 I'll be handing back the quizzes that we took a week and a half ago today at the end of class. I should be done with all of the grading probably in the next couple of days. You should be able to access your midterm grade. I don't know when they'll post them, but I should be done with them by I think Tuesday or Wednesday. If you're super interested in your midterm grade, it should be coming out soon. No, it won't count your lab grade. In fact, your midterm grade will be deflated relative to your final grade because you're not counting your lab stuff, which is usually for most people who post them. You won't have the curve on the OWL assignments and you won't have the curve on the quizzes. Because I can't do that until the end. But you will have the curves on the two tests. Hopefully we'll get the second exam in there. I'm not sure though. I'll tell you next time if it's in there or not. Okay, so that's the deal with grades. Are there any other questions about that type of stuff? Okay, cool. So let's just finish up chapter 6 today and then start chapter 7. There was a couple of computer problems this morning. Sorry that the chapter 7 slides didn't get up, but they'll be up right after class. Anyways, we talked about surface tension and resulting in beading. When you look in the water you can see this really clearly. You can see little beads there. Not the best surface for it, but it's the attraction of surface molecules. Actually, the attraction of the molecules underneath the surface molecules that actually are pulling them downward giving you kind of a beaded shape. So that's why raindrops are spherical in formation because they kind of want to get away from the surrounding atmosphere because if you recall water is polar, remember, and air being made up of mostly nitrogen and oxygen and you guys know how to build those now. But if you look at nitrogen and oxygen and carbon dioxide, can anybody tell me that? Is this polar? No, is this polar? No, is this polar? It tries to get away from the non-polar stuff as much as possible. So of course water is the most common liquid that we're dealing with in this class. So it really will benefit you to start thinking about the properties of invention. And then we talked about surfactants, so substance which decreases the surface tension of water for example. So this kind of just gets in between the water molecules and disrupts them. So vapor pressure, let's see if we have a good picture of vapor pressure. Well it doesn't look like it on the slide, but we can draw something. You can imagine if you've got some sort of blast that's connected to a vacuum and there's nothing in the blast or just air in the blast and if you evacuate that blast, you sucked out all the air then there's nothing less than the blast, right? And then you can imagine some liquid, okay? So the only thing that's in that thing is the liquid, okay? So what will happen is all the molecules won't stay in the liquid phase. What will happen is some of them to occupy the space up here will go from the liquid to the vapor phase like this. And in fact this happens whether you've evacuated the blast or not it just happens in a more dramatic fashion when the plastic is evacuated because there's no air occupying that space up there, okay? So what you'll see is after time the surface level of the liquid will actually decrease, okay? Because some of the molecules went up into the gaseous phase and there'll be now pressure on the liquid, okay? Whereas there wasn't before because we removed all the molecules, okay? So this pressure coming back down on the liquid is known as the vapor pressure, okay? So let's go ahead and describe what we drew there. Place water in a sealed container. Both liquid water and water vapor will exist in the container. How does this happen below the boiling point? Well, the temperature is too low for the boiling inversion. It's just that some of those molecules have enough energy to get over that phase transition, okay? So because the liquid molecules are in continuous motion with their average kinetic energy directly proportioned to the Kelvin temperature so that's the average as not the energy of the individual particles, okay? So since the average kinetic energy of those molecules in the liquid phase is lower than what they would be at 100 degrees Celsius which would let them boil, right? The average gives you that the bulk of them will still stay in the liquid phase, okay? But some of them, very few of them, will have enough energy to get out of that liquid phase and go up into the vapor phase, thus causing this vapor pressure. So the average molecular kinetic energy increases with temperature as you would imagine, right? Energy and temperature are essentially the same thing, okay? So some high energy molecules have sufficient energy to escape the liquid phase and even at cold temperatures some molecules can be converted. So if you've ever set a glass of water out for, I don't know, days at a time, right? Eventually it'll evaporate, all of the water will leave, okay? But it never got above 212 degrees Fahrenheit in your house, right? If it