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In the last video, I talked about some of the weaker

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intermolecular forces or structures of elements.

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The weakest, of course, was the London dispersion force.

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In this video, I'll start with the strongest structure, and

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that's the covalent network.

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So if you have a covalent network crystal-- and let me

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actually define the word crystal.

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Crystal is just when you have a solid, where the molecules

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that make up the solid are in a regular, relatively

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consistent pattern, and this is versus an amorphous solid,

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where everything is kind of just a hodge-podge and there's

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different concentrations of different things, of different

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ions, and different molecules, and different

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parts of the solid.

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So crystal is just a very regular structure.

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Ice is a crystal, because once you get the temperature low

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enough in water, the hydrogen bonds form a crystal, a

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regular structure.

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And we've talked about that a bunch.

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But the strongest of all crystal structures is the

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covalent network.

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And the biggest, or the prime, example of that is carbon when

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it forms a diamond.

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So in the covalent network, carbon has four valence

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electrons, so it always wants four more.

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So when carbon shares with itself, it's very happy.

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So what it can do is it can form four bonds to four more

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carbons, and then each of those carbons can form four

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more bonds to four more carbons.

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And this, one, 1, 2, 3, and it just keeps going on.

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This is the structure of a diamond.

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And the reason why this is such a strong structure is

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because you can almost view the entire-- in fact, you

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should view the entire diamond as one molecule, because they

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all have covalent bonds.

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These are actual sharing of electrons, and these are

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actually the strongest of all molecular bonds.

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So you can imagine if the entire solid is made out of

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this network of carbons, you're going to have an

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extremely strong, extremely high boiling point substance,

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and that's why a diamond is so strong, and that's why it's so

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hard to boil a diamond.

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Now, the next two, and it depends on your special cases

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of the next most solid version of a solid, and it depends

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which case you're talking about, one are the ionic

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crystals, and I'll do them both here, because one isn't

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necessarily-- ionic crystal-- and the next is the metal.

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Well, it's not the next.

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They're kind of the metallic crystal.

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And these bonds, I mean, let's say the most common ionic

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molecule or-- that's not exactly the right word,

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because to some degree, let's say if I had some sodium and

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some chloride-- and just remember, what happens with

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sodium chloride is sodium here really has one extra electron

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that it's dying to lose.

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Chlorine has seven electrons and it's dying

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to get a new one.

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So sodium essentially donates its electron to chlorine, and

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then the chlorine becomes negative, the sodium becomes

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positive, and they want to be near each other, right?

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So you have a positive sodium ion and a negative chlorine

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ion, and the structure of this is going to look something

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like this, where they're all-- so let me do

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the sodium in green.

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So you have a bunch of sodium ions that are positive, and

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then you have a bunch of chlorine ions that are maybe--

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this isn't the exact way that they actually are, but I think

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you get the idea, that one atom is positive and one atom

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is negative, so they really, really want to be close to

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each other.

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And so this is a pretty strong bond, and it has very-- not a

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very high boiling point.

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It can have a pretty high boiling point, and this type

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of structure is actually quite brittle.

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So if you take some dry table salt, not dissolved in water,

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if you have a big block of it and you slam it with a hammer,

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you'll see that you'll get, like, a big slice of it.

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It'll just fall off, right?

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Because you're essentially just cutting it along one of

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these lines really fast. That's the interesting thing.

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Whenever you do something on a macroscale, like cut

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something, you really fundamentally are breaking

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atomic bonds.

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So the strength of the atomic bonds really do tell you about

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how hard or strong something is.

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Now, the metallic crystal we've talked a lot about.

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Metals, they like to get rid of their electrons, or not get

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rid of them, they like to share them.

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So what happens is, let's say in the case of iron, you have

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a bunch of iron atoms. This is all iron.

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And their electrons are allowed to roam free in the

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neighborhood.

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These are all the electrons.

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They're allowed to roam free.

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And because of this, it forms this sea of electrons that are

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negative, and that makes it a very good conductor of

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electricity.

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And, of course, since the iron atoms have allowed their

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electrons to roam, they all become slightly positive.

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And so they're kind of embedded in this mesh or this

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sea of electrons.

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And so the metallic crystals, depending on what cases you

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look at, sometimes they're harder than the ionic

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crystals, sometimes not.

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Obviously, we could list a lot of very hard metals, but we

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could list a lot of very soft metals.

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Gold, for example.

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If you take a screwdriver and a hammer, you know, pure gold,

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24-carat gold, if you take a screwdriver and hit it onto

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the gold, it'll dent it, right?

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So this one isn't as brittle as the ionic crystal.

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It'll often mold to what you want to do with it.

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It's a little bit softer.

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Even if you talk about very hard metals, they tend to not

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be as brittle, because the sea of electrons kind of gives you

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a little give when you're moving around the metal.

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But that's not to say that it's not hard.

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In fact, sometimes that give that a metal has, or that

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ability to bend or flex, is what actually gives it its

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strength because it's allowed to kind of deflect the force.

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So the strength, and I've touched on this, it also goes

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into the boiling point.

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So because these bonds are pretty strong, it has a higher

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boiling point.

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If you just took salt crystal and tried to boil it, you'd

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have to add a lot of heat into the system.

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So this has a higher boiling point than say-- I mean,

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definitely things that have just van der Waals forces like

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the noble gases, but it'll also have a higher boiling

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point than, say, hydrogen fluoride.

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Hydrogen fluoride, if you remember from the last video,

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just had dipole-dipole forces.

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But what's interesting about this is they have a very high

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boiling point unless they're dissolved in water.

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So these are very hard, high boiling point, but the ionic

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crystals can actually be dissolved in water.

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And when they are dissolved in water, they form

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ionic dipole bonds.

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What does that mean?

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Ionic dipole or ionic polar bonds.

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And this is a situation where the sodium-- and this is

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actually why it dissolves in water.

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Because the water molecule, we've gone over this tons of

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times, it has a negative end, because oxygen is hoarding the

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electrons, and then the hydrogen ends are positive

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because the electron's pretty stripped of it.

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So when you put these sodium and chloride ions in the room,

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or in the water solution, the positive sodiums want to get

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attracted to the negative side of this dipole, and then the

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negative chlorides, Cl minus, want to go near the hydrogens.

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So they kind of get dissolved in this.

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They don't necessarily want to be-- they still want to be

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attracted to each other, but they're still also attracted

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to different sides of the water, so it allows them to

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get dissolved and go with the flow of the water.

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So in this case, when you actually dissolve an ionic

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crystal into water, as an ionic crystal, not a good

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conductor of electricity, not a lot of charge that is really

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movable in this state.

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But here, all of a sudden, we have these charged particles

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that can move.

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And because they can move, all of a sudden, when you put

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salt, sodium chloride, in water, that does become

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conductive.

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So anyway, I wanted you to be at least exposed to all of

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these different forms of matter.

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And now, you should at least get a sense when you look at

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something and you should at least be able to give a pretty

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good guess at how likely it is to have a high boiling point,

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a low boiling point, or is it strong or not.

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And the general way to look at it is just how strong are the

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intermolecular bonds.

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Obviously, if the entire structure is all one molecule,

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it's going to be super-duper strong.

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And on the other hand, if you're just talking about

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neon, a bunch of neon molecules, and all they have

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are the London dispersion forces, this thing's going to

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have ultra-weak bonds.

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So a gas is almost its most natural state.

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If you get super, super cold, you might be able to get it to

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a fluid, and then everything in between.

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Anyway, hopefully, you found that useful.

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