Understanding Intermolecular Forces and Crystals

We're going to go to the next slide because we're kind of finished with the hydrogen bonding. And thanks for that clarification question. And now we're looking at just now we're on to liquids. We're still kind of talking about liquids. We're talking about intermolecular forces. But some neat applications of intermolecular forces and what intermolecular forces explain is things like surface tension. Surface tension is the amount of energy required to stretch a surface. And what's really cool about this is that this. top has a somewhat of what's called a skin and this skin on here and there was a really my daughter studying this in her AP chem class and right now and there was this really neat question about I think I might ask it on you guys but I'm gonna give it this is ethylene glycol can ethylene glycol undergo hydrogen bonding yes or no what would you say type it into the chat Yes. Now here's water. Can water undergo hydrogen bonding? Yes. But water has a much higher surface tension. Oh, wait, let me ask you this question. How many hydrogen bonds can water make? If you had to count, remember there are lone pairs here. One, two, three, or four. Okay. I see people saying four. I see people saying two, but now there's kind of a little open-ended thing here. This hydrogen can make a hydrogen bond with something else. So can this one. This lone pair can make a hydrogen bond and so can that one. So I call it four. You'll hear some people say two, but that's only because there's a definition between a hydrogen bonding donor and a hydrogen bonding acceptor. I'm calling them, I'm just looking at the water and saying, well, it can make four, whether it's accepting or donating. So my next question is, how many hydrogen bonds can ethylene glycol make? So type in the chat whether you think it's two, four, or six. See, I'm drawing all these little dots. I'm making, yes, that's six. It can make six hydrogen bonds. So it seems weird that water has a higher surface tension than ethylene glycol. Now, why do you think that's the case? actually I'm not going to talk about it any further. I'm just going to, I think I'm going to ask you that as a post-lab question for our first experiment, because we'll be dealing with ethylene glycol in that experiment and water. So we'll keep that in mind and we'll look at our surface tension measurements for both of those and see what they do. Because I think we figured it out when I was talking to my daughter last night, but I still have a couple of questions about it. So keep that in mind. So anyway, something can have a surface tension because it can pull the top. surface molecules are being pulled down from the bottom and it makes this skin. And of course, that's why water striders can walk on water. So cohesion, the reason that happens is because of these different kind of definitions of cohesion is intermolecular forces between like molecules, whereas adhesion is between unlike molecules. So in our situation here, this is mercury, and mercury likes itself. better than it likes glass. Whereas water likes glass better than it likes itself. So water's adhesive properties are greater than cohesive. So that's why water will rise in a glass tube and mercury will fall. Mercury is really the only one that does that. I've never really seen anything else. So anyway, again, surface tension was a measure of um was a measure of IMFs of that liquid. Viscosity is a measure because it's a resistance to flow. Ethylene glycol, super viscous. Water, not as much. Look at these viscosity numbers. There's viscosity of water. And then I feel like, you know, ethylene glycol is not in here, but what is on here is glycerol. And glycerol is like ethylene glycol, very sticky substance. So when you look at those numbers, the stronger the intermolecular forces, the higher the viscosity. So of course, Glycerol must have higher intermolecular forces. Any questions on that? Oh, there, see, I circled them. Diethyl ether has the lowest surface tension. I'm sorry, viscosity. Diethyl ether is so, you know, you open that, that will just crawl right out of the beaker and you'll smell it very soon because it has very little intermolecular forces. So as far as water is concerned, Water is kind of unique because of its small size and its high polarity and its ability to, when it freezes, it expands. So you've probably heard of that before. The magic number here is four degrees. When your water gets cooled down, it will eventually get to a point where it starts to form that cage structure you see over here and it will then expand and float. And then, of course. people can ice fish and fish don't die from being stuck on the top because that would be a problem. Okay, any questions about because we're kind of finishing that bit there that was about surface tension and about viscosity. Are we good to go with that? Pretty good. Yeah, I had a question about the ethylene glycol and how many hydrogen bonds there were. Okay. Are you just counting the oxygen and hydrogens and the ethylene glycol to get to your number six? Well, I guess, you know, how did I count that is I counted it by lone pairs and the hydrogen. So this hydrogen can hydrogen bond with something and each lone pair can hydrogen bond with a hydrogen on something else. So that's how, that's how I was counting it. Did that answer your question? I'm still trying to like count it to see because I don't see it yet. Let me circle. I'll circle with a different color. Thank you. One, two, say this was going to something else, three, and then again, four, five, six. Oh, okay. Okay. Thank you. Because there's, you know, the H2s. in the middle, so I was thinking that it would have been a bigger number, but they're connected to the carbons? Correct. Okay, thank you so much. So it's almost like, are you thinking about this, because this is one of your assignments for the first lab, is count number of hydrogen bonds per molecule, is that if you have an OH, okay, and then on a group like this, each OH can do three. Now, that is with other things. But if you have something like acetone, like I had on that previous slide, and then these guys do not hydrogen bond, these CH3s, right? Those agents do not hydrogen bond. Acetone does not have any interactions through hydrogen bonding with another acetone molecule because there's no OHs on the acetone, another acetone molecule, but it can hydrogen bond with water. So that's why it likes water so much because it can interact. in that way and dissolve and be missable with it. So anyway, does that make sense? Yeah. Thank you so much. Sure. That was a great question. That is hard. That's a tough one. And people want to see that and they think, oh, H's, H bonds, but you really have to analyze what it's attached to. Very good. All right. Let me go to my next thing here. Here we go. Now we're going to jump in to crystal structures. Now, crystal structures are things like salt, right, or ice. OK, all these things are crystal structures and they have an order to them. So whereas amorphous solids do not have an order. Now, the question is, what is that order? That order is through what's called a unit cell. In a unit cell, there's like a certain individual. But then when you grow it in all directions, you get. a three-dimensional kind of xyz axis thing here and it will grow out and become very ordered and you get that beautiful crystalline structure that catches light in certain ways and that's why it looks pretty like that because light comes in and then bounces off and is diffracted and you know you see like a diamond or something like that or ice. So anyway each one of these little dots here can be either a molecule or an ion, or a metal atom, because all of these things are crystal structures. And they're called lattice points. Atoms, molecules, or ions. And they have different shapes based upon what they're made of. So these little shapes here, these shapes then have parameters that are described by the angle of the box or the length of the sides. So we have... angle. We have lengths of sides, and those lengths of sides are called A, B, and C. And then the angles are called alpha, beta, and gamma. So whatever those angles are, whatever those lengths are in relation to each other. dictates the crystal structure and how it organizes itself. Any questions about this picture? I think this is really fun, but hard at the same time because it's hard to visualize, but it's really cool. And I have in my online, my completely online class, I have a complete discussion. I don't know, Chris, you wanted a discussion. I can give you that discussion. that you have to look up a metal, I mean, or compound, an ionic compound, look at all these things. But we usually only focus on this one. Because when you get to these things, look at these angles. I mean, and look at these angles, 120, even though you can determine like how many are at each point. And then you can figure out like what the density of the substance is by calculation, which is really cool. I love it. It's so fun. So let me show you some more of that. in that previous slide was very similar to just looking at those little dots. Now we're going to grow it out and look at more space filling. And the space filling, when you do that, you can see the lattice, the unit cell is in here. And on each one of these corners, these are corners, and actually a corner is only one eighth of that sphere. So when it's one eighth of that sphere, 1 8 8 is equal to 1. There's really only one atom in that unit cell. And you're probably asking, why do we care? The reason we care is because x-ray crystallography was developed based upon all of this understanding. And that you can use x-rays and you shoot them through and then they diffract out of a crystal. So it's