polarity and specifically how to identify when molecules are polar versus non-polar versus ionic uh that's going to be the subject of this lesson now we're in chapter one of my brand new organic chemistry playlist and this is a gen chem review and it ends with a discussion of molecular properties and we've got to talk about polarity because in the next lesson we're going to talk about intermolecular forces and specifically dipole-dipole forces are for polar molecules so we're going to identify when they're polar and when they're not now again this is my brand new organic chemistry playlist i'll be releasing lessons weekly throughout the 2020-21 school year so if you don't want to miss one subscribe to the channel click the bell notifications you'll be notified every time i release one okay so discussion of electronegativity is at the heart of polarity and uh oftentimes we you know gave you some working definitions uh uh of how to identify if something's ionic or if a compound is covalent and if you got uh you know pure covalent bonds or non-polar covalent bonds and stuff well in gen chem we gave you some really generic you know kind of rules mice and metal and non-metal ionic two non-metals covalent uh and then we just said two non-metals usually if they're different but not carbon and hydrogen well they're probably going to be polar okay well we've got a more exact definition and you're not gonna have a table of electronegativities on you at all times but i do want to point out a couple key things so you see some differences now uh when we gave you those definitions of like metal and non-metal ionic and two nonmetals covalent stuff like that we were oversimplifying it a little bit and we're kind of lying to you in a couple places there are exceptions to that general trend and i want to point a couple of those out that'll be important and significant for organic chemistry so if we look at electronegativity and i've put the polling scale up on your screen there so the way it works is the bigger the difference in electronegativity the more polar a bond gets now if that difference in electronegativity reaches 1.7 or greater it's now become so big that it's so polar we don't even call it polar covalent more we call it ionic at that point but a difference in electronegativity that's less than 1.7 is going to be covalent now anywhere in the range of 0.5 to 1.7 that's going to be a polar covalent bond and then anything less than 0.5 for a difference we'll call that either non-polar covalent or pure covalent all right so let's take a look at some bonds here and i want to center this around carbon since this is organic chemistry and if we just look at two identical atoms well that's going to be a non-polar bond because if you've got two identical atoms that's just a difference of zero and that's about as pure covalent or nonpolar covalent as you can get so but take a look at even a carbon-hydrogen bond it's very important that you recognize that this is considered a non-polar bond as well so if you look carbon's at 2.5 hydrogen's at 2.1 and so that's just a difference of 0.4 less than the 0.5 threshold to be called polar so that's considered a non-polar covalent bond that'll be super important we start looking at properties of molecules in the next section cool so let's look at some polar ones here and let's take a look at the carbon nitrogen bond and for a carbon nitrogen bond carbon's at 2.5 nitrogen is at 3.0 and we're right at the threshold to be considered polar covalent and if we kind of moved our way up to carbon oxygen bond now 2.5 compared to 3.5 1.0 definitely right in the middle of the polar covalent range and if we move this a little further carbon to fluorine so again carbon at 2.5 fluorine at 4.0 on the polling scale a difference of 1.5 and we're starting to approach the threshold of what's considered polar covalent if this bond was just a little more polar it might be considered ionic cool now let's kind of go textbook ionic compound from here and so if we go textbook ionic compound let's say we choose sodium chloride so and if we look at the polling scale there for sodium chloride you'll see that chlorine is at 3.0 and sodium i got to look that one up myself is 0.9 so 0.9 to 3.0 we've got a difference of 2.1 past that threshold and indeed we've got an ionic bond and the bigger the difference in electronegativity is here uh you know the more ionic character we'd say that has and notice we've got a metal and a non-metal and that works out great and most of the time in gen chem so you're not going to have you know all these uh electronegativity values available to you and so we just gave you this simple rule to remember we said metal nonmetal ionic two nonmetals covalent but i want to point out a couple exceptions there and in fact just one really major exception so and i'm going to give you the carbon magnesium bond now carbon's a non-metal magnesium is a metal and in gen chem we would have considered we would have you know trained you to consider this as an ionic bond but if you actually look at the electronegativity value you see that carbon is at 2.5 and magnesium's at 1.2 that is only a difference of 1.3 it's a polar covalent bond and it's a pretty polar polar covalent bond but it technically is not ionic and in second semester we're going to deal with a reagent called a grignard reagent which definitely looks at a carbon magnesium bond as a polar covalent bond so definitely want to kind of cover the fact that we've got some exceptions there now what are you supposed to take away from this well the truth is this again you're still not going to have a table of electronegativities in front of you at all times so what we're going to expect of you is when you're dealing with two non-metals i expect you to know that two identical atoms nonpolar carbon dihydrin non-polar most the time outside of that though if you've got two different non-metals you should probably just kind of you know go with your gut and say that's probably gonna be polar so sweet and then most of the time metal and non-metal it's going to be ionic although we won't encounter that a whole lot in organic chemistry we'll be dealing mostly with covalent bonding and stuff like that so so even that rule of thumb we gave you in gen chem it's still pretty useful to you here in ochem but i did want to point this one bond out because we will play with that a little bit in second semester so now we're going to take a little time identifying when molecules are polar versus nonpolar and also kind of being able to rank relative polarities as this again will be important to our discussion of intermolecular forces in the next lesson all right so first thing you should look for when you're trying to identify if a molecule is polar is you're going to look for polar bonds so if you look at this bottom example here so you've got carbon carbon bonds those are definitely non-polar as we just stated and carbon hydrogen bonds we also so those are below the threshold and those are considered non-polar and so for a molecule that doesn't have any polar bonds it is simply going to be a non-polar molecule so first thing i'll look for is just look for polar bonds so cool but if we look at the example of carbon