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Intermolecular forces

Different types of intermolecular forces (forces between molecules). Created by Jay.

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  • marcimus pink style avatar for user Susan Moran
    Hi Sal,
    I am a 60 year old woman who has been taking Geometry, Chemistry and Algebra in order to help my 15 year old niece. I wondered about something, and there seemed to be no definitive answer when I Googled it. The question is: do thoughts have mass? I think that they must have an infinitesimal, at this time, immeasurable mass.

    Thanks,

    Sue
    P.S. I think that I am addicted to Khan Academy!
    (28 votes)
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    • mr pink red style avatar for user Jeffrey Baum
      thoughts do not have mass. the reason is because a thought merely triggers a response of ionic movement (i.e. Na+, K+ ) these ions already exist in the neuron, so the correct thing to say is that a neuron has mass, the thought is the "coding" or "frequency" of these ionic movements. so a thought does not have mass. i like the question though :)
      (32 votes)
  • blobby green style avatar for user smasch2109
    If you have a large hydrocarbon molecule, would it be possible to have all three intermolecular forces acting between the molecules? Or is it just hydrogen bonding because it is the strongest? How do you determine what forces act when you have big and diverse molecule like an anhydride, e.g. acetic anhydride: Would here be dipole-dipole interactions between the O's and C's as well as hydrogen bonding between the H's and O's? Or just one of the two? What about the london dispersion forces?
    (13 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      You can have all kinds of intermolecular forces acting simultaneously. Usually you consider only the strongest force, because it swamps all the others.
      When you are looking at a large molecule like acetic anhydride, you look at your list of intermolecular forces, arranged in order of decreasing strength. Then you go down the list and stop at the first one that fits your molecule.
      We can omit all the forces involving ions. That leaves
      H-bonding
      Dipole-dipole
      Dipole - induced dipole
      London Dispersion
      What is the strongest .intermolecular force in acetic anhydride?
      H-Bonding ? No, because there are no O-H, N-H, or F-H bonds in the molecule.
      Dipole-dipole ? Yes, because the molecule has polar C=O groups.
      Dipole-induced dipole and London Dispersion forces are also present, but they are small in comparison to the dipole-dipole forces.
      (42 votes)
  • male robot hal style avatar for user Roni
    Can someone explain why does water evaporate at room temperature; having its boiling point at 100ºC?
    (17 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      In water at room temperature, the molecules have a certain average kinetic energy.
      But some of them will be moving faster, and some of them will be slower.
      Some of the molecules at the surface of the water will have enough kinetic energy to escape into the atmosphere.
      The water will slowly evaporate, even though it is at room temperature.
      (31 votes)
  • piceratops ultimate style avatar for user Jack Friedrich
    At , he says that the boiling point for methane is around -164 degrees celsius. So, can methane be in a liquid form, if it is colder than -164 degrees celsius? And if so, is it present in a liquid form naturally anywhere?
    (8 votes)
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  • aqualine seedling style avatar for user Venkata Sai Ram
    how can a molecule having a permanent dipole moment induce some temporary dipole moment in a neighbouring molecule
    (3 votes)
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    • orange juice squid orange style avatar for user awemond
      Suppose you're in a big room full of people wandering around. If I bring a smelly skunk into the room from one of the doors, a lot of people are probably going to move to the other side of the room. When the skunk leaves, though, the people will return to their more even spread-out state. Conversely, if I brought a bunch of cupcakes there might be a rush for my side of the room, though people would spread out again once the cupcakes were gone.

      In this terrible metaphor, the people walking around the room are like the electrons evenly spread around a molecule. If you bring something negatively charged near the molecule (like the negative end of a permanent dipole), then the negatively charged electrons will be repelled and will concentrate on the far side of the molecule, making the near side slightly positively charged and the far side slightly negatively charged (ie you've made an induced dipole!). When the permanent dipole goes away, though, it is more stable for the electrons to spread out again. Conversely, if you bring the positive side of a permanent dipole near a molecule, the electrons in the molecule will rush towards the positive side, leaving the far side of the molecule with fewer electrons and thus a temporary positive charge.
      (37 votes)
  • blobby green style avatar for user cpopo9106
    In the notes before this video they said dipole dipole interactions are the strongest form of inter-molecular bonding and in the video he said hydrogen bonding is the strongest. can you please clarify if you can. Thanks
    (5 votes)
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  • spunky sam blue style avatar for user Tobi
    if hydrogen bond is one of the strongest inter molecular force why is ammonia a gas and hydrogen fluoride (most electronegative of the FON elements) has a boiling point of 19.5 degree C and water 100 degree C?..... quite a wide variation in boiling point and state of matter for compounds sharing similar inter-molecular force
    (6 votes)
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    • female robot grace style avatar for user tyersome
      Good question!

      From your question I assume you've worked out why HF has a higher boiling point than NH₃.

      To see why H₂O has a much higher boiling point, try drawing each molecule and all possible hydrogen bonds they can make. What do you see?