got to be crystalline because it won't diffract nicely if it isn't. And then you can determine the distances from here to here. And from here to here, just based upon... the x-rays that come in and the x-rays that come out. It's really the coolest thing. Now, if we look at this one, now just what I said there, what do you think? I'm going to draw this again, but this guy is in the middle. You see him? He's in the middle. How many atoms are in this, what's called body-centered cubic unit cell? There was one in the simple cubic. How many do you think is in the body-centered? I hear a couple people saying two, and that's correct, because there's one eighth at each corner times eight equals one, and then there's plus one from the center. Perfect. How about this is face-centered, and when we do face-centered, each one of these is called, it's in the face. So when it's in the face, and this is harder to see up there, if you take a slice of it, you slice it in half. So there's a half of a molecule. half of an atom rather, in each face. And there's one at every corner. So how many atoms are in a face-centered cubic arrangement? Very good. I hear people saying four because there is one eighth times eight, which would equal one. And then there are one half times six, which equals Three. Three plus one is four. Very good. And here's the picture of it. So you can see this like sliced up and that's what I was trying to do there. And that would be simple cubic. And then this one would be body centered. You see the one in the middle and then one eighth at each corner. And then you see, oh, I didn't do the same for face centered. Looks like we're going to do the same thing that I was just going over. Any questions on that? Now, I ask you to draw this on your fun sheet. I say draw these things and do some calculations with it. Because if I know that there's four atoms per unit cell, yeah, coordination numbers, we're going to get to that here in a second. If I can count the number of atoms in a unit cell and I know who it is, for example, I can know the grams of one atom. And say there's four atoms. Well, I'd multiply that by four and say, I know this is a cube. How do you get the area of a cube? I'm not the area, I'm sorry, volume. Remember, it's length times width times height. That means I know now the volume and I can get density is mass over volume, mass of the four atoms over the volume of that cube and how many atoms were in it. Now, coordination number, in order to know coordination number, you kind of have to know further out in the lattice. And let's talk about that. But before we get to that, let's talk about closest packing, because closest packing. is a way to, is how those atoms are organized. So closest packing, they organize themselves like in the little indentations. So you have one layer, for example, and then the next layer goes in the indentations, and then the next one goes in the next indentation. So it kind of makes sense. It's making most efficient use of space. So there's one, there's different ways they can do it. See, it can either go in there or it could have gone in these little holes. And that's what that one's doing. That's cubic. And this is what we're looking at most of the time. Hexagonal does a slightly different thing. Let me see if I can whip back real quick. I wanna look at that hexagonal. Bear with me. Don't look, I know it's looking freaky. Where is it? Don't look, sorry. Right there, hexagonal. See that? That hexagonal means you get those 120 degrees. And you get some neat crystals. They're like grown crooked, but they're really pretty and very prismatic. So before we can get to determining coordination numbers and understanding what that is, we have to grow the crystal out. And to grow the crystal out, they have to pack like they do in there. And there's, look at that coordination number, 12 for hexagonal and cubic close packing. Because what ends up happening is a lot of these are face centered. And when they're face centered, they pack really tightly. So they're getting more around, hence the 12. Because there's other coordination numbers of eight and six. So let's look at those and see if we get that coordination number coming up. I'll let you guys look at these. These are kind of hard. I have a video in content that you might want to look at that talks about all of this and zooms you in and out. It's only like six minutes long that you might want to look at to help you understand that. And I'll put it in an announcement. I'm going to get, okay, before we get to, I'm keeping the coordination number in mind. But before we get to that, let's just look at the relationship between these things. The relationship between edge length and radius of an atom. If it's simple cubic, see how easy that is? If you measure the edge length, you can figure out the radius of the atom. Because did you ever wonder how all the sizes of atoms were determined? This was one of the main ways to do it. Because you would shoot an X-ray through, and it would come out, and you could determine distances. And if you would determine distances and you knew what was in here, you can figure out basically atomic radii, which is on the very small level. So these things you just want to review. If you have body-centered or face-centered, there's relationships. And your fun sheet has some of these questions in there to ask you to use this and use Pythagorean's theorem. You know, that's what this is using. Your book goes through that. So let's just talk about this real quick here and do a quick example. So I'll let you read that. So when it's face centered, you know how many are in the unit cell. Do you remember how many are in the unit cell if it's face centered? Type that into the chat if you remember. It's four. That's right. You guys got it. So when it's four, you can know if it's silver. You say it's silver. Well, what's the density of silver if there are four atoms per unit cell? And we need to know this 409. Does anybody know what P stands for? in the pm the m is meter that's right it's picometers so since it's picometers picometers is 10 to the minus 12 so we have to remember that and we have to remember that it's a cube so we can determine the volume and since we can determine the volume we can use that in our calculation so that's what we're doing there this is centimeters this was picometers cubed And now it was converted to centimeters cubed because in density, remember, D equals M over V, we want to have this, it's always in centimeters cubed, which is the same as milliliters. So if we know it is four silver atoms, then we can use the molar mass of silver to determine the mass of those four atoms. And that's what's happening here. And then we get a number and it's going to be teeny, teeny, tiny. You see that? It's a very small mass and then we can plug it into the density formula. So look at these numbers. See why they're so small? Of course it makes sense. And then you're going to get the density of the silver. And sure enough, you compare that with the literature values, and sure enough, that's what the density of silver is. So this density was determined way before the crystal structure of it was. So this was a way to verify what's on the molecular level or atomic level of silver. Any questions on this one? Any questions on how to get from picometer cubed to centimeter cubed? Because that's one of the main stumbling blocks for this. And remembering that you have to use Avogadro's number because Avogadro's number takes it on the atomic level instead of the mole level. OK. All right. Now, how is this done? This is this X-ray crystallography I was telling you about. You have a crystal. You shoot x-rays through it of a certain wavelength. And if you know the wavelength, which you can adjust, and you know the angle that this thing, it's being, it's coming in at, it will come in and it will refract as it goes through the crystal. And then mathematically, that can be determined to figure out where these spots are. And I have a picture here of what a real one looks like. This is what I actually did this in my graduate studies that you can actually. get all these spots and you know the original x-ray crystallographers actually measured they measured with rulers it was just crazy but they were the pioneers for this system this thing i don't know if you wonder there see this like whiteness here this whiteness is the actual like what's holding the crystal it's called the goniometer it's a funny it's a funny word but it's um it's the the x-ray gets it's this and that's why you see it in there. The x-rays are absorbed like your bones in an x-ray, so it's not seen, but the other ones get diffracted in this fashion. So this theta angle that's coming in here and then it's diffracted, this is through, is described through what's called Bragg's equation here, and that is depends on Bragg's law is the angle and the distance between them are related to the wavelength of the incoming x-rays. So if you know the incoming x-ray wavelength and you know the angle for which you shoot it's coming in, you can determine the distance between the layers like this. So there was those, remember we were seeing those pictures, and now you can determine the edge length. And that edge length can give you information about the crystal structure. Right now it's called the Bragg equation. It's very automated. I mean nobody measures anything anymore. So I'm sorry, you know, somebody asked about coordination numbers. I didn't really cover that too much. But in that video that I'm going to put out, it shows you how to do that. It shows an atom and it shows how there's atoms around it. And as far as what coordination number it would be. So you'd have to know the entire crystal structure and see it kind of built out to be able to see the