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Chemistry: Intermolecular Forces

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  • Type: Video Tutorial
  • Length: 17:06
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 182 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Final Exam Test Prep and Review (49 lessons, $64.35)
Chemistry: Condensed Phases: Liquids and Solids (15 lessons, $25.74)
Chemistry: Intermolecular Forces (2 lessons, $4.95)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at http://www.thinkwell.com/student/product/chemistry. The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

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Thinkwell
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If we can understand in detail the interactions between molecules, what holds molecules together, then we can start to be able to predict their physical properties, such as boiling points or melting points. So to do this, let's go ahead and, first of all, broadly go over the types of attractions between molecules, in other words, intermolecular attractions again. The strongest attraction that we would have would be if both of our molecules had a charge, but the opposite charge, in other words. We refer to this as an ion-ion or a charge-charge interaction and I show that in the little cartoon here. Of course, they would have to be, again, opposite charges.
The second strongest interaction between two different molecules would be one involving an ion or charge in one of them and a dipole moment in the other. Okay, so again, what we would refer to as an ion-dipole interaction would be aligning the dipole moments such that its slightly negative end, for instance, was pointing toward the positive end of an ion. Now, to demonstrate these two ideas, ion-ion, and ion-dipole, I've asked my colleague, Professor Gordon Yee to join me. Gordon, if you would please... What I'm first going to demonstrate for you is a charge-charge interaction. I'm going to rub this balloon on Gordon's head and in doing so, I am creating a surplus of negative charge on the balloon and a buildup of positive charge on his head. So when I let go, the two particles stick together now, again, an ion-ion interaction in this case.
Now what I'm going to do, again, is charge up the balloon, but this time, instead of putting it on Gordon's head, I'm going to put it on the wall behind me, which has no permanent charge in it. Nonetheless, we're going to be able to induce a dipole moment in the wall, thereby creating an ion-dipole interaction.
Now, the next weakest interaction is going to be a dipole-dipole interaction and to illustrate this point I'm going to us as an analogy two bar magnets. So again, just to clarify, this is only an analogy. These are magnetic dipoles rather than electrical dipoles, but the idea is the same in that if I have two opposite poles approach each other, there's an attraction between those two poles. Positive and negative charges, again, attract. Neither one of these have a net charge. Neither one of these two molecules have a net charge, but nonetheless their separation in charge is enough to induce an attraction between them. Correspondingly, if I bring up the same charges, if I brought up the two negative ends or the two positive ends, I would have repulsion.
Now the weakest form of attraction, and we'll talk about this in a moment, is what's referred to London/dispersion forces. These are interactions between molecules that don't have any permanent dipole moment or charge. Now actually, London forces exist in all molecules but they are responsible uniquely for holding molecules together that are not polarized at all. Again, we'll talk about this. Collectively, all of these types of interactions are referred to as van der Waals forces.
Okay, so let's talk in detail now about dipole-dipole interactions. I've listed several examples of molecules that have a dipole moment associated with it. These are only a few of many, many, many different molecules that are polarized. Ammonia is a gas at room temperature, but not too much below room temperature, at about 30 below, that's 30 Celsius below, so -30, ammonia turns into a liquid. Compare that to methane, which is about the same size but doesn't reach a liquid until many, many, over a 100 below that. Acetone is a very common solvent, again a small molecule, but unlike many other small molecules that are gasses at room temperature, acetone is liquid at room temperature. It is the solvent that used to be used for nail polish remover as well as many other industrial uses for acetone. Acetone, again, has a strong dipole moment associated with it, pointing in the direction of the oxygen. Another common solvent is methylene chloride used a lot in industry for degreasing things, also used to be used for decaffeinating coffee beans. It would be the solvent that would dissolve caffeine, which, again, is polar molecule, and the methylene chloride will interact with that and remove that caffeine from the coffee bean. It's no longer used so much; water is now kind of a preferred solvent for that process. In all of these situations, the molecules in the liquid phase line up head to tail, such that the positive end, let's say, of the dipole is next to the negative end of the next dipole and so we have this type of ordering in the liquid phase. And once again, the important point here is, the stronger that interaction is the greater these dipole moments, in general, the stronger we predict the attractive forces would be and therefore the higher we would predict the boiling point, although, we'll get into exactly how that works soon, but again, the stronger the forces are holding the molecules together. And again, we'll see that that relates to physical properties.
