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Oregon State University Ecampus Online Chemistry Lecture - Relative Melting Point

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Oh, hi. Up on this board I've gone ahead and compartmentalized some compounds. What we have on the far left are the inert gases, group 18s, such as helium neon, and argon. You can look at a periodic table and see that these are the inert gases. I've got molecules listed, which are a combination of nonmetals only, and broken it up into Roman numeral I,II, and III. Ionic compounds, which are a combination of metal and nonmetal, and totally think of these things as being solid, where what's keeping them together are charges. So we have oppositely charged magnets. Over on the right side, completely new topic for us called network covalent compounds. There are not many of these known, and I would consider these to be the most common three: graphite, diamonds, and quartz. Make a couple of notes here. Graphite is 100% pure carbon. So let me write down pure carbon. And graphite appears as sheets. Let me hold up a model. And this is going to be a sideways model of graphite. Now ignore these little metal rods because they are not supposed to be there; they're part of the model itself. These are not to be taken as bonds. What we have is, if I go ahead and turn it this way, we have hexagonal, hexagon, 6-membered carbon wings. And these wings are fused together. Now I'm going to just take a piece of common plastic, here, and hold it up this way. Saying please go ahead and envision that this top layer is a sheet. It goes on for a long ways in two dimensions. Now here's another sheet, and they're not bonded together. This sheet and the top sheet are not bonded together, those are just pillars keeping the bottle together. There's very little attraction between this sheet and the middle one. And then one more here. Now graphite doesn't exist as three sheets; it exists as a bunch of sheets. But let me take three of these common plastic sheets and tell you that to think of them as being sheets of carbon graphite. The sheets are very, very slick over one another. Very nice model because there is very little interaction between this sheet and this sheet. Carbon graphite makes for a nice lubricant. Locksmiths have it as a crushed black powder and they'll squeeze it out of a tube into a lock so your key slides in nice and easy. If you take carbon graphite and crush it, kind of press it down into a rod, you can insert it into a tube of wood, make yourself a pencil. If you take carbon graphite and crunch into like a block you can have a charcoal briquette. If you take it and put a little bit of a starter in there- And what I have over here is, I've got rifle powder. Rifle powder, or known as gun powder, and this is one really old can that we found at a garage sale, this is probably like a surplus from World War II that's been boxed up and sold to people. Tiny little black pellets of carbon graphite and a starter. That starter gets things going. So this isn't like a barbeque where I have to start it with paper or bring in an exterior fluid. There's a little bit of starter on here and it's giving off a nice amount of heat. The reason it works as like a projectile launcher is because envision this like in a cannon, pressure building up and then shooting a projectile. So we've got carbon graphite, it's known as a network covalent compound because it's a network. It's not one atom. It's not sixty atoms. It's not a hundred atoms. It's a whole bunch, thousands, trillions, trillions, trillions, of atoms. Now this has a very high melting point. So already on this chart let me go ahead and say things to the right, these network covalent compounds the clear winners. These are going to have very high melting points. Now in order to melt graphite you need to put in a tremendous amount of heat. I'm not going to list the temperature, but tell you it's in the range of 4000 centigrade, the temperature of the outside of the sun, just incredible heat. And the reason it's so high is because in order to melt you need to loosen these carbon atoms. You need them to flow over one another, make them melt. And in order to loosen them all up imagine all of these bonds, the energy there that you need to overcome. So you have to put in a tremendous amount of heat to get all of those bonds. That's not true with something, say, of an inert gas. Inert gases just exist as points in space, with very little interaction. These have a tremendous amount of interaction. Next down the list is diamond. And diamond, oh diamond girl's best friend, hardest substance naturally occurring, all these little old-time sayings. We have a box of fake diamonds here. Diamonds are clear, 100% pure carbon. And the deal on diamonds is that we have we have SP3 hydrolyzed carbons. We have one carbon surrounded by four neighbors. And that goes on and on and on. So