coordination number. But face centered coordination number is 12. body centered is eight and face centered is six. I'm sorry, face centered is 12, body centered is eight and simple cubic is six. And so I'll show you that video. I'll put that out in an announcement and let you check it out. But as far as crystal structures go, if you know through this Bragg equation, what you're, you're trying to determine the distance then between layers, assuming n equal one, and when assuming n equal one, that means you have one layer. and then you have another layer. This is n equal 1. This would be n equal 2. And then you could say a third one. That would be n equal 3. Okay, so this is much more complex than this, but this is kind of the simple part of it. So if you know what the wavelength is, and you know what your n is, you can solve, and you know the angle, and you can solve for it. And you can solve for the distance. And in this example, you're getting 314 picometers. Oops, sorry, I meant to just erase that. I know this is a lot to take in because crystallography is complex stuff, but this is just kind of an introduction. But does anybody have any questions about, because I have a fun sheet question about this on there, that where you calculate some things and calculate the distance between layers in the crystal. So what's kind of cool about crystallography is that, see, here's my pen, if you can see me up there. They put the crystal on the top here. And then this thing, either the X-ray source. can comes in and hits the crystal and then they move the crystal. The crystal gets moved in all directions and by moving it in all directions relative to the x-ray source you can take all of these measurements because you get all these spots from all different angles and I have still yet to find a really good video about it and showing how this works but that's how it's done and then you can determine the distance between them. So as far as crystal types, so that was kind of ending X-ray crystallography. Now the different types of crystals are metallic crystals, covalent, ionic, and I think there's one more, covalent network solid. So a crystal type, I'm asking you to name that crystal type if you've looked over this information already. If it's composed of charged species, what kind of crystal do you think that is? Is it ionic, covalent, molecular, or metallic? Type it in the chat if you think you know it. If it's ionic, if it's charged species, it's ionic. Exactly. So that's probably what you're very familiar with because we're so used to things like sodium chloride, etc. And the anions and cations, they're different in size. It's got a very high melting point. The lattice energy is huge because it's very stable. It's pretty brittle. And they're good electrical. They're not good conductors in a solid. But... when you melt them, they will conduct electricity or aqueous. So that's an ionic crystal. So this is something that you might see as an ionic crystal. Can anybody tell me what type of packing that is? Is it simple cubic or simple cubic, body-centered cubic, or face-centered cubic? By that picture, what do you think? I hear somebody, I see somebody saying body. That's true because you see the one in the middle. It's BCC. Very good. This is called zinc sulfide. What do you think that crystal type is? FCC? I see FCC and that's good. But now I'm going to ask you a question here because I think this gets really hard. Because here you notice there's ones in the faces, right? So that means face centered. But you have these guys in the middle. And these guys in the middle, do you think those are the zincs or the sulfides? based upon that picture, those little gray ones? Do you think they're the zinc or the sulfur? Now the next question I would ask myself is how can you tell? How can I tell whether they're zincs or sulfurs? What's the difference between zinc or sulfur? Zinc is positive two, sulfur is negative two. Now do you remember from chapter eight, are cations smaller or larger than anions? This is what you would want to think about here. Because you can see it in the picture. These gray ones are smaller than the green ones. Now, remember, when cations lose electrons, they get smaller. When an atom gains electrons, they get bigger. So it looks to me these sulfides are the green ones. And the zinc is the middle one. Now, let me ask you another question. If I had to count zincs in this unit cell, how many are there? That is correct. You guys see people saying four. So I'm going to write ZN4. If you had to count sulfurs, how many are present in this unit cell? Now, remember, when you just count. But in the unit cell, you have to remember that there's only a piece of it in the unit cell. Yeah, that's the tricky part. So now, you know, there are four. So what is the molecule? What is the formula for zinc sulfide? It's an empirical formula, which is the lowest ratio. What is the formula? And it should match what we know and what is in the unit cell. It's ZNS. Remember, you did. Yeah, exactly. So what about cesium chloride? How many cesiums? And what's the one in the middle over here? Is that cesium or chlorine? So now I'm on the other one. It's cesium. And there's one of them. How many chlorines are them, which apparently are red? One. So it's C-S-C-L, one to one, which it should be, right? We know that about cesium and chlorine because they are one to one. One's an alkali, one's a halogen. So we do know that. Oh, here's another one. This one looks crazy. How many, this is calcium fluoride. How many gray ones are in that unit cell? Now let me ask you this. What is the arrangement of the gray ones? Is it simple cubic, body-centered, or face-centered? So I'm getting some differing opinions here because look at these guys. These silver ones are faces. So it's face-centered, and if it's face-centered, how many gray ones, which I believe are calciums because they're smaller, They're the cation. How many are in that unit cell? Because remember, yeah, you have to only count the pieces that are in there. And there's in the corners. Of course, you have an eighth. And in the faces, you have a half. So that should be four. How about the green ones? How many green ones are inside that unit cell? They're actually inside. So you can just count them. Yeah, F8. And what's the molecular, what's the empirical formula? See that? It's the lowest ratio. And that it should match every time because what chemistry is doing is studying matter and we're zooming in and we're trying to figure out the ratio and what's actually in there. It's like studying anatomy of a person, except we're studying the anatomy of the universe. So that's why it's such a big subject. Anyway, anybody have any questions about what I just did there? Fun stuff. Sorry, I've got a quick question. Can you just explain again how you determine how many calcium atoms or ions that I guess in this case there are in that structure? Right. So what I'm understanding is that like these structures with their little spheres are not exactly like that's not exactly the location of a of an ion of a ion or an atom or is it? It's the location, but not all of it is actually in this unit cell. What does that mean? It means like if we look at, let me go back to this picture. Sorry, maybe I missed a little something. No, it's hard. I think it's hard personally. I'm just right there. So when you look at these pictures here and you say, when we do that counting, you have to take into account that there's only a certain portion of that in that unit cell. So on the picture that we were just looking at, the calcium's were face centered. Oh, okay. Okay. So there, so when they're talking about unit cell, they're taking essentially like this cubic measurement sort of thing. Yes. And they're slicing it. It's being sliced. Okay. So when I show you that, when I share that video with you and you watch it, it'll, it'll make a lot of sense after we've talked about it. And I'll share another link called ChemTube3D. It is like the fabulous thing, ChemTube 3D. And it enables you to click on one of these unit cells and rotate it and really see. It doesn't really slice, but it'll let you do one unit cell and you can move it around and to determine the number of atoms inside. Because some of them get a little bit more complex like that last one, that calcium fluoride we were just looking at. Yeah. Yeah. OK. Thank you. OK. Anybody else have a question about that? You just have to start talking because I really can't see your hand. If you have a question, I was just going to ask it, it looks like the ratios are. By the valence electrons, like, how many balance electrons each like. The calcium and the, and the flooring for instance. Like, calcium has for the electrons, therefore. And pouring needs 1 more because it's really close. So that's why there's 2. fluorine per calcium. Does that have something to do with how you're counting all of this? Well, I think it more has to do with the ratio. Like you're saying, it's a one to two ratio. Calcium has two electrons, two valence electrons, you know, and fluorine has your seven, right? So you only have, it needs one and then it needs another one, right? So these crystal structures, what we're studying right now is ionic and the ionic compounds, they build by size. So they have these different structures because this one's too big, so they can't snuggle right. So they have to have a certain type of structure. We can have molecules can do this too, where you'll have a water molecule at each little place. But the water molecule might look like this, and part of it's over here, and part of it's in the unit cell. So molecular crystals do a little... They they're following face centered, body centered, simple, but they only a part of a molecule, for example, might be in there. So it