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:9:38
AP.Chem:
SAP‑5 (EU)
,
SAP‑5.A (LO)
,
SAP‑5.A.1 (EK)
,
SAP‑5.B (LO)
,
SAP‑5.B.1 (EK)

Video transcript

what we're going to do in this video is start talking about forces that exist between even neutral atoms or neutral molecules and the first of these intermolecular forces we will talk about are London dispersion forces so it sounds very fancy but it's actually a pretty interesting and almost intuitive phenomenon so we are used to thinking about atoms unless they say we have a neutral atom so it has the same number of protons and electrons and so that's all those are all the protons and the and the neutrons in the nucleus and then it'll have a cloud of electrons so I'm just imagining all these electrons kind of jumping around that's how I'm going to represent it and let's imagine and this is definitely not drawn to scale the nucleus would actually be much smaller if it was but let's say that there is an adjacent atom right over here and it's also neutral maybe it's the same type of atom it could be different but we're gonna say it's a neutral and it also has an electron cloud cloud and so if these are both neutral in charge how would they be attracted to each other and that's what London dispersion forces actually explain because we have observed that even neutral atoms and neutral molecules can get attracted to each other and the way to think about it is electrons are constantly jumping around probabilistically they're in this probability density cloud where the electron could be anywhere at any given moment but they're not always going to be evenly distributed you can imagine that there is a moment where that left atom might look like this just for a moment where most of or maybe slightly more of the electrons are spending time on the left side of the atom then on the right side so maybe it looks something like that and so for that brief moment you have a partial negative charge this is the Greek letter Delta lowercase delta which is used to denote partial charge and on this side you might have a partial positive charge because remember when it was evenly distributed the negative charge was offset by the positive charge of the nucleus but here on the right side because there's fewer electrons here maybe you have a partial positive on the left side you have where the most of the electrons are in that moment partial negative now what might this induce in the neighboring atom think about that pause the video think about what might happen in the neighboring atom then well we know that like charges repel each other and opposite charges attract each other so if we have a partial positive charge out here on the right side of this left atom well then the negative electrons might be attracted to it in this right atom so these electrons here might actually be pulled a little bit to the left so they might be pulled a little bit to the left and so that will induce what is called a dipole so now you'll have a partial negative charge on the left side of this atom and then a partial positive charge on the right side of it and we already had a randomly occurring dipole on the left-hand side but then that would have induced a dipole on the right-hand side a dipole is just when you have the separation of charge where you have your positive and negative charges at two different parts of a molecule or an atom or or really anything but in this world then all of a sudden these two characters are going to be attracted to each other or the atoms are going to be attracted to each other and this attraction that happens due to induced dipoles that is exactly what London dispersion forces is all about you can actually call London dispersion forces as induced dipole induced dipole forces there they become attracted to each other because of what could start out as a temporary imbalance of electrons but then it induces a dipole and the other atom or the other molecule and then they get attracted so the next question you might ask is how strong can these forces get and that's all about a notion of polarizability how easy is it to polarize an atom or molecule and generally speaking the more electrons you have so the larger the electron cloud larger electron cloud electron cloud which is usually associated with molar mass so usually molar mass then the higher polarizability you're going to have because you gonna have more electrons to play around with if this was a helium atom which has a relatively small electron cloud you couldn't have a significant imbalance at most you might have two electrons on one side which would cause some imbalance but on the other hand imagine a much larger atom or not a much larger molecule you could have much more significant imbalances three four or five fifty electrons and that would create a stronger temporary dipole which would then induce a stronger dipole in the neighbors that could Domino through the entire sample of that molecule so for example if you were to compare some noble gases to each other and so we can look at the noble gases here on the right-hand side if you were to compare the London dispersion forces between say helium and argon which one do you think would have higher London dispersion forces a bunch of helium atoms next to each other or a bunch of argon atoms next to each other well the argon atoms have a larger electron cloud so they have higher polarizability and so you're going to have higher London dispersion forces and you can actually see that in their boiling points for example the boiling point of helium is quite low it is negative two hundred and sixty eight point nine degrees Celsius while the boiling point of argon it's still at a low temperature by our standards but it's a much higher temperature than the boiling point for helium it's at negative one hundred eighty five point eight degrees Celsius so one way to think about this if you were at say negative 270 degrees Celsius you would find a sample of helium in a liquid state but as you warm things up as you get beyond negative two hundred sixty eight point nine degrees Celsius you're going to see that those London dispersion forces that are keeping those helium atoms together sliding past each other in a liquid state they're going to be overcome by the energy due to the temperature and so they're going to be able to break free of each other and essentially the helium is going to boil and you're going to enter into a gaseous state the state that most of us are used to seeing helium in but that doesn't happen for argon until a good bit warmer still cold by our standards and that's because it takes more energy to overcome the London dispersion forces of argon because the argon atoms have larger electron clouds so generally speaking the larger the molecule because it has a larger electron cloud it'll have higher polarizability and higher london dispersion forces but also the shape of the molecule matters the more that the molecules can come in contact with each other the more surface area they have exposed to each other the more likely that they can induce these dipoles in each other for example butane can come in two different forms it can come in what's known as n butane which looks like this so you have four carbons and ten hydrogen's so two three four five six seven eight nine ten this is known as n butane but another form of butane known as isobutane would look like this so you have three carbons in the main chain and then you have one carbon that breaks off of that middle carbon and then they're all they all have four bow bonds and the leftover bonds you could say are with the hydrogen's so it would look like this and this right over here is isobutane isobutane now if you had a sample of a bunch of n butane versus a sample of a bunch of isobutane which of these you think will have a higher boiling point pause this video and think about that well if you have a bunch of n butanes next to each other rather imagine another n butane right over here it's going to have more surface area to its neighboring butanes because it is a long molecule it can expose that surface area to its neighbors while the isobutane in some ways is a little bit more back it has lower surface here it doesn't have these big long chains and so because you have these longer and butane molecules you're going to have higher london dispersion forces they obviously have the same number of atoms in them they have the same number of electrons in them so they have similar sized electron clouds they have the same molar mass but because of n butanes elongated shape they're able to get closer to each other and induce more of these dipoles so just by looking at the shape of n butane versus isobutane you say higher london dispersion forces and n butane so it's going to have a higher boiling point it's going to require more energy to overcome the London dispersion forces and get into a gaseous state
AP® is a registered trademark of the College Board, which has not reviewed this resource.