- [Instructor] So we've already talked
about multiple types of solids.
We've talked about ionic solids.
That's formed when you have ions
that are attracted to each other,
and they form these lattice structures.
We have seen metallic solids,
and we've thought about them as these positive ions
in this sea of negatively charged electrons.
And we've also seen molecular solids,
which is formed from individual molecules being attracted
to each other through intermolecular forces.
Now, what's different about covalent network solids is
that they're entire networks formed by covalent bonds.
What we see here, for example, is a network
of silicons and carbons,
and this is silicon carbide right over here.
And now, some of you might thinking,
haven't we already seen covalent bonds involved
in a solid before, for example, in molecular solids?
And this right over here is an example
of a molecular solid that we studied in that video.
You have the molecules,
which are made up of atoms bonded with covalent bonds.
But the reason why they form a solid is
because the molecules are attracted to each other
through intermolecular forces.
And if you wanted to melt this molecular solid,
you have to essentially overcome
these intermolecular forces.
Well, in a covalent network solid, the solid,
to a large degree, is made up of these covalent bonds.
And if you wanted to melt this somehow,
you would have to overcome these covalent bonds,
which, generally speaking, are stronger
than these intermolecular forces.
And so you can imagine, covalent network solids
are going to have higher melting points.
You also don't see a sea of electrons here.
So unlike metallic solids,
they're not going to be good conductors of electricity.
But just to understand this point a little bit more clearly,
let's look at some more covalent network solids.
So what you see here on the left,
you might recognize as a diamond.
A diamond is just a bunch
of carbons covalently bonded to each other,
and this is the structure of how these carbons are bonded.
And as you might already know,
diamonds are the hardest solid that we know of.
These covalent bonds, the way that they are structured,
can take a lot of stress, a lot of pushing and pulling.
It's very hard to break it.
Now, what's interesting is that same carbon can form
different types of covalent network solids.
For example,
this right over here is graphite,
and graphite is probably something
you're quite familiar with.
When you write with a pencil,
you're essentially scraping graphite
onto that piece of paper.
And so this is what graphite looks like.
It's these covalent network sheets,
and each of these sheets actually are attracted each other
through intermolecular forces.
And that's why it's easy to scrape it,
because these sheets can slide past each other.
But if you really wanted to melt graphite,
you would have to break these covalent bonds.
And so you can imagine, to overcome the covalent bonds
and melt, say, diamond or graphite,
it takes a very, very high temperature.
Graphite, for example, sublimes
at 3,642 degrees Celsius.
The silicon carbide that we looked at
at the beginning of this video,
it decomposes at 2,830 degrees Celsius.
This right over here is a picture of quartz,
which is a very common form of silicon dioxide,
another covalent network solid,
and this has a melting point
of 1,722 degrees Celsius.
So the big takeaway over the last several videos is
there's many different ways of forming a solid.
It could be with ions, it could be with metals,
it could be with molecules that are attracted
to each other with intermolecular forces,
or you could have a network
of atoms formed with covalent bonds.