Now the weakest of all forces, again I mentioned, are these London forces, London/dispersion forces. These are forces in molecules that have no net dipole moment, that have no net charge, yet they still offer, I mean yet there's still an attraction. So, for instance, if we consider just the atoms helium versus argon versus xenon, all of those, especially the latter two form liquids at low temperatures. Now helium also will form a kind of a liquid but we have to get very, very close to absolute zero before we actually see that, whereas these other materials at much higher temperatures, form liquids. Well, what is it that holds those atoms together? London forces. And what is a London force? Well, a London force has to do with the fact that the electron cloud in larger atoms especially, is not held very tightly by the nucleus. Remember, as we get further and further away from the nucleus, there's not as much attraction on those electrons. Being looser held they are more polarizable. They're more squishy. They're more flubbery, if you will. You can kind of push them around a little bit like Jell-O. And at any instant in time, you can induce a dipole moment, or rather, at any instant in time the molecule may not be perfectly symmetrical and have a small dipole moment associated with it that's not a permanent dipole. This is a very important difference. There's no net dipole moment, remember, but it has a small, instantaneous dipole moment and that, in turn, induces a dipole moment in the next molecule or atom, which induces a dipole moment in the next one, and so on. So that we have these attractions due to the small dipole moment in one which, again, induces a dipole moment in the next, and so we get this attractive force. If you look over here, I show you a picture of chlorine, of bromine, and of iodine, all at room temperature. You'll notice a pretty significant difference here. Chlorine is a gas at room temperature. Bromine is a liquid at room temperature, and iodine, in fact, is a solid at room temperature. None of these molecules have dipoles. None of these molecules are charged. The only difference is that there are much stronger London forces in the iodine than there are in the bromine or the chlorine, once again, because the electrons in the iodine are more squishy. They're more polarizable. Being further from the nucleus, they are not restricted as much.
Now finally, let's talk about the last type of interaction and I have a riddle for you. What would be the one atom on the periodic table that readily forms bonds with other atoms, yet has no core electrons at all? The answer, of course, is hydrogen. Hydrogen uniquely bonds with other things but has no core electrons. In other words, its valence electrons or electron is the only electron it's got. Now, the important significance about this is that when it's engaged in bonding, its nucleus is partially exposed, the nucleus itself. So other molecules that have partially negative charges can get very, very close to that positive charge and that results in a strong interaction. So let's look at what's going on in hydrogen bond and I'll use the example here of a hydrogen bond in water and I've illustrated here for you, again, to clarify the difference between this kind of an OH, the OH within a water molecule, and this type of an OH interaction, the interaction between a molecule. So once again, the two words intermolecular versus intramolecular interactions. This is much, much stronger than this one, but this bond is polarized, very strongly polarized and we have a partially exposed nucleus here, no electrons blocking it. So that is going to be very attracted to buildup of negative charge such as a lone pair on this oxygen.
We have two requirements for a hydrogen bond: first, that the hydrogen actually be bound to a very electronegative atom. Usually this is something like oxygen, chlorine or fluorine. It's got to be a very electron-deficient molecule that's going to pull in electron density. But the second important requirement is that the hydrogen needs to, or is this piece of it, what's referred to as the hydrogen bond donor or, excuse me, let me back up, the hydrogen bond acceptor, that is the lone pair on the oxygen. That's available to interact again with this dipole moment.
So again, the term the hydrogen bond acceptor is actually this lone pair that can interact with the hydrogen and this would be the hydrogen bond donor, if you would. It's the one that's actually giving the hydrogen for that interaction. So, again, there's our strong interaction. If you imagine what is going on in ammonia, ammonia has a lone pair, certainly, and it's got a hydrogen, but it doesn't form good hydrogen bonds because the bond between nitrogen and hydrogen is not very polarized, and as a result of not being so strongly polarized, the hydrogen is not as exposed and so it doesn't have as strong of an interaction with that lone pair coming from nitrogen. On the other hand, ammonia would be a great hydrogen bond acceptor. If I replace this water molecule with ammonia, we'd get a good hydrogen bond but ammonia doesn't form good hydrogen bonds with itself. Other kinds of things that participate in hydrogen bonding would be alcohols like methanol or ethanol, carboxylic acids--I show you an example here of acetic acid. Let's look at a model of this really quickly. Acetic acid, again, has got this polarized OH bond, okay. This is, in fact, unusually acidic for an OH bond and you might be able to figure out why if you think about the resonance structure you could draw for acetic acid. So we have a very exposed positive charge here and that, again, can interact with lone pairs from other molecules of acetic acid, let's say. A very similar idea is this molecule.
This is an example of an amide. In this case, that is an NH bond, but it's much more acidic. In other words, that hydrogen is much more exposed than the hydrogen would be on ammonia, for example, and, again, if you consider a different resonance structure than the one I've drawn here for this molecule, you can start to understand why that hydrogen is more acidic than normal. I'll leave you with that puzzler.
But, again, these are several common types of molecules that participate in hydrogen bonding, amides, acids, and alcohols; that we want to be aware of. And these guys all, typically, have very high boiling and melting points associated with them. Why? Because what causes something to boil or melt has to do with how easily you break the bonds and these being the strongest type of intermolecular interaction, that's going to give us then, that's going to require, rather, the highest amount of energy to break those bonds and so we'll see we have to go to higher temperatures for that. So okay, we've kind of set the stage now. We're ready to make the next leap to going from what's happening on a molecular level to what are the consequences of these types of interactions.
Condensed Phases: Liquids and Solids
Intermolecular Forces
Intermolecular Forces Page [2 of 3]

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