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Covalent networks, metallic crystals, and ionic crystals

Video transcript
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. 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, 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 sth. 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 neighborhood. 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 electricity. 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 conductive. 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.