does have an effect, like you're saying, that the valence electrons do mean something because as far as the ratio is concerned and as far as the size and then which structure they take based upon their sizes. Because let me ask you this. Isn't cesium huge? Think about where cesium is on the periodic table. He's down the lower left. He's monstrous. So he takes up that. big center area and then you got these chlorines you know around him whereas calcium is not as big as cesium so look how many they can fit into the unit cell does that make sense and zinc is a little larger yeah so well so i guess um i just can't this is just ionic so we would still have to do like the counting and everything with the um you the crystal structures and figure out which crystal it is to, and, and how many, um, I guess how many like angles there are for regular molecules, because they're not going to follow like the same rules as an ionic structure. Well, you know, I think I'm just messing up left and right. I can't stand it. I'm going to have to show you this chem tube 3d thing, because once you see this, it really helps. I think, um, to visualize what we're talking about here. So let me just run over and I can, I know exactly I've been using this site for quite a while now. I used it to create a discussion based upon counting of atoms and like this, because it was hard. So ChemTube3D, it's out of University of Sheffield in England, and it's just the coolest thing. And you go here and then you can go to any one of these structures. Can you guys see what I'm looking at? Okay. And, and say for, I just want to do a simple one, cubic close packed. We just learned about that. And then it'll bring up a structure for me. Hopefully it must be thinking. Can you guys see that moving around there? Okay. Now it looks like to me, is this, what, what is this? Is this, is a simple cubic body centered or face centered? What do you think? That's right, it's face. So you're seeing it. So there's even more to this. It gets crazy. You can click on this stuff and it shows you like other stuff outside of that unit cell. But what I wanted to show you was kind of that what you were asking about the molecules, because I feel like let's see if we can find some molecules in here. This stuff has so much good stuff that you can't stop looking. You know, I don't know. It does have water in here. Search. water. You can do a search in here and you can search water. And we want some ice water in the solid state. And then you can look at this a little bit. Now, I don't know. I don't feel like this is helping me. I feel like this is just a freaking mess to try to figure out what it is looking. I know I can't see the unit cell. I would need a unit cell. Sorry about that. I shouldn't have showed you that. Well, I don't know. I'm not really finding. This only has a lot of the inorganic things that we were just learning about, like that zinc blend, zinc sulfide. We just counted on the PowerPoint. That's right here. So you can look at it, and it gives you this structure. And you can click on the CCP, and it gives you the unit cell only. And then you can move it around and go, oh, I see there's four gray ones in there. And that's what we just did. And I see that there's, you know, it's face centered. So I see. Some of the yellow ones in the faces and the yellow ones on the corners. So that was ZN4S4. Remember, we saw. So, oh, there it is. See, ZNS4. And sometimes you can even see coordination. Remember, somebody was asking about coordination number. And you can see coordination. So you can click on it and it starts giving you some weird things. See that? There's things in there. It's talking about the coordination around each atom. So it gives it like a tetrahedron. Because that's how it's organized. There's a lot to this, I know. But anyway, I just had to show you that because it was kind of cool. Anybody have any questions on that before I get out of here? Your book has a question, and I encourage you to look at it because it's helpful. It has a question about the unit cell and number of molecules of iodine, and iodine's a molecular solid. So where we are right now is we are talking about the different types. Of solids, and those solids, we just did ionic and we went through the different. Counting and you really only find that in this ionic situation. We just counted all of these. In this book, it doesn't really show it too much for the other types. Now let's name that crystal type. We want to move on to another crystal type. If the lattice points are occupied by atoms only. Thanks, Shannon. Yeah, what time is it? We need to get out of here soon. I got about 10 more minutes. If what crystal type is it, do you think? Type in the chat if you know. If the lattice points are occupied by atoms. So they're not ions. They're not molecules. They're atoms. And they're held to death. together by covalent bonds. What else? It has a high melting point. And it's a poor, oh, see, this gives it away right here. If it's a poor conductor of heat and electricity, could that be a metal? Yes or no? No. So it can't be a metal. It's not a metallic. It's called a covalent network solid. So a covalent network solid or a net. network solid or covalent solid. I mean, you hear different books call it different things, but I always, I learned it as a covalent network solid because it was networked, but it was all covalent bonds, but it was a bad conductor because all those electrons are localized. They can't move. So that's why it takes that form. Now this also has a certain unit cell. Look at this diamond. I don't know if you can see a tetrahedron right there. So it, on that chem tube 3D. there is a diamond structure for you to look at and you can move it around. So that is called a covalent network solid. Now we're at another one. What type of crystal is this? If lattice points are molecules held together by intermolecular forces, low melting point. And so again, a poor conductor, what type of solid do you think that is? Hopefully you have your book next. Yo, somebody does. Molecular. Very good. That's a molecular solid. And water is a molecular solid. So is benzene. It's funny though, benzene, solid benzene sinks, solid water floats. Okay, one more here, lattice points occupied by metal atoms. Well, that gives that one away, right? That means it must be a metallic solid. And sure enough, it conducts electricity. It has a high, it can have hard, it can be hard or soft, the metal. It can have a low or high, just depends on the metal. And this is like a sea of electrons. And these electrons just swim around here. And that means that it's a good conductor. Okay. So the crystal structure of metals has all these different types. You have the body-centered, the face-centered, the hexagonal. So all these metals take those different unit cells. And here's some cool crystal structures of different substances. And what do you think that one in the middle is there? Do you think that's an ionic compound, a metallic? substance or covalent. Yeah, it's metallic. That's easy to see because it's like, it looks like a metal, right? Which is shiny, right? And it's lustrous and that type of thing. This first one is copper sulfate, hexahydrate, pentahydrate, I'm sorry. And I think there's an example of them right over here. Maybe not. It doesn't give that to you. No, but it's an ionic solid. It has water molecules in it. This one, metallic, this one is sulfur. So sulfur, that yellow substance is sulfur. So that one is under a covalent solid. It's not up there as an example, though. Quartz is a covalent solid. So anyway, and this is just another example of that. And these are posted for you, these PowerPoints, if you want to look at them. But I think we're kind of done for today. Does anybody have any questions before we quit? Because I know that this is a lot to take in, but chemistry is a lot to take in, I think. How do you determine coordination numbers from just like the structure, like the simple cubic and the face centered or the body centered? I don't think you really can. Like you say, how do you determine? It's like saying. What I always akin this to is I always say we're visiting the chemical zoo. And what's going on in the chemical zoo? When I go visit zebras, they have stripes. And I just have to know that. When I have a crystal structure that's body centered, I just kind of have to know that it has a certain coordination number. Now, can you count it? If you saw a drawing? Can you guys hear me? Yes. Okay, because my whole thing is locked up. I mean, my PowerPoint has locked up and stuff. And it's spinning. But either way, you can look at that and you could count if you saw an extended lattice work. And you could say, oh, there's a sodium and I count six around. Well, that means it's got a coordination number of six and it must be a simple cubic. So you would have to have a picture of it is what I'm saying. And you'd have to be able to count it that way. Now, you could also just memorize that simple cubic has six, body centered is eight, face centered is 12. And if it does indeed have that, that must be a piece of evidence that it's face centered or whatever. Did that make sense? Yep. Thank you. Okay, but that's a good question. I'm going to go see if I can find after I log off here. Find a picture of what you're talking about and point at them and just count them. And that video that I'm going to share with you shows it in 3D almost. So it would help you. All right. Well, I'm locked up. My PowerPoint won't stop sharing. So I'm just going to say goodbye. I don't even know if I can quit the meeting. So you guys are just going to have to leave because it's locked up on my end. Okay, so I'm going to stop talking here and I'll see you guys on Monday. Bye. Thank you. Have a good one. That's sweet.

