Chemistry: Types of Solids

This video is too basic

02/04/2012

~ Lee15

I thought cubic crystals will be covered in much more depth in this video. Unfortunately they didn't. Since mindbites give you only 60 seconds preview you cant know this, but its a very basic video so if you need to know about SCC, BCC, FCC, and more structurally oriented chemistry like I did then this is not the video for you.

Also this is something that you could read in a book. No explanation of complicated concept what so ever. Why is this worth 3 credits and not two then? =(

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So here's another riddle for you; what do a candle, diamonds, carbon dioxide, butter, glass, mothballs, aluminum, a crystal, Styrofoam and a silver spoon all have in common? They're all solids! They're all solids, but they have very different properties. We know a wax is kind of a soft material and easily bends, certainly melts very easily. We know diamonds are extremely hard - we borrowed these from the Smithsonian just for this, in fact, so I've got to get these back. We know butter, again, very low melting material. Carbon dioxide has the property of readily subliming, so very, very different range of types of properties that we have. Metal, we know, is very shiny. It has the property of conducting electrons very easily, conducting heat very easily. So the goal here is now that we're experts on understanding the molecular level and the interactions between molecules, how do we make sense of the various different properties that solids give us?
Let's clear the decks here and talk a little bit about the molecular level and the different kinds of solids that we have. There are two general types of solids and they are kind of classified as crystalline solids versus amorphous solids. So let's talk about crystalline solids first. The point of crystalline materials is that there is, on a molecular level, a very ordered array. Now, this ordered array can be on an atomic level, in other words, for instance, a crystal of solid xenon could have a nice ordered array like this, or this might be a solid crystal of iron that has a nice ordered array of atoms. They can be ionic materials, in other words, this might be sodium chloride, for instance, positive and negative ions stuck together in an ordered array. They might be at a molecular level ordered, in other words, I show here the molecular structure of naphthalene and naphthalene is the chemical used in one of the two forms of mothballs. And although it's a neutral molecule, it, too, has considerable ordering in the solid. In other words, all the naphthalene molecules are lined up together and, in fact, let's see if I can get a picture here for you of that. So again, this is just a little bit better so that you can see it on camera here. You can see the ordering in the molecule, naphthalene, this two-ring system.
So my point again is not for you to memorize what naphthalene is, but simply recognize that the common feature of crystalline materials is that they have this high degree of ordering. An analogy I like to make is with a brick wall, that it goes on and on and on and you have this very long-range ordering as a result. Now, there are certain properties that result from the long-range ordering, and one of the important ones that we mentioned earlier is the idea that electrons or x-rays can be diffracted from these crystals. In fact, that is an important technique for determining what the spacing is between the atoms, even on a level of where all the atoms are in individual molecules. We can use that as a technique to actually get a picture of what these molecules look like. You may have been finding yourself wondering, "How do they know all of this stuff? How will I know that they're not just pulling the wool over my eyes?" Well, we know this from primarily x-ray diffraction, as far as the technique that lets us know what the molecular structure of atoms is. In order to do x-ray diffraction, we must have this long-range ordering to set up these planes that we can use to diffract the x-rays or electrons. So again, that's the common feature, but they come in many different flavors, if you will. There are molecular crystals, there are atomic crystals and there are ionic crystals.
Now, in contrast to that, we have amorphous materials, amorphous materials would include principally waxes, such as candles, and glasses, such a glass, for instance. But not only glass, plastics are considered to be glasses, polymers, Styrofoam cups, for instance, are polymers. These are all amorphous materials and basically an amorphous material is just something that doesn't have any kind of a long-range order. In the case of glasses, in particular, glass that we're familiar with, this is silicon and oxygen. There is considerable ordering around the silicon and oxygen atoms that we, in fact, could predict from what we know about VSEPR theory, which you've seen a little bit about. But there is no long-range order. In other words, there might be a little bit of ordering here, but it's going to be completely different than the kind of ordering here. So it's different than the type of long-range order that we saw previously, where we had these long, long arrays of atoms all in perfect order.