did, you'd probably be dead, you know? But that's the temperature that water boils at, you know? So how did that happen? Because you would expect that they would only leave when they get enough energy. Well, what happens is that every individual particle or every individual molecule of water has its own energy spectrum, if you will, right? And if that molecule can get above the energy needed, that would be equivalent to getting it to 212 degrees Fahrenheit, then it can go into the vapor phase, okay? So even at cold temperatures some of these molecules can be converted and over time, of course, they all will eventually. And you can see here's a pretty good picture of the average kinetic energy of the molecules. So you see that here when it's cold, most of the molecules, the average number of molecules, the highest peak, which would be where most of the molecules are, have kind of an average kinetic energy, but you can see the whole of all of the molecules. So some of them that have a lot of energy are way up here, but there's not very many of them, right? But if we increase that temperature, right, the average kinetic energy increases as well, right? But the distribution of molecules increases too. So you've got more molecules that are at the higher range here, right? So that's why it just boils faster. So as you can imagine, molecules in the vapor phase can lose energy and convert it back into the liquid phase. So these two processes are known as evaporation. That's the process of conversion of the liquid to a gas at a temperature too low to boil, okay? So we call that evaporation. And condensation is the conversion of the gas back to the liquid. So you can see here, we actually give off energy from a gas to a liquid, okay? If you recall the acetone experiment that we did in lab where I poured some acetone in the beaker and I put the two beakers side by side, the one without acetone and the one with acetone, and then you felt it and it was colder, if you recall this, right? It's because of this process here, the opposite of this process, right? So going from a gas to a liquid, that's an exothermic process, so it gives off energy so you'll feel it get warmer. But if you recall, when we went from a liquid to a gas, we lost energy, right? So it's the opposite of that. So it's an exothermic process condensation, okay? Evaporation is an endothermic process, okay? And this is essentially what we've described here. Okay, so this is an evacuated beaker with some sort of petri dish on it, right? And so we just put our liquid in there. So you can see the rate of evaporation is very high. So these arrows depict, red depicts evaporation, blue depicts condensation, at time zero. It's only evaporation and it's very high. Time one, right? Some starts coming back down, okay? Time two, a little bit more, and then time infinity or whatever. It's equal, okay? So you're having equal amounts evaporate and condense, okay? So that beaker that shows that there's an equal condensation and evaporation, this beaker is known as, well, it's called that it's at equilibrium, okay? So we say that this beaker is at equilibrium because there's just the rate of evaporating to the rate of condensation is equivalent, okay? So if you recall the acetone beaker. So in the case of the liquid gas equilibrium, the point at which vaporization equals condensation is known as the vapor pressure of that liquid, okay? So when you're at equilibrium, then you can feel how much pressure that liquid or the gas on top of the liquid is giving off and that pressure is known as the vapor pressure. Not anytime here, here, here. We have to wait until we're at equilibrium to figure out what the vapor pressure is. We'll be talking about vapor pressure more later, okay? I just wanted to introduce it. So boiling point, so remember, evaporation and boiling are two different things. Evaporation occurs below the boiling point, okay? Boiling point is the temperature at which the vapor pressure of the liquid becomes equal to the atmospheric pressure, okay? So the normal boiling point is the temperature at which the vapor pressure of the liquid is equal to 1 atm, why 1 atm? Because that's the normal atmospheric pressure. So what happens when you go to a mountain where the atmospheric pressure is lower than 1 atm? Does anybody know? Has anybody gone up to a very high mountain and tried to boil water or go anywhere to take the temperature of it? What you find is that it decreases the temperature, okay? And if you were underwater or whatever, it would increase the temperature, you know? That it would take to boil water. In fact, I think if you go to like Mount Everest and you try to boil water at the top of Mount Everest, you could boil water at like 78 degrees Celsius as opposed to 100 degrees Celsius. Boiling point, what? It feels true. Yeah, it feels true. Boiling point is dependent on the intermolecular forces. So polar molecules have higher boiling points than non-polar molecules. It's because polar molecules are like magnets. They like to stick together. So it's harder to like pull them apart. And non-polar molecules, they don't like