States of matter
Covalent networks, metallic, and ionic crystals Covalent Networks, Metallic, and Ionic Crystals: Some of the strongest molecular structures.
Covalent networks, metallic, and ionic crystals
- 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
- 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
- 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
- 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.
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