- [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.
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.