in this video we're going to focus on London dispersion forces which is one of the three kinds of intramolecular forces along with dipole-dipole forces and hydrogen bonding remember that all intimately intermolecular forces are IMF's are forces of attraction that pull separate molecules together okay so what do we want to do in this video here we want to talk specifically about the strength of London dispersion forces LDF we want to be able to compare two kinds of molecules and identify which of those molecules experience stronger LTFS and we want to be able to explain why one kind of molecule will experience stronger LDS than another so in order to do that let's examine two molecules that only experience London dispersion forces the two molecules we're going to look at today are f2 + CL - so f2 + CL - well if we look at the periodic table you'll see that fluorine and chlorine f and cl are both very similar to one another they're both in group 17 so therefore they both have 7 valence electrons and of course they're right next to each other in group 17 so therefore the lewis dot diagram that's going to result from these molecules is going to be pretty similar both bonded by single bonds and you'll notice because both the atoms in the bond or sorry both the atoms in the molecule are identical we're not going to have any polar bonds and because we don't have any polar bonds in the molecule the molecules will be nonpolar and as you know nonpolar molecules only experience London dispersion forces well if that's what's similar about these atoms what's different if we look back at the periodic table we'll see that each fluorine atom because it's atomic number nine has nine protons and nine electrons whereas each chlorine atom has 17 protons and 17 electrons so just looking at the board diagram for each atom notice this is a Bohr diagram for the atom not for the whole molecule will see that chlorine is very obvious that chlorine is much larger than fluorine it's got almost twice the number of protons and electrons as fluorine does and remember it's electrons that give an atom its size so one way we can show this for the molecule is by drawing a electron cloud diagram for both fluorine and chlorine so remember electron cloud diagrams show the relative location of the electrons in space and what they also show is that the electrons are zipping all around remember electrons are not stationary they're in constant motion so there's zooming all around and some of the electrons are shared as we'll see here but we can see the electron cloud for chlorine cl2 is much larger than the electron cloud for f2 okay so what does this have to do with London dispersion forces LDS well as you may remember the way nonpolar molecules experience intermolecular forces is when the electrons in the atom or sorry in the molecule zoom or off to one side than the other so in this case here the atom on the right just happens to have more electrons around it than the atom on the left and therefore what we have is a temporary dipole remember nonpolar molecules can have temporary dipoles if the electrons zoom over to one side so in this case here we'd have a temporary dipole pointing to the right and this atom on the right would have a partial negative charge therefore this would have a partial positive charge and notice the same thing for this chlorine cl2 molecule over here also there would be a temporary dipole pointing to the right with a partial negative and a partial positive charge because more of the electrons happen to have zoomed over to this atom on the right here now what's different well you'll notice the electron cloud for the co2 molecule is much larger than the electron cloud for the f2 molecule what this shows is that or electrons can happen to zoom around the atom on the right as compared to the f2 molecule so therefore the temporary dipole in this CL 2 molecule is going to be stronger than the temporary dipole in this F 2 molecule so let's write that down here temporary dipole is stronger what I mean by stronger is that this partial negative charge will be more negative than this partial negative charge because CL 2 has more electrons okay so how would this play out well if I another molecule of f2 were to get near this first molecule of f2 we would experience some LDF some London dispersion forces some forces of attraction between these two molecules here and same thing if another molecule of cl 2 got near this first molecule of cl 2 what i've tried to diagram on these other molecules here is really the strength of the dipoles so as you'll see the fluorine has this weaker dipole and i tried to diagram the CL 2 with a stronger dipole as you see a bigger arrow here so essentially just the force of attraction is going to be stronger here and just weaker here so what we're gonna have is a weaker force of attraction between the f2 to F 2 molecules as compared to the CL 2 s because the dipole is stronger in CL 2 okay so let's summarize things here first what can we say here molecules with more electrons are more polarizable I'm not sure if I've heard the term polarizable in any context except for this one here in chemistry so I'm assuming it's a made-up chemistry work but what is polarizable mean it means how polar we're able to make a nonpolar molecule so how much the electrons are able to zoom out so if we have more electrons more of the electrons are able to go on one side of the molecule than the other they are able to zoom over there so it's more polarizable so it's able to become more polar all right so molecules with more electrons are more polarizable and therefore they're going to have stronger partial negative and partial positive charges and they're gonna form stronger attractions between separate molecules and form stronger IMS in this case the intermolecular force that's being experienced is London dispersion forces okay so let's examine now some data to see if this plays out in reality and what we're gonna do is look at two molecules our f2 and CL 2 molecules and we're going to look at the boiling point remember when something boils it goes from liquid phase to gas phase and so what happens is that the intermolecular forces are completely broken so we can take a look at the boiling point to see how much energy it takes to completely break apart all the intermolecular forces in the compound so let's take a look here for f2 the boiling point is negative one hundred eighty eight degrees Celsius for CL 2 the boiling point is negative 34 degrees Celsius so both of these are cold temperatures but the boiling point for f2 is incredibly cold no we're outside of a laboratory are you gonna find temperatures on earth at least of negative one hundred and eighty eight degrees Celsius whereas thirty negative 34 degrees Celsius all those super cold you could find that on some on earth so therefore it doesn't take as much energy to boil f2 it takes a higher temperature this is a less negative number so it takes a higher temperature it takes more energy to break apart all the intermolecular forces in co2 let's go back to our periodic table and see some other elements in group 17 here well another element ripe and elk loin is bromine and B are two can bond very similarly to how f2 in CL 2 bond and same thing with iodine and we can get an I 2 molecule well the boiling point for BR 2 is 59 degrees Celsius and the boiling point for I 2 is 184 degrees Celsius so what we can see here is as our atoms get larger and larger and larger in terms of their number of electrons the boiling point is going to increase it takes more energy to break apart these molecules notice bromine has over twice the number of electrons as chlorine and iodine has a significantly larger number of electrons than bromine so what we can say here is it takes more energy to break I am yes I remember what's the reason why as we go from f2 - I - we have more electrons so the molecule is more polarizable okay so hopefully what we've been able to do in this video here is compare two molecules and identify which experience is the stronger into London dispersion forces and again it's going to be the molecule with the most electrons explain why that kind of molecule will experience stronger LDS is because the more electrons we have the more polarizable our molecule is and so therefore this stronger intermolecular forces it will experience