But some neat applications of intermolecular forces and what intermolecular forces explain is things like surface tension. Surface tension is the amount of energy required to stretch a surface. And what's really cool about this is that this.

top has a somewhat of what's called a skin and this skin on here and there was a really my daughter studying this in her AP chem class and right now and there was this really neat question about I think I might ask it on you guys but I'm gonna give it this is ethylene glycol can ethylene glycol undergo hydrogen bonding yes or no what would you say type it into the chat Yes. Now here's water. Can water undergo hydrogen bonding? Yes. But water has a much higher surface tension.

Oh, wait, let me ask you this question. How many hydrogen bonds can water make? If you had to count, remember there are lone pairs here. One, two, three, or four. Okay.

I see people saying four. I see people saying two, but now there's kind of a little open-ended thing here. This hydrogen can make a hydrogen bond with something else. So can this one.

This lone pair can make a hydrogen bond and so can that one. So I call it four. You'll hear some people say two, but that's only because there's a definition between a hydrogen bonding donor and a hydrogen bonding acceptor.

I'm calling them, I'm just looking at the water and saying, well, it can make four, whether it's accepting or donating. So my next question is, how many hydrogen bonds can ethylene glycol make? So type in the chat whether you think it's two, four, or six. See, I'm drawing all these little dots. I'm making, yes, that's six.

It can make six hydrogen bonds. So it seems weird that water has a higher surface tension than ethylene glycol. Now, why do you think that's the case? actually I'm not going to talk about it any further. I'm just going to, I think I'm going to ask you that as a post-lab question for our first experiment, because we'll be dealing with ethylene glycol in that experiment and water.

So we'll keep that in mind and we'll look at our surface tension measurements for both of those and see what they do. Because I think we figured it out when I was talking to my daughter last night, but I still have a couple of questions about it. So keep that in mind. So anyway, something can have a surface tension because it can pull the top.

surface molecules are being pulled down from the bottom and it makes this skin. And of course, that's why water striders can walk on water. So cohesion, the reason that happens is because of these different kind of definitions of cohesion is intermolecular forces between like molecules, whereas adhesion is between unlike molecules.

So in our situation here, this is mercury, and mercury likes itself. better than it likes glass. Whereas water likes glass better than it likes itself. So water's adhesive properties are greater than cohesive. So that's why water will rise in a glass tube and mercury will fall.

Mercury is really the only one that does that. I've never really seen anything else. So anyway, again, surface tension was a measure of um was a measure of IMFs of that liquid.

Viscosity is a measure because it's a resistance to flow. Ethylene glycol, super viscous. Water, not as much.

Look at these viscosity numbers. There's viscosity of water. And then I feel like, you know, ethylene glycol is not in here, but what is on here is glycerol. And glycerol is like ethylene glycol, very sticky substance. So when you look at those numbers, the stronger the intermolecular forces, the higher the viscosity.

So of course, Glycerol must have higher intermolecular forces. Any questions on that? Oh, there, see, I circled them.

Diethyl ether has the lowest surface tension. I'm sorry, viscosity. Diethyl ether is so, you know, you open that, that will just crawl right out of the beaker and you'll smell it very soon because it has very little intermolecular forces. So as far as water is concerned, Water is kind of unique because of its small size and its high polarity and its ability to, when it freezes, it expands.

So you've probably heard of that before. The magic number here is four degrees. When your water gets cooled down, it will eventually get to a point where it starts to form that cage structure you see over here and it will then expand and float. And then, of course. people can ice fish and fish don't die from being stuck on the top because that would be a problem.

Okay, any questions about because we're kind of finishing that bit there that was about surface tension and about viscosity. Are we good to go with that? Pretty good. Yeah, I had a question about the ethylene glycol and how many hydrogen bonds there were. Okay.

Are you just counting the oxygen and hydrogens and the ethylene glycol to get to your number six? Well, I guess, you know, how did I count that is I counted it by lone pairs and the hydrogen. So this hydrogen can hydrogen bond with something and each lone pair can hydrogen bond with a hydrogen on something else. So that's how, that's how I was counting it. Did that answer your question?

I'm still trying to like count it to see because I don't see it yet. Let me circle. I'll circle with a different color. Thank you.

One, two, say this was going to something else, three, and then again, four, five, six. Oh, okay. Okay. Thank you. Because there's, you know, the H2s.

in the middle, so I was thinking that it would have been a bigger number, but they're connected to the carbons? Correct. Okay, thank you so much. So it's almost like, are you thinking about this, because this is one of your assignments for the first lab, is count number of hydrogen bonds per molecule, is that if you have an OH, okay, and then on a group like this, each OH can do three.

Now, that is with other things. But if you have something like acetone, like I had on that previous slide, and then these guys do not hydrogen bond, these CH3s, right? Those agents do not hydrogen bond. Acetone does not have any interactions through hydrogen bonding with another acetone molecule because there's no OHs on the acetone, another acetone molecule, but it can hydrogen bond with water.

So that's why it likes water so much because it can interact. in that way and dissolve and be missable with it. So anyway, does that make sense?

Yeah. Thank you so much. Sure. That was a great question.

That is hard. That's a tough one. And people want to see that and they think, oh, H's, H bonds, but you really have to analyze what it's attached to. Very good. All right.

Let me go to my next thing here. Here we go. Now we're going to jump in to crystal structures. Now, crystal structures are things like salt, right, or ice. OK, all these things are crystal structures and they have an order to them.

So whereas amorphous solids do not have an order. Now, the question is, what is that order? That order is through what's called a unit cell. In a unit cell, there's like a certain individual. But then when you grow it in all directions, you get.

a three-dimensional kind of xyz axis thing here and it will grow out and become very ordered and you get that beautiful crystalline structure that catches light in certain ways and that's why it looks pretty like that because light comes in and then bounces off and is diffracted and you know you see like a diamond or something like that or ice. So anyway each one of these little dots here can be either a molecule or an ion, or a metal atom, because all of these things are crystal structures. And they're called lattice points. Atoms, molecules, or ions.

And they have different shapes based upon what they're made of. So these little shapes here, these shapes then have parameters that are described by the angle of the box or the length of the sides. So we have...

angle. We have lengths of sides, and those lengths of sides are called A, B, and C. And then the angles are called alpha, beta, and gamma.

So whatever those angles are, whatever those lengths are in relation to each other. dictates the crystal structure and how it organizes itself. Any questions about this picture? I think this is really fun, but hard at the same time because it's hard to visualize, but it's really cool.

And I have in my online, my completely online class, I have a complete discussion. I don't know, Chris, you wanted a discussion. I can give you that discussion. that you have to look up a metal, I mean, or compound, an ionic compound, look at all these things.

But we usually only focus on this one. Because when you get to these things, look at these angles. I mean, and look at these angles, 120, even though you can determine like how many are at each point.

And then you can figure out like what the density of the substance is by calculation, which is really cool. I love it. It's so fun.

So let me show you some more of that. in that previous slide was very similar to just looking at those little dots. Now we're going to grow it out and look at more space filling.

And the space filling, when you do that, you can see the lattice, the unit cell is in here. And on each one of these corners, these are corners, and actually a corner is only one eighth of that sphere. So when it's one eighth of that sphere, 1 8 8 is equal to 1. There's really only one atom in that unit cell. And you're probably asking, why do we care? The reason we care is because x-ray crystallography was developed based upon all of this understanding.

And that you can use x-rays and you shoot them through and then they diffract out of a crystal. So it's got to be crystalline because it won't diffract nicely if it isn't. And then you can determine the distances from here to here. And from here to here, just based upon...

the x-rays that come in and the x-rays that come out. It's really the coolest thing. Now, if we look at this one, now just what I said there, what do you think?

I'm going to draw this again, but this guy is in the middle. You see him? He's in the middle.