dioxide here so in this case you want to look for polar bonds and a carbon oxygen bond is indeed polar we'll find out that actually double bonds are significantly more polar than single bonds when they're polar anyways in this case this guy's got polar bonds however that's not enough you then have to look at the orientation of those polar bonds to see if they cancel or not so to speak so if you look at that carbon oxygen bond we often represent a bond dipole like so and we've got another one right on the other side and this carbon's sp hybridized the bond angles are 180 degrees and so if you add these two bond dipoles together as vectors the vector sum is going to be zero and we say that they cancel and as a result this molecule is going to be nonpolar cool now if you look at the difference here with the one i've got below so very similar molecule but one of the oxygens is replaced by a sulfur and the carbon sulfur bond there's a smaller difference in electronegativity and a smaller difference in electronegativity is going to lead to a less polar bond and so as a result because these are not equal in polarity notice i've driven kind of drawn a smaller arrow on purpose there with a smaller arrow so or a smaller bond dipole anyways these are not going to add up to zero and as a result overall this molecule is going to be polar and in fact we can after drawing these individual bond dipoles you add them up it's still gonna overall the molecule is still going to have a bond or i'm sorry a dipole overall pointing to the left and that overall dipole would look like so and you might be asked to draw an overall molecular dipole so these are individual bond dipoles and then in blue there's an overall molecular dipole cool so look for polar bonds you don't have any it's non-polar if you do though you have to see if they add up to a vector sum of zero if they cancel if they do that's a non-polar molecule but if they don't add up to zero as the case here that's going to be a polar molecule cool but in addition to just being able to see if a molecule is polar or not you might have to compare relative polarities which is why i've got this series of molecules up here and relevant to this discussion a measure of polarity is what's often called the dipole moment it's symbolized by the greek letter mu right here and it's equal to the difference in electronegativity or the i shouldn't really say the difference electronegativity but the partial charges that result from the difference in electronegativity so times the distance of separation between those two atoms so in this case a bigger partial charge which is the result of a bigger difference in electronegativity is going to lead to a more polar bond and that's the part of this definition we really want to focus on here so all right if we look at this series of molecules we'll start with methane here and there's no polar bonds in this whatsoever and therefore the overall dipole moment is zero it's a nonpolar molecule okay no no surprise there but in this next example that carbon chlorine bond that's a polar bond chlorine is significantly more electronegative than carbon and so having a polar bond we have to ask yourself okay are there multiple polar bonds and do they cancel well no that's the only polar bond which means there's nothing to cancel it out and overall the entire molecule has a dipole running right up the center of it the bottom part of this molecule the way i've drawn is partially positive the top part's partially negative that's what happens in a polar molecule one side will be partially positive and other side be partially negative cool bond dipole here is 1.87 du by and that's just a common you you use for uh dipole moments i wouldn't worry too much about it in this case we look at the next example here now we've got two electronegative atoms two chlorines but you might notice that the dipole moment actually decreases from one point eight seven dubai to 1.6 to buy so and if i ask most of my freshmen and sophomores like which one of these is more polar nine out of ten students are going to choose the second one to be more polar like it's got two chlorines chad come on so well let's take a look at it investigate and see why indeed it actually is less polar well there are two polar bonds we can draw in those bond dipoles so and when you've got multiple bond dipoles you have to add them together and in this case if they point the same direction that increases the polarity when they point exactly opposite directions they cancel each other out and stuff like this as long as they get equal polarity associated with them as they do in this case so the question is are these pointing more in the same direction then they're going to be additive and the polarity should increase or are they pointing more in opposite directions then that should cause a decrease in polarity so in this case the bond angles are 109.5 this is an sp3 hybridized carbon here and being 109.5 so notice 90 is the threshold if you have a bond angle smaller than 90 between them they'd point more in the same direction and then to add together if the bond angle as it is in this case is bigger than 90 degrees then actually these cancel each other out more than they add and that's why the overall dipole moment of this molecule went down now if we want to draw the overall molecular dipole moment here it'd just be the average of these two which would run right down the center of them and again you could be expected to draw an overall molecular dipole like this guy now if we look at this next one here we've got three chlorines now and if you notice our dipole moment went down even further 1.01 dubai and we can draw those three bond dipole moments and again the key is every angle between any two chlorines you choose is 109.5 they're all bigger than 90. and so they cancel each other out more than they add all three of them and your dipole moment overall is going down a little further and if you take the average of all three of these individual bond dipoles you have a dipole moment running right down the middle of this molecule right in between all three of these so to speak a little hard to draw from this perspective cool finally now you get the carbon tetrachloride here where you've got four carbon chlorine bonds and now they actually add up to a vector sum of zero just like we saw in carbon dioxide up here these four all add up to zero and that's why now this is a non-polar molecule with a dipole moment of zero cool and again in summary the big point of this lesson is for you to be able to identify when molecules are polar versus nonpolar and also maybe to compare some relative polarities in certain cases so again one of those big intermolecular forces we're going to talk about the next lesson are dipole-dipole forces and the key is polar molecules we'll have them and non-polar molecules won't and so being able to predict when they're polar and when they're not or which is more polar will be an important skill you'll need to have for the next lesson if you found this lesson helpful consider giving me a like and a share and if you're looking for practice problems and study guides to go along with this check out my premium course on chatsprep.com happy studying