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# Molecular polarity

AP.Chem:
SAP‑4 (EU)
,
SAP‑4.C (LO)
,
SAP‑4.C.1 (EK)
,
SAP‑4.C.2 (EK)

## Video transcript

now that we understand how to draw dot structures and we know how to predict the shapes of molecules let's use those skills to analyze the polarity of molecules using what's called the dipole moment so to explain what a dipole moment is let's look at this situation over here on the right where we have a positively charged proton all right some distance away from a negatively charged electron and let's say they're separated by a distance of D here we know that a proton and an electron have the same magnitude of charge so both have a magnitude of charge Q equal to 1.6 times 10 to the negative 19 so of course a proton would have positively charged Q so let's go ahead and make this positively charged Q and an electron would have a negatively charged Q like that if we were to calculate the dipole moment right the definition of a dipole moment symbolized by the Greek letter mu dipole moment is equal to the magnitude of that charge Q times the distance between those charges D so Mew is equal to Q times D and we're not really going to get into math in this video but if you were to go ahead and do that calculation you would end up with units of Dubai's so you get a number and that number would be in device here so we're more concerned with analyzing dipole moments in terms of the molecular structure so let's go ahead and look at the dot structure for HCl all right so if I look at this covalent bond between the hydrogen and the chlorine I know that that covalent bond consists of two electrons and chlorine is more electronegative than hydrogen which means that those two electrons are going to be pulled closer to the chlorine so I'm going to go ahead and show that here with this arrow all right the arrow is pointing in the direction of movement of electrons if you will so those electrons and yell are going to move closer to towards the chlorine so chlorine is going to get a little bit more electron density around it and so we represent that with a partial negative charge all right so I'm going to a lowercase Greek Delta here and it's partially negative since it has increased in electron density that's one way of thinking about it and since hydrogen is losing a little bit of electron density right it's losing a little bit of negative charge and so it is partially positive so go ahead and draw a partial positive sign here and so we're setting up a situation where we are polarizing the molecule alright so so so this part of the molecule over here on the right right is increasing an electron density and so that is our partial negative side so that's one Pole and then this other side here right is losing some electron density and so it's partially positive so we have it like that so that's where the positive sign comes in you can think about on this arrow here right this little positive sign giving you the distribution of charge in this molecule and so you have these two poles right a positive pole and a negative pole if you think about those two poles as having a center of mass you could calculate you can have a distance between them and you could calculate the dipole moment for this molecule and so when you calculate the dipole moment for HCl mu turns out to be equal to approximately one point one one devised and so we have a polarized bond we have a polarized molecule and so therefore we can say the HCl is relatively polar it has a dipole moment all right so that's that's kind of how to think about analyzing these molecules let's do let's do another one here let's do carbon dioxide alright so I know that the co2 molecule is linear so after you draw the dot structure right you're going to get a linear shape which is going to be important when we're trying to predict the dipole moment if I analyze the electrons in this carbon oxygen bond alright so we have a double bond between carbon and oxygen oxygen is more electronegative than carbon so oxygen is going to try to pull those electrons closer to itself and so we can go ahead and draw our arrow or our vector pointing towards the right here and so we have a we have a bond we have a bond dipole situation here on the Left we have the exact same situation right oxygen is more electronegative than carbon and so these electrons are going to be pulled closer to this oxygen so we draw another another arrow or another vector in this case so even though we have these individual bond dipoles if you think about this molecule as being linear and you can see we have these two vectors that are equal in magnitude but opposite in direction those two vectors are going to cancel out and therefore we would not expect to have a dipole moment for the molecule there's no molecular dipole here so Mew turns out to be equal to zero a simplistic way of thinking about this would be like a tug-of-war right you have these really strong atoms if you will right these oxygens but they're equally strong and if they're pulling with equal force in opposite directions it's going to cancel out so the individual bond dipoles cancel out and so there's no overall dipole moment for this molecule and carbon dioxide is considered to be nonpolar let's go ahead and analyze a water molecule over here on the right so B the electrons in this covalent bond between the hydrogen and oxygen oxygen is more electronegative than hydrogen so those electrons going to be pulled closer to the oxygen right same thing for this bond over here and we also have lone pairs of electrons on our central atom to think about and that's of course going to increase the electron density going in this direction for that lone pair and in this direction for that lone pair and so even though it's it's we know the geometry of the water molecule is bent and it's hard to represent that on this two-dimensional surface here if you use a Molly mod set you will kind of see that your your net dipole moment right would be directed upward in this case and so the the individual bond dipoles are going to add to give you a molecular dipole in this case pointed up and so therefore we can have a dipole moment associated with your water molecule so miu turns out to be approximately one point eight five and we can consider water to be a polar molecule all right let's do two more examples so on the left is ccl4 or carbon tetrachloride and so you can see that we have carbon bonded to chlorine here and so this is a straight line this means in the plane of the page if you will and so we know that geometry is tetrahedral around this carbon all right so let's go ahead and analyze that as well so I have I have a wedge drawn here which means this chlorine is coming out at you in space all right and then I have a dash back here meaning this chlorine back here is going away from you in space so that's how to think about it but it's really much easier to go ahead and make this using a Molly mod set and you can see that however you rotate this molecule it's going to look the same in all directions right so a tetrahedral arrangement of four of the same atoms around a central atom right key you can turn the molecule over it's always going to look the same in three dimensions and that's really important when you're analyzing the dipole moment for this molecule so let's go ahead and do that which we're start with our electronegativity differences right so if I look at this top carbon chlorine bond right these two electrons in this top carbon chlorine bond chlorine is more electronegative than carbon and so we could think about those electrons being pulled closer to the chlorines let me go ahead and use green for that right so those two electrons are going in this direction and it's the same thing for all of these chlorines chlorine is more electronegative than carbon so we can draw these individual bond dipoles all right we can draw four of them here and in this case we have four dipoles but they're going to cancel out in three dimensions so again this is a tough one to visualize on a two-dimensional surface but if you have the molecule in front of you it's a little bit easier to see that if you keep rotating the molecule right it looks the same and so these individual bond dipoles cancel there's no dipole moment for this molecule and so mu is equal to zero and we would expect the carbon tetrachloride molecule to be nonpolar let's look at the example on the right where we have substituted in a hydrogen for one of the chlorines and so now we have chcl3 or chloroform so now if we analyze the molecule alright so let's think about this bond in here all right two carbon is actually a little bit more electronegative than hydrogen so we can show the electrons in that bond in red moving towards the carbon this time and once again carbon versus chlorine chlorine is more electronegative so we're going to have a bond dipole in that direction which we can do for all of our chlorines here and so hopefully it's a little bit easier to see in this case right in this case the individual bond dipoles are going to combine to give you a net dipole located in the downward direction for this molecule so I'm tempting draw the molecular dipole right the dipole for the entire molecule going a little bit down in terms of how I've drawn this molecule and so since we have a hydrogen here right there's no upward pull in this case to balance out the downward pull and so we would expect this molecule to have a dipole moment and so mute urns out to be approximately one point zero one for chloroform so it is certainly more polar than our carbon tetrachloride example