to stick together. So it's not very hard to pull them apart. Okay, so let's talk about some of these, what we call van der Waals forces in liquids. Well, the physical properties of liquids are explained in terms of their intermolecular forces. Remember what we were talking about. We were just talking about polar molecules that stick to each other. Well, let's look at what are actually doing that. Kind of these two sticking together. Okay, so it's not a bond, but it's an intermolecular force. And there's different types of intermolecular forces. They're collectively known as these van der Waals forces. Well, these van der Waals forces are different a little bit different than this. This one is actually a hydrogen bond. Very strong. But it's not a bond, even though it's called a hydrogen bond. Okay? It's a misnomer. This is a strong intermolecular interaction. These van der Waals forces, the hydrogen bond is a very strong dipole-dipole interaction. So it's the attractive forces between these two polar molecules. So this is a hydrogen bond because it's between H and O and or S. Okay? So if you see this H between one of these things, it's called a hydrogen bond. Okay? If I was just looking at kind of a polar molecule like acetone, so it doesn't hydrogen bond. It doesn't hydrogen bond. But if you took the difference of electronegativities between these two atoms, what you would find is that the overall dipole moment is like that. Okay? So what does that mean? This is the negative side. This is the positive side. Right? So what you'll find is when acetone is in the liquid phase, right, you'll have two acetone and I'm just going to draw a skeletal structure here. Two acetone molecules kind of align like this. Okay? But not actually having an hydrogen bond. But this is a dipole-dipole interaction because this is a dipole, right, that thing. So dipole-dipole interaction. This isn't a formal hydrogen bond, okay? So it's less strong than a hydrogen bond. So this one is the strongest intermolecular force when you have a hydrogen bonded to those. Second strongest would be between polar molecules, right, that don't have this sort of situation. And that's called a dipole-dipole. And then the last one is these London forces. It's named after some guy named London. So as the electrons are in continuous molecule, a non-polar molecule could have an instantaneous dipole. So what does that mean? So if we look at a non-polar molecule, so notice this is a polar molecule, this is a very polar molecule. Let's look at a non-polar molecule. Okay, so in carbon dioxide or in nitrogen or oxygen or anything, it doesn't help us really as much to look at the structure here, but let's kind of idealize the structure as a cartoon. Okay, so if we were a little person, small enough to be able to see a carbon dioxide molecule, it wouldn't look like this to us. In fact, it would look more kind of like a curvy hot dog or something like that, okay? So where the atoms are, where it kind of bumps out, let's blow that up. Carbon dioxide molecule, right? Carbon dioxide has a bunch of electrons flowing around through a lot, right, however many carbon has, plus as many as each oxygen has, right? And those electrons, you know, are able to flow throughout the molecule. I know we've talked about it concentrating around an atom, right? But there are a probability of these electrons to be on this side of the molecule or on this side of the molecule. What you'll find is that when two carbon dioxide molecules or two non-polar molecules kind of come into contact with each other, what they'll do is they'll have their electrons like flowing, flowing, flowing, right? And then say they're kind of concentrated over here, right? The electron has negative charges that I'm showing. And so what happens is this carbon dioxide molecule also has its electrons flowing around here, right? But when they come into contact with each other, it feels that this one has a concentration of them here and kind of pushes this on to the other side. It doesn't like it, okay? So negative and negative don't like each other. This is known when this happens, so now, right, if that pushed its concentration over here, so now this side of the molecule is more positively charged, right? So now the kind is very, very, very, very slightly attracted to each other, even though they're nonpolar, okay? But very, very slightly. That's why carbon dioxide is so hard to get into the solid phase, okay? Like dry ice, you know, melt or sublime, I guess, much lower than zero degrees. In fact, if you put your hand on dry ice, you'd probably give yourself a severe burn, right? It's because of this interaction here that it's very hard to keep it in the solid phase, okay? Because this is not a very strong interaction. Remember how I said these are like, these guys are like those magnets that you have together, or, well, I guess ionic ones are more like those magnets that you can't like pull, you can't even pull them apart, right? But these ones would be more like, you got these two magnets, they're far away from each other, and you let them go and they smack on each other. These