[MUSIC] In this video I will introduce one of my favorite concepts, intermolecular forces. Recall from a previous video that there are two factors that determine the state of a given sample of material: the thermal energy, which was related to the temperature, and the interaction energy, which is related to the intermolecular forces. What are these intermolecular forces? Well there's different types of intermolecular forces. I'll go through these one at a time and give examples. The first type is called dispersion forces. Sometimes this is also called induced polar interactions. Everything has dispersion forces. All matter experiences some interactions with other types of matter when they get close enough together. The next type of interaction energy is dipole-dipole forces. This is a special case where this type of interaction is only there for polar molecules. And I'll talk about how to determine polarity in this lecture. The third type of interaction energy is hydrogen bonding. This is a very powerful but very specialized type of interaction that is only coming into play for molecules that have an O-H bond, an N-H bond, or an F-H bond. So there are many types of matter that don't have hydrogen bonding. But water has hydrogen bonding because it has an O-H bond. And that's part of the reason we spend so much time thinking about it. Marching through these one at a time, the first one on the list was dispersion forces, which, remember, can also be called London dispersion forces or Van Der Waals forces. These are present in all molecules and atoms, and they increase with increasing number of electrons. So this parallels the trend in atomic number. The larger a molecule is, the more dispersion forces the molecule feels when it's near other molecules of the same type. This is from temporary induced dipole states. So let's look at a simple molecule, a simple diatomic molecule. Could be H2 or N2 or I2. Take your pick. That molecule has an oval shaped cloud with two nuclei that are here and here and if three of those molecules are next to each other we can look at the electron cloud and see that the average electron cloud is about an oval shape. But there are moments in time where there are more electrons on one side of the molecule than on the other, where the cloud is not symmetrically oval. And at that moment in time, there is a partial positive charge on the side of the molecule that doesn't have a normal allocation of electrons, and there is a partial negative charge on the side of the molecule that has more than it should, in terms of electron cloud at that moment of time. But remember there's coulombic attraction and repulsion. So the electrons on this neighboring molecule would notice that there's more electron density near them than normal, and they would move away from that electron density. So, what happens is that at a moment in time, a snapshot in time, all the electron clouds have these temporary induces dipoles. This is a temporary state because the electrons of course are moving very rapidly. But at that moment in time when I have the pluses and minuses lined up to give a maximized coulombic attraction, right, they have the minus lined up with the plus on each one of these, that gives rise to these dispersion forces. The more electrons there are, the greater the temporary induced dipole can be, right? Because if you have 12 electrons on one side of the molecule and none on the other, that's going to be a bigger dipole than if you just have one electron which is moved to the, a single side of that molecule. So everything has intermolecular forces. A lot of books describe them as being fairly weak. But I think that's inaccurate because intermolecular forces are strong enough for example, to make iodine I2 molecule a solid at room temperature and pressure. So there must be a lot of intermolecular interaction in I2, in order to make it a solid, with the amount of thermal energy that we have at room temperature. The second type of intermolecular forces is dipole-dipole forces. This only occurs when there is a permanent dipole in the molecule. In order to determine this, we need to determine if the bonds are polar and what is the molecular shape. So single atom species, such as argon gas do not have any dipole-dipole forces, because there's no bond present. But if we have any bonds, we can look at the electronegativity difference between the atoms, which remember we can say is delta chi, and we can subtract the electronegativity of the smaller electronegativity atom from the larger one. So, for example, let's look at the molecule HF. The electronegativity difference there would be 4.0 for fluorine minus 2.1 for hydrogen. So, the electronegativity difference there is 1.9. And that's a large enough electronegativity difference that we say that HF is a polar bond. Here are some examples of bond polarities for other types of bonds. If both atoms in the bond are the same type of element, for example, over here, with the chlorine-chlorine