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