ones would be like two refrigerator magnets, right? Had to like force together, whatever. That's these ones. These also have low boiling points, not super low. These have high boiling points relative to their molecular weight. These have boiling points relative to molecular weight, and these have very, very, very, very low boiling points relative to their molecular weight. In fact, like liquid nitrogen's boiling point is negative 78 degrees Celsius, so very cool. This dry ice is negative 30 degrees Celsius, okay? And remember, water's zero degrees Celsius. That's a melting point, right? This is a boiling point. So this thing here, this could happen. This is called an induced dipole here. It's inducing this to occur. And this one's called an instantaneous. Moving, moving, and at any instant it can be here, here, here, here, here. And when the instantaneous gets near another one, it'll induce the other one to have its own dipole. So London forces, this is called London forces here. If the electrons are in constant motion, the non-polar molecule could have an instantaneous dipole induce the dipole here. So London forces exist between all molecules, but it's the only attractive force between non-polar atoms and four molecules, okay? This is why non-polar stuff has very, very, very low freezing points, very, very low melting points. So there's your instantaneous dipole electrons can be at an instant a range in such a way that they have a dipole and the temporary dipole interacts with other temporary dipoles like here to cause attraction. There's hydrogen bonding in a three-dimensional array. You can see it much better. So it's not really considered a Van der Waals force like I alluded to earlier. It's really a very, very strong dipole-dipole interaction. So it's a special type of dipole-dipole? Yeah, so you... this really causes the expected boiling point and melting point to be much higher. If you think about the molecular weight of water, right, this boiling point is 100 degrees Celsius. The molecular weight is 18.02. This one is negative 30 degrees Celsius, right? And its molecular weight is what? 16 plus 16, that's 32 plus 12, right? And that's 44.01. One of your things we thought, well, they're probably... are they usually going to have higher boiling points than lighter things, right? Just like we talked about throwing a bowling ball up into the air or throwing a tennis ball up into the air, which one's easier, right? Throwing the tennis ball up into the air because it doesn't take as much energy, right? But if you think, right, then why is this thing that, you know, over twice as light got, you know, a 170-degree temperature difference in the boiling point? You wouldn't expect something like that, right? It's because of these hydrogen bonds. Very important things that happen. Okay, so... and these are some of the molecules that... so hydrogen bonding has an extremely important influence on the behavior of many biological systems like every biological system since every biological system is compulsive. About 70% water, at least, you know, and water contains these hydrogen bonds, right? These other molecules, ammonia, especially, is another one that you find in biological systems. HF is not something that you find in biological systems, but it's the only molecule that hydrogen bonds that has a fluorine. Okay, so let's talk about the solid phase now. We'll leave the liquid phase behind us. The solid phase is characterized by high density, of course. You've got to rock. You've got to rock. You put it in water or the things, right? High density. Definite shape, right? You definitely have this shape with these crystals here that's independent of the container. And small compressibility, you can imagine, trying to take a hammer and beat a piece of wood or something. It won't compress very much. And very small thermal expansion. You can imagine doing the same thing, right? Trying to put something like wood into the oven without burning it and trying to expand it. It won't expand very much relative to like gases or all liquids don't have a very big thermal expansion. So when we're looking at the particles in solid, of course those were the macromolecular properties. Let's look at the properties of the particles in the solid. So they're highly organized and very defined. The melting point just depends on the strength of the attractive forces. So polar solids melt at higher temperature than non-polar solids. A crystalline solid is a regular restraining structure. So both the sodium chloride crystal here and the diamond crystal. So this is just carbon, Adam. Those are regular repeating structures. An amorphous solid is more like sulfur that we saw in the picture before. It doesn't have a regular repeat. No organized structure. Okay, so there's types of crystalline solids that I'd like you to know. Two of them I've got models of right here on the board and they're the same two that are right there. So ionic solids, we know all this stuff about ionic solids already. They're held together by electrostatic forces, the positive and negative charges of the ions. They're high