it's a nice model, but it's not very good in the sense that it's small. It's supposed to be like miles long. showing you trillions upon trillions upon trillions of carbon atoms. And again why is it hard, why does it melt at a high temperature? It's because in order to like break it or interrupt these interactions you have to put in a tremendous amount of energy. So there we have it. Pure carbon graphite, pure carbon diamond. Let me write pure carbon next to diamond, but also point out the third one is quartz. So next to diamond pure carbon as well. And people have said things like "well graphite is stable and what the earth has done over years is heated it up and pressurized it into this configuration, diamonds." Quartz is not pure carbon, but it is a network covalent compound. Quartz is made of SiO2 units. One could say that it's similar in structure to diamond, in that what we have is silicon surrounded by four oxygens. They're not flat as in the board they make more or less like pyramids. Where we have three oxygens down below and one oxygen on top. Go ahead and sneak a silicon in the middle there and you can show the four bonds if you'd like. And we have these triangle based pyramids. Now these don't stop going "okay this is all we have" because it doesn't even match the formula, SiO2. Si, there's four oxygens, but instead what's happening is we have these neighbors, and it goes on in three dimensions. So lets explain the name. Network because we have 3-dimensional, of in the case of graphite 2-dimensional, but the idea is ongoing structure. Covalent because it turns out these are made out of nonmetals. So we don't have positive or negative charges we have covalent or sharing of electrons bonds. These melt at extremely high temperature, call it 4000 centigrade. Way over on the left, the inert gases we don't have that at all. Let me list a few of them. These inert gases should be though of as being specks or discrete particles. They don't have neighbors. If they did have neighbors they would be near each other and the interactions would have to be strong enough to keep them together, but at room temperature and pressure we don't observe that. If we take helium down to extremely low temperatures, like 4 Kelvin, they will go ahead and get together. You've taken out all the energy. They're not going to be zooming around in the gas phase. And we can show a slight attraction between them. But where we live, certainly already melted and then boiled. They are gases at room temperature and pressure. Molecules, remembering that these are not metals, fall into three different categories to separate low melting points to high melting points. Type I are going to be your nonpolar. These nonpolar molecules, just like the inert gases, have very little attraction between them. We call them dispersion or London forces. Let me make a little picture up on the board of dispersion or London forces. These London forces, or dispersion forces, are found in inert gases and these nonpolar molecules. Let me show a methane molecule as just written out as CH4. And another one as CH4. These are tetrahedral, nonpolar, so there is no like positive and negative end like a polar molecule, so we have very little attraction. This dispersion, or London force, that's present here, or in just the inert gases, is said to have come about because we have a distribution of electrons. This molecule has electrons, whether localized or delocalized, and at any given moment the electrons are not perfectly balanced. And so this will be offset at one moment, this will be offset at one moment. We could have a slight attractive force. It's extremely weak. Type II molecules- We're going to go ahead and see a polar effect. And so if we have a molecule that is polar, and let me hold one up, that's going to be methanol, in this case. Water is polar, it undergoes the same thing. With methanol what we have is we have a negative end over here because of the oxygen and two lone pairs of electrons and a rather positive end over here without an electron pig. So I'm going to point my hand going toward this side. This is where the electrons are. Think of this as a magnet of sorts. Negative end, positive end. So a neighbor could come up next to it and it's not going to align itself this way, with two negative ends towards each other, there's a preferred orientation where we're going to have positive, negative, positive, negative. They're going to line up. So they're going to act like magnets and line up. There's a nice attraction. So we call this a Roman numeral II, or polar attraction. Lets go ahead and make a little note here that this happens with water, and how about this one? Very similar to methane, nice and balanced CH4, except it's no longer nice and balanced. CH3, and somebody's replaced one of the hydrogens with a fluorine. So we have a polar molecule. Type III, and this is going to be the strongest attraction that we