bond, then that's a non polar bond. If the electronegativity difference is greater than .4, say up to about 2.0, we say that that is a polar covalent bond. And remember there's a lot of grey area here. Some people draw the line for where you transition from covalent to ionic in a different place, but really it's a continuum, so drawing a line is a bit artificial. So for my non polar molecule the electronegativity difference with these are zero or very close to zero. So that's what I'm calculating up here is delta chi. We already did the H-F bond and we saw that that was 1.9. So that gives us a partial positive and a partial negative pull, you probably can't see that, it's too little, on that molecule. If the electronegativity difference is greater than 2.0, generally we say that's an ionic bond, and an example that you're most familiar with is sodium chloride, which has an electronegativity difference of 2.1. There are ionic compounds that have even larger electronegativity differences. Here are some of those on this slide. So for example, the rubidium fluoride bond is even more polar than the sodium chloride bond. And so when we get above a certain percent ionic character, here we say we draw the line at 50% ionic character and we plot that versus electronegativity difference, right, we see that the 50% ionic character happens somewhere between say 1.7 and 2.0. And things that are around that electronegativity difference are very polar, and things that are down near zero like phosphorus, hydrogen, and carbon hydrogen bonds, or a bond between two atoms of the same type of element, like a carbon-carbon bond, we say those things are non polar. I think for our purposes, the three most important non polar bonds are carbon-hydrogen, phosphorous-hydrogen and carbon-carbon, or any time we have two atoms of the same type of element. Right, that's also really important. So a fluorine-fluorine bond is also non polar, because the electronegativity difference would be zero. I've shown you these little arrows that illustrate dipole moments before. We put the positive end of the arrow on the less electronegative element, the on that's slightly electron deficient and we put the pointy part of the arrow on the more electronegative element, the part that is partially positive. Here's an O-H bond drawn out really big. And if we look at the electron cloud, around that water, we see that there's more electron density in the cloud around the oxygen, and there's more electron density around an OH bond's oxygen than there would be around an oxygen that was bonded to another oxygen. So oxygen here has a partial negative charge, hydrogen in of, of an OH bond has a partial positive charge. These are not full charges. So, they're not ions, like we saw with ionic compounds. These charges come from asymmetric electron clouds. So, instead of the cloud being an oval shape for this sigma bond, it's slightly skewed to one side. Dipole moment, which is given the symbol mu, is a measure of the magnitude of this polarity, and it's measured in the unit called Debye. So, dipole-dipole forces, we have to ask ourselves two questions and so far I've only introduced or examined the first question. Are the bonds polar? We determine that again by looking at electronegativities of the different atoms making the bond. But the second question that is really important is: what is the molecular shape, and how does that shape influence the arrangement in space of any bond dipoles? We learned about molecular shapes in another week's lectures. So let's apply that information now. Let's determine if there is a permanent dipole moment in the molecule carbon dioxide. So we're looking for dipole-dipole forces, which is a type of intermolecular forces. Our example is CO2, and the questions are, are the bonds polar? Well yes! A carbon oxygen bond is polar. It gives me little bond dipoles that look like that. The second question is what is the molecular shape? Well, we did this example several times, so hopefully you know that it's linear, but even if you've forgotten, you can draw the Lewis dot structure, determine the arrangement of the electron regions around the simple central atom, and come up with linear as your answer. In this case, the bond dipoles are equal in magnitude, because they're both for carbon-oxygen bonds, and they are pointed in opposite directions. A bond dipole moment is a vector quantity. So if you added these two vectors, they would add up to 0. In other words, the bond dipoles cancel each other out. So, even though carbon dioxide has polar bonds, the overall molecule is non polar, and carbon dioxide does not have dipole-dipole forces. Let's do another example. How about water? Is water polar? We can draw the Lewis dot structure, and when we do that, it reminds us that oxygen has two lone pairs on it, which forced the hydrogens down toward each other into a bent shape. Each oxygen-hydrogen bond is polar. And in this case, the bond dipoles port, point towards