How many atoms are in this, what's called body-centered cubic unit cell? There was one in the simple cubic. How many do you think is in the body-centered? I hear a couple people saying two, and that's correct, because there's one eighth at each corner times eight equals one, and then there's plus one from the center. Perfect.

How about this is face-centered, and when we do face-centered, each one of these is called, it's in the face. So when it's in the face, and this is harder to see up there, if you take a slice of it, you slice it in half. So there's a half of a molecule. half of an atom rather, in each face. And there's one at every corner.

So how many atoms are in a face-centered cubic arrangement? Very good. I hear people saying four because there is one eighth times eight, which would equal one.

And then there are one half times six, which equals Three. Three plus one is four. Very good.

And here's the picture of it. So you can see this like sliced up and that's what I was trying to do there. And that would be simple cubic. And then this one would be body centered.

You see the one in the middle and then one eighth at each corner. And then you see, oh, I didn't do the same for face centered. Looks like we're going to do the same thing that I was just going over. Any questions on that? Now, I ask you to draw this on your fun sheet.

I say draw these things and do some calculations with it. Because if I know that there's four atoms per unit cell, yeah, coordination numbers, we're going to get to that here in a second. If I can count the number of atoms in a unit cell and I know who it is, for example, I can know the grams of one atom.

And say there's four atoms. Well, I'd multiply that by four and say, I know this is a cube. How do you get the area of a cube?

I'm not the area, I'm sorry, volume. Remember, it's length times width times height. That means I know now the volume and I can get density is mass over volume, mass of the four atoms over the volume of that cube and how many atoms were in it.

Now, coordination number, in order to know coordination number, you kind of have to know further out in the lattice. And let's talk about that. But before we get to that, let's talk about closest packing, because closest packing. is a way to, is how those atoms are organized.

So closest packing, they organize themselves like in the little indentations. So you have one layer, for example, and then the next layer goes in the indentations, and then the next one goes in the next indentation. So it kind of makes sense. It's making most efficient use of space.

So there's one, there's different ways they can do it. See, it can either go in there or it could have gone in these little holes. And that's what that one's doing. That's cubic.

And this is what we're looking at most of the time. Hexagonal does a slightly different thing. Let me see if I can whip back real quick. I wanna look at that hexagonal. Bear with me.

Don't look, I know it's looking freaky. Where is it? Don't look, sorry.

Right there, hexagonal. See that? That hexagonal means you get those 120 degrees. And you get some neat crystals. They're like grown crooked, but they're really pretty and very prismatic.

So before we can get to determining coordination numbers and understanding what that is, we have to grow the crystal out. And to grow the crystal out, they have to pack like they do in there. And there's, look at that coordination number, 12 for hexagonal and cubic close packing. Because what ends up happening is a lot of these are face centered. And when they're face centered, they pack really tightly.

So they're getting more around, hence the 12. Because there's other coordination numbers of eight and six. So let's look at those and see if we get that coordination number coming up. I'll let you guys look at these.

These are kind of hard. I have a video in content that you might want to look at that talks about all of this and zooms you in and out. It's only like six minutes long that you might want to look at to help you understand that.

And I'll put it in an announcement. I'm going to get, okay, before we get to, I'm keeping the coordination number in mind. But before we get to that, let's just look at the relationship between these things. The relationship between edge length and radius of an atom.

If it's simple cubic, see how easy that is? If you measure the edge length, you can figure out the radius of the atom. Because did you ever wonder how all the sizes of atoms were determined? This was one of the main ways to do it. Because you would shoot an X-ray through, and it would come out, and you could determine distances.

And if you would determine distances and you knew what was in here, you can figure out basically atomic radii, which is on the very small level. So these things you just want to review. If you have body-centered or face-centered, there's relationships.

And your fun sheet has some of these questions in there to ask you to use this and use Pythagorean's theorem. You know, that's what this is using. Your book goes through that. So let's just talk about this real quick here and do a quick example.

So I'll let you read that. So when it's face centered, you know how many are in the unit cell. Do you remember how many are in the unit cell if it's face centered?

Type that into the chat if you remember. It's four. That's right. You guys got it. So when it's four, you can know if it's silver.

You say it's silver. Well, what's the density of silver if there are four atoms per unit cell? And we need to know this 409. Does anybody know what P stands for? in the pm the m is meter that's right it's picometers so since it's picometers picometers is 10 to the minus 12 so we have to remember that and we have to remember that it's a cube so we can determine the volume and since we can determine the volume we can use that in our calculation so that's what we're doing there this is centimeters this was picometers cubed And now it was converted to centimeters cubed because in density, remember, D equals M over V, we want to have this, it's always in centimeters cubed, which is the same as milliliters. So if we know it is four silver atoms, then we can use the molar mass of silver to determine the mass of those four atoms.

And that's what's happening here. And then we get a number and it's going to be teeny, teeny, tiny. You see that?

It's a very small mass and then we can plug it into the density formula. So look at these numbers. See why they're so small?

Of course it makes sense. And then you're going to get the density of the silver. And sure enough, you compare that with the literature values, and sure enough, that's what the density of silver is.

So this density was determined way before the crystal structure of it was. So this was a way to verify what's on the molecular level or atomic level of silver. Any questions on this one? Any questions on how to get from picometer cubed to centimeter cubed?

Because that's one of the main stumbling blocks for this. And remembering that you have to use Avogadro's number because Avogadro's number takes it on the atomic level instead of the mole level. OK. All right.

Now, how is this done? This is this X-ray crystallography I was telling you about. You have a crystal. You shoot x-rays through it of a certain wavelength.

And if you know the wavelength, which you can adjust, and you know the angle that this thing, it's being, it's coming in at, it will come in and it will refract as it goes through the crystal. And then mathematically, that can be determined to figure out where these spots are. And I have a picture here of what a real one looks like. This is what I actually did this in my graduate studies that you can actually.

get all these spots and you know the original x-ray crystallographers actually measured they measured with rulers it was just crazy but they were the pioneers for this system this thing i don't know if you wonder there see this like whiteness here this whiteness is the actual like what's holding the crystal it's called the goniometer it's a funny it's a funny word but it's um it's the the x-ray gets it's this and that's why you see it in there. The x-rays are absorbed like your bones in an x-ray, so it's not seen, but the other ones get diffracted in this fashion. So this theta angle that's coming in here and then it's diffracted, this is through, is described through what's called Bragg's equation here, and that is depends on Bragg's law is the angle and the distance between them are related to the wavelength of the incoming x-rays.

So if you know the incoming x-ray wavelength and you know the angle for which you shoot it's coming in, you can determine the distance between the layers like this. So there was those, remember we were seeing those pictures, and now you can determine the edge length. And that edge length can give you information about the crystal structure.

Right now it's called the Bragg equation. It's very automated. I mean nobody measures anything anymore. So I'm sorry, you know, somebody asked about coordination numbers. I didn't really cover that too much.

But in that video that I'm going to put out, it shows you how to do that. It shows an atom and it shows how there's atoms around it. And as far as what coordination number it would be.

So you'd have to know the entire crystal structure and see it kind of built out to be able to see the coordination number. But face centered coordination number is 12. body centered is eight and face centered is six. I'm sorry, face centered is 12, body centered is eight and simple cubic is six. And so I'll show you that video.

I'll put that out in an announcement and let you check it out. But as far as crystal structures go, if you know through this Bragg equation, what you're, you're trying to determine the distance then between layers, assuming n equal one, and when assuming n equal one, that means you have one layer. and then you have another layer.