melting point and boiling point, right? If you recall, they're about 800, 700, 800 to about, you know, 1400 degrees Celsius to melt, right? So very, very high melting. Hard and brittle, right? If you've ever gotten a big chunk of salt and you tried to smash it with a hammer and break all over the book, it would. And if they dissolve in water, well, we haven't really... We talked a little bit about electrolytes about how we can carry the charge on the ions. We'll talk more about this electrolyte in the next chapter, okay? But if it dissolves in water, we call it an electrolyte. NACL is an example of that. So there's NACL. There's another picture of NACL. And there's some... So this is the chemical structure of it. That's the macromolecular structure of it, okay? So if I talk about particles, I want you to tell me about particles. I want you to tell me about this, right? I don't want you to tell me the volume is regular or the volume, if I were talking about liquids, the volume takes the shape of the container, you know? You don't want to talk about that. That would be this stuff. If you want to tell me about this. So covalent solids, how is that different than an ionic solid? Well, it's made up of covalent bonds, okay? Just like diamonds. A diamond is made up of a number of covalent bonds. You can see all of those carbon atoms. It's pretty interesting, I guess, if you're into carbon or whatever. But you can see, like me, I don't know. You can see all of those atoms have tetrahedral structure around each of the atoms, okay? And it's because carbon, you know, would prefer to have a tetrahedral structure if you remember CH4, right, as a tetrahedron. So these are held together by entirely covalent bonds. So it's a big, you know, conglomerate of covalent bonds. So it's like one big molecule. It's like this. High melt, high melting point and boiling point. As you can imagine, it's probably not very easy to melt or get a diamond into a gaseous phase, right? And very, very hard. Infinity, whatever. However many carbon atoms are in that particular sample. So there's a few different what we call allotropes of carbon. Diamond is one of them. If you've heard of buffments or fullerines or buckyballs, that's another one. So it kind of looks like these geogenic domes. Have you ever seen these kind of domes? Like the biodome? Anybody ever seen that polyshore biodome or whatever that thing? That's what, that's like half of buckyballs. Or graphite. Do you use graphite? Yeah, that's carbon too. So those are the three allotropes of carbon. So diamond looks like this where it's all tetrahedral, right? Graphite looks like this. I'm not going to draw a buckyball because have you ever heard of that place, the Louvre? Anybody know that place? Outside of the Louvre, right? There's these kind of structures that look kind of like buckyballs. So you go look at pictures of them. But the structure of graphite, kind of flat sheets of carbon that are like this. Every one of these is carbon. In fact, when you start talking about carbon a lot, carbon is so prevalent in different structures, especially like organic and biological systems, you know that you start condensing the structures and not representing the carbon atoms at all. So this is all by short end. But you can see these, this is actually a plane. All of these atoms are in the same plane. But if you compare that to the diamond up there, they're not all in the same plane. It's a trigonal plane if you want to think about it. Yeah. So you can see, right? So if we cut it up, we'll just cut it up. And all of them, if you think of this as being the central carbon there, right? These are all in the same plane. The bond angle there must be 120 degrees. And there's one, two, three of them, right? So it's a trigonal plane. Okay, so molecular solids, these are just solids that are cooled down below their freezing point, like ice, or in this case, frozen methane. So they're just kind of, instead of that whole conglomerate structure, there are a bunch of little things stuck to each other. And then of course there's metallic solids, like this lightning rod here, and then stroke. So these are metal atoms held together with what we call metal bonds, which is just metal atoms kind of squished together. So if you've ever seen like a copper pipe or something like that, there's no real bonds between those copper atoms. They're just pushed together, okay? Pushed together in what we call a sea of electrons, okay? So they all share their electrons with each other. That's why when, you know, you take a copper pipe and touch an electric fence, right? You feel it like flow through you or whatever. Okay, so you have those overlap of orbitals of the metal atoms. So they kind of just kind of stick together and they share their electrons together. Overlap causes regions of high electron density where the electrons are extremely mobile and they'll conduct electricity. So that's that. That bad, huh? We only have a few more minutes, and what I want to do, let's just pass out the quickness. I was going to, let's see if I can cheat. Is that going to cheat? It'll be right there.