Covalent network solids | Intermolecular forces and properties | AP Chemistry | Khan Academy

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 seen thought about them as these positive ions in this sea of negatively charged electrons. 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 there are 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.

Now, some of you might think, "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. Each of these sheets actually are attracted to 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 36 degrees Celsius. The silicon carbide that we looked at at the beginning of this video decomposes at 2830 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 1722 degrees Celsius.

So, the big takeaway over the last several videos is that there are 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.