This is n equal 1. This would be n equal 2. And then you could say a third one. That would be n equal 3. Okay, so this is much more complex than this, but this is kind of the simple part of it. So if you know what the wavelength is, and you know what your n is, you can solve, and you know the angle, and you can solve for it.

And you can solve for the distance. And in this example, you're getting 314 picometers. Oops, sorry, I meant to just erase that.

I know this is a lot to take in because crystallography is complex stuff, but this is just kind of an introduction. But does anybody have any questions about, because I have a fun sheet question about this on there, that where you calculate some things and calculate the distance between layers in the crystal. So what's kind of cool about crystallography is that, see, here's my pen, if you can see me up there. They put the crystal on the top here. And then this thing, either the X-ray source.

can comes in and hits the crystal and then they move the crystal. The crystal gets moved in all directions and by moving it in all directions relative to the x-ray source you can take all of these measurements because you get all these spots from all different angles and I have still yet to find a really good video about it and showing how this works but that's how it's done and then you can determine the distance between them. So as far as crystal types, so that was kind of ending X-ray crystallography.

Now the different types of crystals are metallic crystals, covalent, ionic, and I think there's one more, covalent network solid. So a crystal type, I'm asking you to name that crystal type if you've looked over this information already. If it's composed of charged species, what kind of crystal do you think that is?

Is it ionic, covalent, molecular, or metallic? Type it in the chat if you think you know it. If it's ionic, if it's charged species, it's ionic. Exactly.

So that's probably what you're very familiar with because we're so used to things like sodium chloride, etc. And the anions and cations, they're different in size. It's got a very high melting point. The lattice energy is huge because it's very stable. It's pretty brittle.

And they're good electrical. They're not good conductors in a solid. But...

when you melt them, they will conduct electricity or aqueous. So that's an ionic crystal. So this is something that you might see as an ionic crystal. Can anybody tell me what type of packing that is?

Is it simple cubic or simple cubic, body-centered cubic, or face-centered cubic? By that picture, what do you think? I hear somebody, I see somebody saying body.

That's true because you see the one in the middle. It's BCC. Very good. This is called zinc sulfide.

What do you think that crystal type is? FCC? I see FCC and that's good.

But now I'm going to ask you a question here because I think this gets really hard. Because here you notice there's ones in the faces, right? So that means face centered. But you have these guys in the middle. And these guys in the middle, do you think those are the zincs or the sulfides?

based upon that picture, those little gray ones? Do you think they're the zinc or the sulfur? Now the next question I would ask myself is how can you tell? How can I tell whether they're zincs or sulfurs?

What's the difference between zinc or sulfur? Zinc is positive two, sulfur is negative two. Now do you remember from chapter eight, are cations smaller or larger than anions? This is what you would want to think about here.

Because you can see it in the picture. These gray ones are smaller than the green ones. Now, remember, when cations lose electrons, they get smaller.

When an atom gains electrons, they get bigger. So it looks to me these sulfides are the green ones. And the zinc is the middle one.

Now, let me ask you another question. If I had to count zincs in this unit cell, how many are there? That is correct.

You guys see people saying four. So I'm going to write ZN4. If you had to count sulfurs, how many are present in this unit cell? Now, remember, when you just count. But in the unit cell, you have to remember that there's only a piece of it in the unit cell.

Yeah, that's the tricky part. So now, you know, there are four. So what is the molecule?

What is the formula for zinc sulfide? It's an empirical formula, which is the lowest ratio. What is the formula? And it should match what we know and what is in the unit cell. It's ZNS.

Remember, you did. Yeah, exactly. So what about cesium chloride?

How many cesiums? And what's the one in the middle over here? Is that cesium or chlorine?

So now I'm on the other one. It's cesium. And there's one of them.

How many chlorines are them, which apparently are red? One. So it's C-S-C-L, one to one, which it should be, right?

We know that about cesium and chlorine because they are one to one. One's an alkali, one's a halogen. So we do know that.

Oh, here's another one. This one looks crazy. How many, this is calcium fluoride.

How many gray ones are in that unit cell? Now let me ask you this. What is the arrangement of the gray ones?

Is it simple cubic, body-centered, or face-centered? So I'm getting some differing opinions here because look at these guys. These silver ones are faces.

So it's face-centered, and if it's face-centered, how many gray ones, which I believe are calciums because they're smaller, They're the cation. How many are in that unit cell? Because remember, yeah, you have to only count the pieces that are in there.

And there's in the corners. Of course, you have an eighth. And in the faces, you have a half.

So that should be four. How about the green ones? How many green ones are inside that unit cell? They're actually inside. So you can just count them.

Yeah, F8. And what's the molecular, what's the empirical formula? See that? It's the lowest ratio. And that it should match every time because what chemistry is doing is studying matter and we're zooming in and we're trying to figure out the ratio and what's actually in there.

It's like studying anatomy of a person, except we're studying the anatomy of the universe. So that's why it's such a big subject. Anyway, anybody have any questions about what I just did there? Fun stuff. Sorry, I've got a quick question.

Can you just explain again how you determine how many calcium atoms or ions that I guess in this case there are in that structure? Right. So what I'm understanding is that like these structures with their little spheres are not exactly like that's not exactly the location of a of an ion of a ion or an atom or is it?

It's the location, but not all of it is actually in this unit cell. What does that mean? It means like if we look at, let me go back to this picture.

Sorry, maybe I missed a little something. No, it's hard. I think it's hard personally.

I'm just right there. So when you look at these pictures here and you say, when we do that counting, you have to take into account that there's only a certain portion of that in that unit cell. So on the picture that we were just looking at, the calcium's were face centered. Oh, okay. Okay.

So there, so when they're talking about unit cell, they're taking essentially like this cubic measurement sort of thing. Yes. And they're slicing it.

It's being sliced. Okay. So when I show you that, when I share that video with you and you watch it, it'll, it'll make a lot of sense after we've talked about it. And I'll share another link called ChemTube3D.

It is like the fabulous thing, ChemTube 3D. And it enables you to click on one of these unit cells and rotate it and really see. It doesn't really slice, but it'll let you do one unit cell and you can move it around and to determine the number of atoms inside. Because some of them get a little bit more complex like that last one, that calcium fluoride we were just looking at.

Yeah. Yeah. OK.

Thank you. OK. Anybody else have a question about that? You just have to start talking because I really can't see your hand.

If you have a question, I was just going to ask it, it looks like the ratios are. By the valence electrons, like, how many balance electrons each like. The calcium and the, and the flooring for instance.

Like, calcium has for the electrons, therefore. And pouring needs 1 more because it's really close. So that's why there's 2. fluorine per calcium. Does that have something to do with how you're counting all of this? Well, I think it more has to do with the ratio.

Like you're saying, it's a one to two ratio. Calcium has two electrons, two valence electrons, you know, and fluorine has your seven, right? So you only have, it needs one and then it needs another one, right?

So these crystal structures, what we're studying right now is ionic and the ionic compounds, they build by size. So they have these different structures because this one's too big, so they can't snuggle right. So they have to have a certain type of structure.

We can have molecules can do this too, where you'll have a water molecule at each little place. But the water molecule might look like this, and part of it's over here, and part of it's in the unit cell. So molecular crystals do a little... They they're following face centered, body centered, simple, but they only a part of a molecule, for example, might be in there.

So it does have an effect, like you're saying, that the valence electrons do mean something because as far as the ratio is concerned and as far as the size and then which structure they take based upon their sizes. Because let me ask you this. Isn't cesium huge?