see in molecules, is called hydrogen bonding. Hydrogen bonding occurs when we have extremely electronegative element. Let me list those. When we have oxygen, fluorine, or nitrogen present and off of one of these is a hydrogen atom. If we have hydrogen attached to one of the big three electronegative elements we find that this is an extremely polar bond. The electrons are pulled toward the oxygen, fluorine, or nitrogen and taken away from the hydrogen. Let me show you the effect using water of this. I'm going to draw a water molecule. And on the left I don't want to subtract, so I'm going to make a bond there in the form of a single line. Let me make this bond in the form of a pair of electrons. We know that this pair of electrons is, well, pulled towards the oxygen because oxygen is very electronegative compared to hydrogen. Now that's going to mean that this oxygen is rather electronegative. It has the electrons pulled towards it. And this hydrogen is having the electrons ripped away from it. So we put a little note there that it's electropositive. Now the result is, this is one water molecule. Say another water molecule is positioned right near it, it's going to orient itself like this. Pair of electrons being negative are highly attracted to the positive right there. And we form a hydrogen bond. Hydrogen bonds are always denoted with dotted lines like this. And it's a very strong attraction. All three of these attractions, considered for molecules, dispersion, or London forces, are actually in inert gases. They're in nonpolar molecules. They're also in polar molecules, but very insignificant compared to the polar effect. And it's in hydrogen bonding molecules, but insignificant compared to hydrogen bonding. Polar occurs when you have a polar molecule, and hydrogen bonding occurs when you have hydrogen attached directly to an electronegative element. As we go to the right the melting and the boiling points go ahead and increase. Instead of writing ionic compounds you are free to write ionic solids because these are now solids at room temperature and pressure. The weakest would be a +1,-1 metal nonmetal combination. Think of these as being weak magnets, +1 -1. Over here we get the strongest, +2 -2. These are stronger magnets, force of attraction is greater. So calcium oxide melts at a higher temperature than sodium chloride because of the strength of the charges. Two and two instead of one and one. In the middle here, in intermediate, +2 -1, is in between sodium chloride and calcium oxide. Over here upon first inspection these look to be identical. +1 -1, +1 -1 Their melting points are extremely close, but we find that sodium fluoride has a melting point that's just a little bit higher than sodium chloride. If you will, since these have the same charges, +1 -1 +1 -1, it's as though we think of this as being in need of a tiebreaker going "well which one is going to win?" Size wins. Here it goes. Sodium + and sodium + are present in both of these. So lets put sodium + down . They are going to be sodium ions of exactly the same size. The chloride is matched up with this one. And let me exaggerate and show chloride as being really large. And fluoride. Let me exaggerate and make it pretty small compared to the chloride. The idea is because the periodic table of the elements shows chlorine as being lower it's going to be bigger than fluorine. Take a look at a periodic table. Fluorine sits above chlorine in the periodic table in group 7. Now chlorine is really big compared to fluorine, which means that the center of the charge of chlorine is going to be further away from sodium than the center of charge in the fluorine. The analogy of the magnets works really nice. Two magnets that are far apart- that's not too much force. It's easy to pull them apart. Two magnets that are closer, greater force, harder to pull apart. So we have ourselves a trend: inert gases, Type I,Type II, Type III interactions, we call these intermolecular forces. Forces between the molecules. Nonpolar gives us dispersion, or also known as London forces. Polar are stronger, higher melting points. Hydrogen bonding are the strongest out of these. Hydrogen bonding gives us the strongest intermolecular force. And that's present with oxygen, fluorine, or nitrogen if we have a hydrogen bonded to it. Higher in temperature are the ionic solids charges +1 -1 +2 -2. And you can have things like +3 -3, even stronger. When you need a tiebreaker the smaller ion has the closer charge. More energy to bust this apart. And then finally your network covalents, very high in temperature.

Video Details

Duration: 16 minutes and 1 second
Country: United States
Language: English
Genre: None
Views: 240
Posted by: umarket on Sep 24, 2009

Videos were taken from an Online Chemistry course provided by Ecampus at Oregon State University. The professor is Dr. Richard Nafshun.

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