the central oxygen atom. If we added up these little bond dipoles that are shown here in blue, they add up to a net dipole that is pointed straight up if the water is in this orientation. So, water has polar bonds that's the first criteria that must be met to have a polar molecule, and it has a bent shape. So it does experience dipole-dipole intermolecular forces. And those energies of interaction combine with the other intermolecular forces that are present in water to make it the liquid state at room temperature and pressure, even though it's a fairly small molecule. Let's review polar and non polar. Water was polar. It has a net dipole moment. It has polar bonds and a bent shape. Carbon dioxide was non polar. It has no net dipole moment. It's linear, and so the bond dipoles cancel out. There can be other shapes where bond dipoles can cancel out. All that has to happen is if you add up the vectors they would need to add up to zero. Another way to think of this is, is there equal pull on the electrons in all, in enough directions for them to cancel out, and not have any net pull. When I say pull I'm saying P-U-L-L. Let's do some practice to review this concept. Predict whether each of the following molecules is polar or non polar. The first molecule I'm giving you is trifluoromethane. The structure is shown here. Is trifluoromethane polar or non polar? Thank you for submitting your answer. The trifluoromethane has bond dipoles between the carbons and the hydrogens in these directions. And it has a very, very tiny bond dipole between the carbon and the hydrogen. But it's very tiny. But hopefully you can see these bond dipoles do not cancel out and therefore, this particular molecule was polar. What about difluormethane? So now we're doing the one on the right. Is difluormethane polar or non polar? Thank you for submitting your answer. This one was quite a bit more difficult to answer. It looks like, in the flat drawing of the Lewis dot structure, that these two carbon-fluorine bond dipoles cancel out. But you have to remember that the shape is not really as shown with 90 degree bond angles. The shape is really tetrahedral isn't it? If I drew it tetrahedral hopefully you can see that the bond dipoles between the carbons and the fluorines don't, are not 180 degrees from each other, are they? I can redraw this on the paper with the two fluorines in the plane of the paper. Could draw it any way I want. Now you can really see it. Here's a one bond dipole. Here's one bond dipole, here's another one. So there's a net bond dipole, I'm sorry, there's a net molecular dipole in this direction that intersects the two fluorines. The last intermolecular force that we need to discuss is a very special case, hydrogen bonding. This is only true for molecules that have an O-H, N-H, or F-H bond. In order for a molecule to experience hydrogen bonding, it must have hydrogen, and it must have either oxygen, nitrogen, or fluorine. And, and this is really important, the oxygen, nitrogen, or fluorine must be directly bonded to a hydrogen. In hydrogen bonding, the lone pairs of one atom are making a partial bond to the hydrogen of a neighboring molecule. And in fact, the hydrogens in a sample of water are being exchanged. They're being passed around from oxygen to oxygen. So this is a stronger interaction than static cling, for example. But it's not quite as strong as a covalent bond. This does give water a very high surface tension. You can observe that if you overfill your glass of water. And it also gives water a relatively high boiling point when we consider the small formula mass of water. Let's review the intermolecular forces. The first one was dispersion forces. Everything has these. Some books will say that they're relatively weak. But I question that, because some of, some molecules that are large have very large dispersion forces. This was the induced dipole forces. The second type is dipole-dipole forces, then that is specific for polar molecules. And the third type of hydrogen bonding, which is a very specific case only for molecules with certain types of bonds. All of these forces are additive. So you would, you could look at how much dispersion force contributes to the interaction energy. If it's a polar molecule, you can add to that how much the polarity contributes to the interaction energy etcetera. Just to give you an idea of the scale of these numbers ,showing the strength of these forces, how many kilojoules per mole of force is there, I'll give you the reference that covalent bonds are between one hundred and one thousand kilojoules per mole. Many of the covalent bonds that we talked about in this class are in the range of a few hundred kilojoules per mole. This concludes the lecture on intermolecular forces. Remember the types and extent of intermolecular forces determine the interaction energy. And that factor combines with the thermal energy at a given set of conditions in a particular system, to determine the state of the material.