Think about where cesium is on the periodic table. He's down the lower left. He's monstrous. So he takes up that.

big center area and then you got these chlorines you know around him whereas calcium is not as big as cesium so look how many they can fit into the unit cell does that make sense and zinc is a little larger yeah so well so i guess um i just can't this is just ionic so we would still have to do like the counting and everything with the um you the crystal structures and figure out which crystal it is to, and, and how many, um, I guess how many like angles there are for regular molecules, because they're not going to follow like the same rules as an ionic structure. Well, you know, I think I'm just messing up left and right. I can't stand it.

I'm going to have to show you this chem tube 3d thing, because once you see this, it really helps. I think, um, to visualize what we're talking about here. So let me just run over and I can, I know exactly I've been using this site for quite a while now. I used it to create a discussion based upon counting of atoms and like this, because it was hard.

So ChemTube3D, it's out of University of Sheffield in England, and it's just the coolest thing. And you go here and then you can go to any one of these structures. Can you guys see what I'm looking at?

Okay. And, and say for, I just want to do a simple one, cubic close packed. We just learned about that.

And then it'll bring up a structure for me. Hopefully it must be thinking. Can you guys see that moving around there? Okay.

Now it looks like to me, is this, what, what is this? Is this, is a simple cubic body centered or face centered? What do you think?

That's right, it's face. So you're seeing it. So there's even more to this. It gets crazy. You can click on this stuff and it shows you like other stuff outside of that unit cell.

But what I wanted to show you was kind of that what you were asking about the molecules, because I feel like let's see if we can find some molecules in here. This stuff has so much good stuff that you can't stop looking. You know, I don't know.

It does have water in here. Search. water. You can do a search in here and you can search water.

And we want some ice water in the solid state. And then you can look at this a little bit. Now, I don't know. I don't feel like this is helping me.

I feel like this is just a freaking mess to try to figure out what it is looking. I know I can't see the unit cell. I would need a unit cell.

Sorry about that. I shouldn't have showed you that. Well, I don't know.

I'm not really finding. This only has a lot of the inorganic things that we were just learning about, like that zinc blend, zinc sulfide. We just counted on the PowerPoint.

That's right here. So you can look at it, and it gives you this structure. And you can click on the CCP, and it gives you the unit cell only.

And then you can move it around and go, oh, I see there's four gray ones in there. And that's what we just did. And I see that there's, you know, it's face centered.

So I see. Some of the yellow ones in the faces and the yellow ones on the corners. So that was ZN4S4. Remember, we saw.

So, oh, there it is. See, ZNS4. And sometimes you can even see coordination.

Remember, somebody was asking about coordination number. And you can see coordination. So you can click on it and it starts giving you some weird things.

See that? There's things in there. It's talking about the coordination around each atom.

So it gives it like a tetrahedron. Because that's how it's organized. There's a lot to this, I know.

But anyway, I just had to show you that because it was kind of cool. Anybody have any questions on that before I get out of here? Your book has a question, and I encourage you to look at it because it's helpful.

It has a question about the unit cell and number of molecules of iodine, and iodine's a molecular solid. So where we are right now is we are talking about the different types. Of solids, and those solids, we just did ionic and we went through the different.

Counting and you really only find that in this ionic situation. We just counted all of these. In this book, it doesn't really show it too much for the other types. Now let's name that crystal type.

We want to move on to another crystal type. If the lattice points are occupied by atoms only. Thanks, Shannon. Yeah, what time is it? We need to get out of here soon.

I got about 10 more minutes. If what crystal type is it, do you think? Type in the chat if you know. If the lattice points are occupied by atoms. So they're not ions.

They're not molecules. They're atoms. And they're held to death.

together by covalent bonds. What else? It has a high melting point.

And it's a poor, oh, see, this gives it away right here. If it's a poor conductor of heat and electricity, could that be a metal? Yes or no? No.

So it can't be a metal. It's not a metallic. It's called a covalent network solid. So a covalent network solid or a net.

network solid or covalent solid. I mean, you hear different books call it different things, but I always, I learned it as a covalent network solid because it was networked, but it was all covalent bonds, but it was a bad conductor because all those electrons are localized. They can't move.

So that's why it takes that form. Now this also has a certain unit cell. Look at this diamond.

I don't know if you can see a tetrahedron right there. So it, on that chem tube 3D. there is a diamond structure for you to look at and you can move it around. So that is called a covalent network solid. Now we're at another one.

What type of crystal is this? If lattice points are molecules held together by intermolecular forces, low melting point. And so again, a poor conductor, what type of solid do you think that is? Hopefully you have your book next.

Yo, somebody does. Molecular. Very good. That's a molecular solid.

And water is a molecular solid. So is benzene. It's funny though, benzene, solid benzene sinks, solid water floats. Okay, one more here, lattice points occupied by metal atoms.

Well, that gives that one away, right? That means it must be a metallic solid. And sure enough, it conducts electricity. It has a high, it can have hard, it can be hard or soft, the metal.

It can have a low or high, just depends on the metal. And this is like a sea of electrons. And these electrons just swim around here.

And that means that it's a good conductor. Okay. So the crystal structure of metals has all these different types. You have the body-centered, the face-centered, the hexagonal. So all these metals take those different unit cells.

And here's some cool crystal structures of different substances. And what do you think that one in the middle is there? Do you think that's an ionic compound, a metallic?

substance or covalent. Yeah, it's metallic. That's easy to see because it's like, it looks like a metal, right?

Which is shiny, right? And it's lustrous and that type of thing. This first one is copper sulfate, hexahydrate, pentahydrate, I'm sorry.

And I think there's an example of them right over here. Maybe not. It doesn't give that to you. No, but it's an ionic solid. It has water molecules in it.

This one, metallic, this one is sulfur. So sulfur, that yellow substance is sulfur. So that one is under a covalent solid. It's not up there as an example, though. Quartz is a covalent solid.

So anyway, and this is just another example of that. And these are posted for you, these PowerPoints, if you want to look at them. But I think we're kind of done for today.

Does anybody have any questions before we quit? Because I know that this is a lot to take in, but chemistry is a lot to take in, I think. How do you determine coordination numbers from just like the structure, like the simple cubic and the face centered or the body centered? I don't think you really can. Like you say, how do you determine?

It's like saying. What I always akin this to is I always say we're visiting the chemical zoo. And what's going on in the chemical zoo? When I go visit zebras, they have stripes.

And I just have to know that. When I have a crystal structure that's body centered, I just kind of have to know that it has a certain coordination number. Now, can you count it? If you saw a drawing? Can you guys hear me?

Yes. Okay, because my whole thing is locked up. I mean, my PowerPoint has locked up and stuff.

And it's spinning. But either way, you can look at that and you could count if you saw an extended lattice work. And you could say, oh, there's a sodium and I count six around. Well, that means it's got a coordination number of six and it must be a simple cubic.

So you would have to have a picture of it is what I'm saying. And you'd have to be able to count it that way. Now, you could also just memorize that simple cubic has six, body centered is eight, face centered is 12. And if it does indeed have that, that must be a piece of evidence that it's face centered or whatever.

Did that make sense? Yep. Thank you.

Okay, but that's a good question. I'm going to go see if I can find after I log off here. Find a picture of what you're talking about and point at them and just count them.

And that video that I'm going to share with you shows it in 3D almost. So it would help you. All right. Well, I'm locked up. My PowerPoint won't stop sharing.

So I'm just going to say goodbye. I don't even know if I can quit the meeting. So you guys are just going to have to leave because it's locked up on my end.

Okay, so I'm going to stop talking here and I'll see you guys on Monday. Bye. Thank you. Have a good one.

That's sweet.

Transcript for:Understanding Intermolecular Forces and Crystals

Transcript for:
Understanding Intermolecular Forces and Crystals