      Does this help you answer your question?
      (3 votes)
  • blobby green style avatar for user Sastha Rajamanikandan
    At , he says "double bond." I know where he gets that term from but what is exactly a double bond?
    (3 votes)
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  • primosaur ultimate style avatar for user FrenchHornCat
    At , a textbox shows up on the screen clarifying that the partial positive charge is on the carbonyl carbon rather than on the CH3 carbon. Why is this? Is this because the net ionic charge on the carbonyl carbon is more positive (would I need to actually make a dot structure diagram to figure that out if I theoretically wanted to)?
    (3 votes)
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    • leaf red style avatar for user Richard
      Two reasons why.

      First is the electronegativity difference between carbon and oxygen. Oxygen is more electronegative than carbon so the electrons in their covalent bond would rather spend more of their time around the oxygen as opposed to the carbon. Electrons are negatively charged so the oxygen with more electrons around it at any given time means it'll have a partial negative charge as a result. And carbon which has a lack of electrons around it because of this will experience a partial positive charge because of the lose of electron density around it.

      Second is the contribution of resonance. If you draw a resonance structure of the acetone molecule in you'll find you can push pi electrons of the double bond of the carbonyl into the oxygen so it has three lone pairs around it now instead of just two. If you add up for the formal charges on this second resonance structure you'll find the oxygen has a 1- charge and the carbonyl carbon has a 1+ charge. Now the actual structure of acetone is a hybrid of combination of these two resonance structures, what we call the resonance hybrid. The resonance structure with the formal charges also helps add some partial negative charge to the oxygen and partial positive charge to the central carbon.

      So it's a combination of both these reasons as to why we get those partial charges on acetone.

      Hope that helps.
      (4 votes)
  • duskpin ultimate style avatar for user Isabel
    Why does oxygen need to take 4 electrons from Carbon? It only needs 2 electrons to fulfill it's octet so shouldn't there just be one line that is drawn?
    (3 votes)
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    • leaf red style avatar for user Richard
      So for acetone's Lewis dot structure, if you just kept a single line (or single bond) between the central carbon and the oxygen you would indeed give oxygen an octet since it already has the 6 electrons from the lone pairs. However, this results in two things we have to consider. First with only a single bond to the oxygen, the central carbon now lacks an octet since it will only have 3 single bonds (1 to the oxygen and 2 to the other carbons). This means the central carbon only has 6 electrons from the three single bonds. The second thing we have to consider are the formal charges around the oxygen and the central carbon atom. The central carbon with only three single bonds will have a formal charge of +1. The oxygen with 3 lone pairs and a single bond will have a formal charge of -1. These formal charges of +1 and -1 still add to an overall charge of 0 for the molecule which makes sense since it's neutral. So what this means together is that you can draw the Lewis structures of acetone with a C=O bond or a C-O and both are valid structures. When we can draw multiple Lewis structures for the same molecule we call these resonance structures. Resonance structures individually don't represent how the molecule actually looks, rather the molecule looks like combination of the two resonance structures called a resonance hybrid. However the resonance structure with a C=O bond has less charge separation (no + or - formal charges on any of the atoms) so we say that it is the dominant resonance structure of the two. So conventionally we draw acetone with with a C=O bond more so when drawing its Lewis structure for simplicity sake. So drawing acetone your way isn't wrong, it's just less conventional. Long explanation, but I hope that helps.
      (2 votes)

