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Chemistry: Intermolecular Forces, States of Matter


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  • Type: Video Tutorial
  • Length: 12:02
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 129 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
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 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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Let me ask you a question. I have a cup of water here, just ordinary water. Why don't all the water molecules just fly out of this cup and disappear? That may seem like a goofy question, but think about it for a moment. Nature strives for ever-increasing disorder. It would be a heck of a lot more disordered if all of this water just disappeared, just took off into the atmosphere. What is it that holds those water molecules together? Clearly there are some forces at work, holding these molecules tightly together. That, in fact, is the reason why we have something called a condensed phase, liquid in other words. We have attractive forces between these molecules.
What we're going to explore now is, what is the origin of those forces? And how does molecular structure ultimately dictate how strong those forces are? To begin with, lets review a little bit about what it is when we talk about a solid or a liquid or a gas. We have an understanding of a solid as being an ordered substance, typically molecules held tightly together. A liquid is a very similar notion, but we have a sense that there's motion now in the molecules. They can move over each other. If they can't completely get free, they can at least travel atop each other. There are still a lot of forces between the molecules, but not quite so many, in general, as the forces in a solid between molecules. Then we have a gas. Of course, we understand the gas as the molecules being completely liberated, freed from each other and in which we've had to break all of those bonds between the molecules.
So what is it that does that? Of course, energy is what's required. We have to put in energy to melt a solid to get a liquid. We have to put in still more energy to break the bonds in a liquid to give us gas.
How exactly does that work? Well, at a given temperature, we have a notion that molecules have some amount of energy in them. When we raise the temperature, suppose that we raise the temperature of water from 20 degrees to 40 degrees, by raising that temperature, somehow we're putting more energy into the water. This shows up in the water as random thermal energy. By random thermal energy, I mean the random motions of individual water molecules, not only how they move from one direction to another, or one location to another, but how they vibrate, how they rotate, how they wiggle, how they bounce off of each other. Every conceivable motion that is going on in those molecules, that is all associated with random kinetic energy. It is a direct function of what the temperature is. So, at a low temperature I have a distribution of different energies. Remember, at one temperature I have a whole broad range of different energies if I were to sample different molecules. They'd have a lot of different energies. A small fraction of them would have sufficient energy to be able to break the bonds holding the molecule to the rest of the molecules, liberating it into the gas phase. It's only a small percentage of the molecules in water at 20, let's say. By raising the temperature to 40 we get a broader distribution. The average energy of the molecules, as far as its thermal energy, increases. Most importantly, a greater percentage of those molecules now have sufficient energy, escape energy if you will, to break out of the surface and again go into the gas phase, turning to water vapor, in other words. By adjusting temperature, we get a notion that that changes the rate at which molecules can go from the liquid into the gas.
We certainly know that that is one of the important components in causing a phase change. The other one is, how much energy do we need? Do we need a lot of energy to break the bonds or do we need just a little energy to break the bonds? On this curve, instead of here, suppose that we have very weak attractions. That is going to move us way over here. Now a lot more of the molecules are going to have enough energy to break into the gas phase.
What is it that holds molecules together, especially when we talk on a molecular level? First of all, let's just make sure that we're all on the same page here and make a distinction between intermolecular forces and intramolecular forces. Intramolecular forces are the type of forces we've been talking about with chemical bonding. What is it that holds atoms together, the sharing of electrons. This is very high energy to break those bonds, much higher energy than the kind of energy needed to break bonds between molecules. This typically is about 1/10^th of the energy required per mole compared to the bond dissociation energy, the energy required to break a carbon/oxygen bond or carbon/hydrogen bond. We're going to focus again on intermolecular forces, interactions between molecules rather than within molecules. To do that, we need to understand what a dipole moment is. A dipole moment is defined as, if you had a charge and you took a second charge of the same magnitude but the opposite sign, and you separated those charges by a specific distance, you would create a dipole moment, a separation of charge. The magnitude of the dipole moment, how strong the dipole moment was, would be a product of whatever the charge was that you were separating, how much it is, "q" in other words, the distance between the two charges. The greater the distance between the two charges, the higher the dipole moment. The bigger the charges that are separated, the higher the dipole moment. Remember, those are the two components that go into a dipole. We can write that mathematically. The Greek letter mu () is the dipole moment, a measure of how big the dipole moment is and it's a product of the charge that you are separating times the distance. We don't have any molecules that look like this, two point charges separated in space by nothingness. But we have a lot of molecules that look like this.
