrogen bonds
versus a total molecular dipole
when we talk about dipole-dipole
interactions in general.
And so you could imagine,
it doesn't even just
have to be hydrogen bonds
between a like molecule.
You could have hydrogen bonds
between an ammonia molecule
and a water molecule or
between a water molecule
and a hydrogen fluoride molecule.
And I mentioned that these are
really important in biology.
This right over here is a closeup of DNA.
You can see that the base pairs in DNA,
you can imagine the rungs of the ladder,
those are formed by hydrogen
bonds between base pairs.
So those hydrogen bonds are strong enough
to keep that double helix together,
but then they're not so strong
that they can't be pulled apart
when it's time to replicate
or transcribe the DNA.
Hydrogen bonds are also
a big deal in proteins.
You learn in biology class
that proteins are made up
of chains of amino acids,
and the function is heavily influenced
by the shape of that protein.
And that shape is
influenced by hydrogen bonds
that might form between the amino acids
that make up the protein.
So hydrogen bonds are everywhere.
There are many hydrogen bonds
in your body right now mainly,
not just because of the DNA,
mainly because you're mostly water.
So life, as we know it, would not exist
without hydrogen bonds.
