Covalent networks, metallic, and ionic crystals Video Lecture - Class 11

In the last video, I talked
about some of the weaker
intermolecular forces or
structures of elements.
The weakest, of course, was the
London dispersion force.
In this video, I'll start with
the strongest structure, and
that's the covalent network.
So if you have a covalent
network crystal-- and let me
actually define the
word crystal.
Crystal is just when you have
a solid, where the molecules
that make up the solid are
in a regular, relatively
consistent pattern, and this is
versus an amorphous solid,
where everything is kind of just
a hodge-podge and there's
different concentrations of
different things, of different
ions, and different molecules,
and different
parts of the solid.
So crystal is just a very
regular structure.
Ice is a crystal, because once
you get the temperature low
enough in water, the hydrogen
bonds form a crystal, a
regular structure.
And we've talked about
that a bunch.
But the strongest of all crystal
structures is the
covalent network.
And the biggest, or the prime,
example of that is carbon when
it forms a diamond.
So in the covalent network,
carbon has four valence
electrons, so it always
wants four more.
So when carbon shares with
itself, it's very happy.
So what it can do is it can form
four bonds to four more
carbons, and then each of those
carbons can form four
more bonds to four
more carbons.
And this, one, 1, 2, 3, and
it just keeps going on.
This is the structure
of a diamond.
And the reason why this is such
a strong structure is
because you can almost view
the entire-- in fact, you
should view the entire diamond
as one molecule, because they
all have covalent bonds.
These are actual sharing of
electrons, and these are
actually the strongest of
all molecular bonds.
So you can imagine if the entire
solid is made out of
this network of carbons,
you're going to have an
extremely strong, extremely high
boiling point substance,
and that's why a diamond is so
strong, and that's why it's so
hard to boil a diamond.
Now, the next two, and it
depends on your special cases
of the next most solid version
of a solid, and it depends
which case you're talking
about, one are the ionic
crystals, and I'll do them both
here, because one isn't
necessarily-- ionic crystal--
and the next is the metal.
Well, it's not the next.
They're kind of the
metallic crystal.
And these bonds, I mean, let's
say the most common ionic
molecule or-- that's not
exactly the right word,
because to some degree, let's
say if I had some sodium and
some chloride-- and just
remember, what happens with
sodium chloride is sodium here
really has one extra electron
that it's dying to lose.
Chlorine has seven electrons
and it's dying
to get a new one.
So sodium essentially donates
its electron to chlorine, and
then the chlorine becomes
negative, the sodium becomes
positive, and they want to be
near each other, right?
So you have a positive sodium
ion and a negative chlorine
ion, and the structure of this
is going to look something
like this, where they're
all-- so let me do
the sodium in green.
So you have a bunch of sodium
ions that are positive, and
then you have a bunch of
chlorine ions that are maybe--
this isn't the exact way that
they actually are, but I think
you get the idea, that one atom
is positive and one atom
is negative, so they really,
really want to be close to
each other.
And so this is a pretty strong
bond, and it has very-- not a
very high boiling point.
It can have a pretty high
boiling point, and this type
of structure is actually
quite brittle.
So if you take some dry table
salt, not dissolved in water,
if you have a big block of it
and you slam it with a hammer,
you'll see that you'll get,
like, a big slice of it.
It'll just fall off, right?
Because you're essentially just
cutting it along one of
these lines really fast. That's
the interesting thing.
Whenever you do something on
a macroscale, like cut
something, you really
fundamentally are breaking
atomic bonds.
So the strength of the atomic
bonds really do tell you about
how hard or strong
something is.
Now, the metallic crystal we've
talked a lot about.
Metals, they like to get rid of
their electrons, or not get
rid of them, they like
to share them.
So what happens is, let's say in
the case of iron, you have
a bunch of iron atoms.
This is all iron.
And their electrons are allowed
to roam free in the
neighborhood.
These are all the electrons.
They're allowed to roam free.
And because of this, it forms
this sea of electrons that are
negative, and that makes it
a very good conductor of
electricity.
And, of course, since the iron
atoms have allowed their
electrons to roam, they all
become slightly positive.
And so they're kind of embedded
in this mesh or this
sea of electrons.
And so the metallic crystals,
depending on what cases you
look at, sometimes they're
harder than the ionic
crystals, sometimes not.
Obviously, we could list a lot
of very hard metals, but we
could list a lot of
very soft metals.
Gold, for example.
If you take a screwdriver and a
hammer, you know, pure gold,
24-carat gold, if you take a
screwdriver and hit it onto
the gold, it'll dent
it, right?
So this one isn't as brittle
as the ionic crystal.
It'll often mold to what
you want to do with it.
It's a little bit softer.
Even if you talk about very hard
metals, they tend to not
be as brittle, because the sea
of electrons kind of gives you
a little give when you're
moving around the metal.
But that's not to say
that it's not hard.
In fact, sometimes that give
that a metal has, or that
ability to bend or flex, is
what actually gives it its
strength because it's allowed to
kind of deflect the force.
So the strength, and I've
touched on this, it also goes
into the boiling point.
So because these bonds are
pretty strong, it has a higher
boiling point.
If you just took salt crystal
and tried to boil it, you'd
have to add a lot of heat
into the system.
So this has a higher boiling
point than say-- I mean,
definitely things that have just
van der Waals forces like
the noble gases, but it'll also
have a higher boiling
point than, say, hydrogen
fluoride.
Hydrogen fluoride, if you
remember from the last video,
just had dipole-dipole forces.
But what's interesting about
this is they have a very high
boiling point unless they're
dissolved in water.
So these are very hard, high
boiling point, but the ionic
crystals can actually be
dissolved in water.
And when they are dissolved
in water, they form
ionic dipole bonds.
What does that mean?
Ionic dipole or ionic
polar bonds.
And this is a situation where
the sodium-- and this is
actually why it dissolves
in water.
Because the water molecule,
we've gone over this tons of
times, it has a negative end,
because oxygen is hoarding the
electrons, and then the hydrogen
ends are positive
because the electron's pretty
stripped of it.
So when you put these sodium and
chloride ions in the room,
or in the water solution, the
positive sodiums want to get
attracted to the negative side
of this dipole, and then the
negative chlorides, Cl minus,
want to go near the hydrogens.
So they kind of get
dissolved in this.
They don't necessarily want to
be-- they still want to be
attracted to each other, but
they're still also attracted
to different sides of the water,
so it allows them to
get dissolved and go with
the flow of the water.
So in this case, when you
actually dissolve an ionic
crystal into water, as an ionic
crystal, not a good
conductor of electricity, not a
lot of charge that is really
movable in this state.
But here, all of a sudden, we
have these charged particles
that can move.
And because they can move, all
of a sudden, when you put
salt, sodium chloride, in
water, that does become
conductive.
So anyway, I wanted you to be
at least exposed to all of
these different forms
of matter.
And now, you should at least get
a sense when you look at
something and you should at
least be able to give a pretty
good guess at how likely it is
to have a high boiling point,
a low boiling point, or
is it strong or not.
And the general way to look at
it is just how strong are the
intermolecular bonds.
Obviously, if the entire
structure is all one molecule,
it's going to be super-duper
strong.
And on the other hand, if you're
just talking about
neon, a bunch of neon molecules,
and all they have
are the London dispersion
forces, this thing's going to
have ultra-weak bonds.
So a gas is almost its
most natural state.
If you get super, super cold,
you might be able to get it to
a fluid, and then everything
in between.
Anyway, hopefully, you
found that useful.