Video transcript

In the video on electronegativity, we learned how to determine whether a covalent bond is polar or nonpolar. In this video, we're going to see how we figure out whether molecules are polar or nonpolar and also how to apply that polarity to what we call intermolecular forces. Intermolecular forces are the forces that are between molecules. And so that's different from an intramolecular force, which is the force within a molecule. So a force within a molecule would be something like the covalent bond. And an intermolecular force would be the force that are between molecules. And so let's look at the first intermolecular force. It's called a dipole-dipole interaction. And let's analyze why it has that name. If I look at one of these molecules of acetone here and I focus in on the carbon that's double bonded to the oxygen, I know that oxygen is more electronegative than carbon. And so we have four electrons in this double bond between the carbon and the oxygen. So I'll try to highlight them right here. And since oxygen is more electronegative, oxygen is going to pull those electrons closer to it, therefore giving oxygen a partial negative charge. Those electrons in yellow are moving away from this carbon. So the carbon's losing a little bit of electron density, and this carbon is becoming partially positive like that. And so for this molecule, we're going to get a separation of charge, a positive and a negative charge. So we have a polarized double bond situation here. We also have a polarized molecule. And so there's two different poles, a negative and a positive pole here. And so we say that this is a polar molecule. So acetone is a relatively polar molecule. The same thing happens to this acetone molecule down here. So we get a partial negative, and we get a partial positive. So this is a polar molecule as well. It has two poles. So we call this a dipole. So each molecule has a dipole moment. And because each molecule is polar and has a separation of positive and negative charge, in organic chemistry we know that opposite charges attract, right? So this negatively charged oxygen is going to be attracted to this positively charged carbon. And so there's going to be an electrostatic attraction between those two molecules. And that's what's going to hold these two molecules together. And you would therefore need energy if you were to try to pull them apart. And so the boiling point of acetone turns out to be approximately 56 degrees Celsius. And since room temperature is between 20 and 25, at room temperature we have not reached the boiling point of acetone. And therefore, acetone is still a liquid. So at room temperature and pressure, acetone is a liquid. And it has to do with the intermolecular force of dipole-dipole interactions holding those molecules together. And the intermolecular force, in turn, depends on the electronegativity. Let's look at another intermolecular force, and this one's called hydrogen bonding. So here we have two water molecules. And once again, if I think about these electrons here, which are between the oxygen and the hydrogen, I know oxygen's more electronegative than hydrogen. So oxygen's going to pull those electrons closer to it, giving the oxygen a partial negative charge like that. The hydrogen is losing a little bit of electron density, therefore becoming partially positive. The same situation exists in the water molecule down here. So we have a partial negative, and we have a partial positive. And so like the last example, we can see there's going to be some sort of electrostatic attraction between those opposite charges, between the negatively partially charged oxygen, and the partially positive hydrogen like that. And so this is a polar molecule. Of course, water is a polar molecule. And so you would think that this would be an example of dipole-dipole interaction. And it is, except in this case it's an even stronger version of dipole-dipole interaction that we call hydrogen bonding. So at one time it was thought that it was possible for hydrogen to form an extra bond. And that's where the term originally comes from. But of course, it's not an actual intramolecular force. We're talking about an intermolecular force. But it is the strongest intermolecular force. The way to recognize when hydrogen bonding is present as opposed to just dipole-dipole is to see what the hydrogen is bonded to. And so in this case, we have a very electronegative atom, hydrogen, bonded-- oxygen, I should say-- bonded to hydrogen. And then that hydrogen is interacting with another electronegative atom like that. So we have a partial negative, and we have a partial positive, and then we have another partial negative over here. And this is the situation that you need to have when you have hydrogen bonding. Here's your hydrogen showing intermolecular force here. And what some students forget is that this hydrogen actually has to be bonded to another electronegative atom in order for there to be a big enough difference in electronegativity for there to be a little bit extra attraction. And so the three electronegative elements that you should remember for hydrogen bonding are fluorine, oxygen, and nitrogen. And so the mnemonics that students use is FON. So if you remember FON as the electronegative atoms that can participate in hydrogen bonding, you should be able to remember this intermolecular force. The boiling point of water is, of course, about 100 degrees Celsius, so higher than what we saw for acetone. And this just is due to the fact that hydrogen bonding is a stronger version of dipole-dipole interaction, and therefore, it takes more energy or more heat to pull these water molecules apart in order to turn them into a gas. And so, of course, water is a liquid at room temperature. All right. Let's look at another intermolecular force. And this one is called London dispersion forces. So these are the weakest intermolecular forces, and they have to do with the electrons that are always moving around in orbitals. And even though the methane molecule here, if we look at it, we have a carbon surrounded by four hydrogens for methane. And it's hard to tell in how I've drawn the structure here, but if you go back and you look at the video for the tetrahedral bond angle proof, you can see that in three dimensions, these hydrogens are coming off of the carbon, and they're equivalent in all directions. And there's a very small difference in electronegativity between the carbon and the hydrogen. And that small difference is canceled out in three dimensions. So the methane molecule becomes nonpolar as a result of that. So this one's nonpolar, and, of course, this one's nonpolar. And so there's no dipole-dipole interaction. There's no hydrogen bonding. The only intermolecular force that's holding two methane molecules together would be London dispersion forces. And so once again, you could think about the electrons that are in these bonds moving in those orbitals. And let's say for the molecule on the left, if for a brief transient moment in time you get a little bit of negative charge on this side of the molecule, so it might turn out to be those electrons have a net negative charge on this side. And then for this molecule, the electrons could be moving the opposite direction, giving this a partial positive. And so there could be a very, very small bit of attraction between these two methane molecules. It's very weak, which is why London dispersion forces are the weakest intermolecular forces. But it is there. And that's the only thing that's holding together these methane molecules. And since it's weak, we would expect the boiling point for methane to be extremely low. And, of course, it is. So the boiling point for methane is somewhere around negative 164 degrees Celsius. And so since room temperature is somewhere around 20 to 25, obviously methane has already boiled, if you will, and turned into a gas. So methane is obviously a gas at room temperature and pressure. Now, if you increase the number of carbons, you're going to increase the number of attractive forces that are possible. And if you do that, you can actually increase the boiling point of other hydrocarbons dramatically. And so even though London dispersion forces are the weakest, if you have larger molecules and you sum up all those extra forces, it can actually turn out to be rather significant when you're working with larger molecules. And so this is just a quick summary of some of the intermolecular forces to show you the application of electronegativity and how important it is.