When you are describing a dipole moment for, let's say, the molecule hydrogen fluoride, what we're really saying is, imagine in your mind a model for hydrogen fluoride that has a separation in charge, what would that charge need to be and what would the distance need to be to have a dipole moment that was equivalent to the actual dipole moment in this molecule? What's really important for you to understand is that when we talk about a dipole moment of a molecule, we're simply asking how much charge is separated and by how much distance is that charge separated, again q x d = (dipole moment). We know that this has to do with how electronegative these elements are. The greater the difference in electronegativity, the more polarized that bond becomes. The greater separation of charge the higher the dipole moment. We can predict which way it's pointing. It's going to be pointing, in the case of hydrogen fluoride, towards the fluorine because it is more electronegative. It gets more electron density, higher dipole moment pointing that direction.
What about a water molecule? We know that in a water molecule, we have a very polarized bond between hydrogen and oxygen. We know that we could justify focus on that bond for a moment, a dipole pointing in this direction and a dipole pointing in this direction. Think about the sum of those two dipole moments. I'm going to put this on the piece of paper here for a moment. I'm going to push in from both of those sides in the direction of the dipole and look at the motion of the molecule. As I am pushing in, I am moving the molecule in this direction. I'm just simply saying that the sum of these two vectors is a vector pointing in this direction. I have a dipole moment overall, that points towards the oxygen and right up through the middle of the molecule. I have a very polarized molecule in water. We're going to see that that is very important.
When we go to ammonia, this is another example where we have a polarized bond, a set of three polarized bonds. The bond is polarized towards the more electronegative nitrogen. Ammonia is not flat. If it were completely flat there would be no net dipole because the individual bond dipoles would all cancel. Because ammonia is pyramidal, meaning it's pushed up--let's remind ourselves that a pair of electrons up here pushes these guys down, the dipole moments actually point in this kind of diagonal direction. They cancel each other out in this way but they don't cancel each other out up and down. There's a net dipole moment going right up the middle of ammonia pointing in the direction, again, through nitrogen, away from the hydrogens. Back on our picture here, we have a dipole moment again for ammonia.
Now, let's think about the molecule methane. In methane, we also have a polarized bond, not very polarized any more. There is just a small difference in the electronegativity of carbon and hydrogen. However, we do have a polarized bond. However, we also have a polarized bond here, here and here. If you think about the molecule methane, with tetrahedral geometry, it's a beautifully symmetrical molecule. These four dipoles completely cancel each other out. There's no net dipole in this molecule whatsoever. So this molecule would be considered non-polar. It has no net dipole moment at all. The important reason for talking about a dipole moment is that the stronger the dipole moment in the molecule, as we'll talk about more, the greater the attraction and the harder it is going to be to break those molecules apart to make gas molecules.
Finally, let's talk about how, in the same way we can take a solid, add energy to it to make a liquid, add more energy to it to make a gas, we can take that gas and condense it to form a liquid. When we do that, we lose the amount of energy, the system liberates this amount of energy, which is the same amount of energy we would have had to put in to vaporize it. That has to be lost again. It's a reversible process. That liquid, furthermore, could be converted into a solid by removing still more energy. How much energy? The same exact amount of energy that we had to put in to melt it. All of the phases and the changes between the phases--that whole series of processes are fully reversible. Put in energy. Put in energy. Take it back out. Take it back out.
We're now going to explore this relationship between how much energy we put in and what the attraction is between molecules.
Condensed Phases: Liquids and Solids
Intermolecular Forces
Introduction to Intermolecular Forces and States of Matter Page [1 of 2]

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