So we would subdivide again as a start. Materials we categorize in those two basic forms, so let's talk now a little bit more about ordered solids, the crystalline materials, and look again into the types of different lattices that we would have for crystalline materials. Basically, three types of lattices, metallic crystalline lattices, so this would be what we'd have, for instance, in silver metal, a crystal of silver metal, lead, nickel, any number of different metals that we're familiar with. What characterizes a metallic crystal is that it has what we refer to as delocalized electrons. Now, what I mean by that is it's as if all of the individual atoms gave up a couple of electrons and put in kind of a community pool and agreed amongst themselves to share that pool of electrons over all of them. And so, I color in blue here, kind of this delocalized electron density, which is kind of moving freely throughout all of these atoms here. And I'm symbolizing it with the positive charge here, the remaining nucleus and core electrons that are left behind. So that's the basic notion that we have of metallic systems. One of the properties of metallic systems is that they conduct electrons easily and you can get a better sense of why in that some of these electrons are free agents. They easily kind of move throughout the solid, they're not restricted to being around one atom or another atom.
Ionic lattices, as the name implies, involves ions. Some common examples of this would be sodium chloride, potassium fluoride. It doesn't have to be monoatomic anions, so, for instance, ammonium chloride would be an ionic lattice. Sodium sulfate would be considered an ionic lattice. Magnesium chloride is another example; notice in magnesium chloride or sodium sulfate there is a 2:1 ratio of ions. All of these things still have long-range ordering and are considered to be ionic lattices.
Now, the third category are things that we call covalent networks, very much like the idea of an ionic lattice, but the difference is, in a covalent network, there's not sufficient difference in electronegativity between the atoms to cause a strong polarization of the bonds. Remember our notion of an ionic bond means that it's so polarized that we like to consider all the electrons almost completely with one partner and not the other, whereas the notion of a covalent bond is that there are more shared electrons. And so, it's not as easy to break the bonds in a covalent network into individual ions. And so, examples of this would be, for instance, sand, silicon dioxide, alumina, very common in minerals. Diamonds are a form of a covalent network. Graphite, which is just a different form of carbon, also would be a covalent network in that case. So, in contrast to ionic lattices, this does not easily break apart into ions. So that's the fundamental difference between these two things.
So finally, let's try to make some sense out of the general physical properties of these different materials. Again, let's start with molecular solids. Carbon dioxide would be one example, but there's a huge range of molecular solids. The only requirement is that we have discrete molecular units rather than a lattice. Again, think back to the brick wall idea. Molecular crystals can have hydrogen bonds, dipole-dipole interactions or just London forces holding them together, so a wide range of different types of interactions holding them together. Their electrical properties, in general, they are very poor at conducting electricity or heat. They're considered to be good insulating materials. And general features, these materials, compared to these other types of crystals, tend to be soft and relatively low melting. And again, you can understand why, because the bonds between the molecules are very weak, and so it's easy to break them into their pieces.
A covalent network - now, remember what that is. A covalent network is where we've got strong covalent bonds in all directions. Once again, think of the brick wall, so we have strong covalent bonds. In other words, the bonds between the molecules, if you will, are exactly the same as the bonds within the molecules. There are no real molecules. It's one big hug molecule, in other words. So this also is an insulator. There aren't easily available ionic charges that you can get access to and, as you might expect, these materials are going to be very hard, because they have a very well established rigid structure to them and no easy breaking points. They tend to have very high melting points and even higher boiling points. So think about sand as a classic example of that covalent network.
Ionic materials are held together by electrostatic attractions primarily, they're ion-ion interactions. They also tend to be insulators. Now, that might throw you off a little bit, because they're ions after all so they're charged, but they're trapped ions. They're stuck next to each other and they cannot freely move around. I have the asterisks here, because when you melt these materials, then you have an excellent conductor, because now the charges are free to move. As long as these materials are in their solid form, they are insulators. And finally, these things also tend to be hard, brittle, and relatively high melting, very often not as high melting as the covalent networks. So high melting compared to molecular crystals, but not nearly as high melting as the covalent networks.
Finally, we have metallic networks or crystals, and remember the key about that is these delocalized electrons, these free agent electrons that wander throughout the solid. That gives them the property of being good conductors, as we know metals are, and it also gives them the property of being very ductile, so they're easily bent and formed and shaped, because there aren't rigid bonds holding them together. There's just this kind of sea of mobile electrons holding them together. So they don't particularly care in what their general shape is. So we can understand why that feature is so different than these other materials.
So, once again, the lesson is the same - by understanding what happens at the molecular level, we can make sense of our macroscopic world. We can explain physical properties of common, everyday materials.
Condensed Phases: Liquids and Solids
Solid State: Structure and Bonding
Types of Solids Page [1 of 2]