If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

Main content
Current time:0:00Total duration:8:36

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 can 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 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 we have four electrons right in this double bond between the carbon and the oxygen so you're 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 is losing a little bit of electron density and this carbon is becoming partially positive like that and so for this 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 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 and 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 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 liquid and it has to do with the intermolecular force of dipole-dipole interactions holding those two 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 I think about these electrons here which are between the oxygen and the hydrogen I know oxygens more electronegative than hydrogen so oxygen is 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 were 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 at one time it was thought that this it was possible for hydrogen to form an extra bond and and that's that 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 is to see what the hydrogen is bonded to and so in this case we have 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 right 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 our fluorine oxygen and nitrogen and so the the mnemonic that students use is phone so if you remember phone as the electronegative atoms that could 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 right so higher than what we saw for acetone and this just is due to the fact that hydrogen bonding is a little bit this is a stronger version of dipole-dipole interaction and therefore 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 this is a these are the weakest intermolecular forces and they have to do with the the electrons that are always moving around in orbitals and and even though even though the methane molecule here if we look at it right we have a carbon surrounded by four hydrogen's for methane it's hard to tell and 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 right these hydrogen's 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 of course this one's nonpolar and so there's there's no dipole-dipole interaction no there's no hydrogen bonding the only intermolecular force it's holding to 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 let's say for the molecule on the left for the molecule at the left if we're 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 this for this molecule right the electrons can be moving in 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 its 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 somewhere around 20 to 25 obviously methane has already boiled and if you will and turned into a gas so methane is obviously a gas at room temperature and pressure so methane is a gas now if you increase the number of carbons you're going to increase the the number of attractive forces that are possible and and if you do that you can actually increase the boiling point of other hydrocarbons dramatically and so even though London version forces of the weakest if you have larger molecules and you sum up all of those extra forces it can actually turn out to be rather significant when you're working with larger molecules and so that's just a quick a quick summary of some of the intermolecular forces and to show you the application of electronegativity and how important it is