formal charges that's what this next lesson in my brand new organic chemistry playlist is going to be about now this uh second lesson in this first chapter and this first chapter is all a review of gen chem we started off with lewis structures now we'll talk about formal charges we'll move on to hybridization and valence bond theory followed by molecular orbital theory which may not be so much review uh then we'll finish it off with a discussion of polarity and intermolecular forces now again this is my brand new organic chemistry playlist i'll be releasing this throughout the weekly throughout 2020 2021 school year and if you don't want to miss one of these subscribe to my channel click the bell notifications and you'll be notified every time i put one of these up okay so formal charges and uh you might recall that formal charges are defined a little bit differently than oxidation states they are not the same thing so uh but again this should be something you saw in your your typical general chemistry course so this shouldn't be brand new uh but there's a lovely little formula for formal charges and i'm going to present that formula and i'll work out an example with it and then we'll not remember that formula ever again because you probably don't remember it now either so but if you look at the way formal charge works uh we talked in the last lesson on lewis structures that you know typically when an atom is making its normal number of bonds it won't have a formal charge a formal charge of zero but when an atom is making something other than its typical number of bonds almost always there is one you know exception you might encounter but almost always it's going to end up with what we call a formal charge here so if you look you might recall that carbon having four valence electrons so is four short of a filled octet and therefore typically wants to make four bonds and in this case one two three four and that's why he's not gonna have a formal charge same thing with this carbon here one two three four bonds no formal charge there now oxygen's got six valence electrons that's too short of a filled octet and that's why you're typically going to see oxygen making two bonds like this one right here one two but this one on the other hand has only one bond and that should clue you and be like oh i bet that atom's gonna end up with a formal charge and that's gonna be the case here as well uh now there's a formula for formal charge and that formula says take the normal number of valence electrons so if we do this for oxygen here his normal number of valence electrons would be six and then you're going to subtract a sum and that sum is going to be half of the number of bonding electrons he has well he's only got the one bond and there's two electrons in that so we half of those two electrons so plus its number of non-bonding electrons and for this oxygen it's two four six and that's the formula for formal charge and it's a fairly formal formula for formal charge here so if we look at this uh one half of two is one and one plus six is seven and six minus seven is negative one and so right up next auction here we'll typically put a minus sign to show you that it's got that negative formal charge now we'll find out that none of the other atoms here actually have any sort of formal charge so like this oxygen here would be six minus half of these four well half of four is two and then these four on top so plus another four and so it'll be half of four plus four which would be a total of six and six minus six comes out to zero which is why he's got no formal charge cool we'll find out carbon here also no formal charge so in this case all he's got is eight bonding electrons and if you take his normal number of four minus half of the eight which is four and four minus four is also zero but again any atom that's making its typical number of bonds is not going to end up with a formal charge it's when they have an abnormal or non-typical number of bonds that you should usually clue in and say oh probably gonna have a formal charge and do that formula now rather than using this formula i hate this formula and it's not that it's a bad formula it's just that it's hard to remember so much easier to remember instead of doing normal number of valence electrons minus the sum of half the bonding electrons plus the non-bonding electrons what you can really do is just say it's the normal number of valence electrons minus the dots and lines i'm not even going to recognize the dots or lines as electrons just say minus the dots and lines so for example in this auction right here we'd say his normal number is still six and then we're going to subtract the sum of the dots and lines and in this case he's got six dots around him and one line and six dots plus one line is a total of seven and six minus seven gets us once again negative one and so i like to just say like six minus one two three four five six seven notice i count this as a line i don't count it as two electrons or anything i just count it as a line and before we just counted half of the electrons in that bond whereas here i'm just counting as a line and so six minus dots and lines still lets me arrive at negative one for carbon over here let's say four minus one two three four lines and no dots four minus four is zero for this auction up here six minus one two three four five six is also zero and so it's a very quick way and also a very easy formula to remember for getting formal chart now if we look at the second example here so carbon is making four bonds and that's his typical number and he's not gonna have a formal charge because his normal valence for minus four lines is zero no formal charge but nitrogen with five valence electrons usually expect nitrogen to make three bonds this one's making four that should clue you in again that he's got a formal charge and so in this case his normal valence is five in this case minus one two three four lines and no dots so five minus four is plus one and he's got a positive formal charge which we signify just like so cool we'll finally move on to this last example and in this case uh chlorine has seven valence electrons and so being one short of a filled octet we typically expect chlorine to have one bond and all of these chlorines have one bond and if we actually formally go through the process here we say seven the normal number of valence electrons minus one two three four five six seven so one line six dots and seven minus seven is zero so making the normal number of bonds and end up with no formal charge but aluminum on the other hand you might recall that aluminum was one of our octet rule exceptions so boron aluminum often go under the octet rule and only have six electrons around them not eight well this aluminum here has got eight around him instead so when aluminum and boron only have six around them they're only making three bonds this one's got four and when you've got an abnormal number of bonds you're typically going to end up with a formal charge and so if we take aluminum's number of valence electrons which is three and subtract the number of lines four and no dots and three minus four gets us negative one cool and that's i highly recommend how you remember the formula for formal charge so and some people you know use balls and sticks or you know i like dots and lines personally so but normal valence minus dots and lines life is good but you definitely have to be able to assign formal charges and to do it pretty quickly for your typical first exam in organic chemistry so if you've benefited from this lesson please give me a like and a share and if you're interested in practice problems involving formal charges and if you're looking for the study guides that go with this course check out my premium course on chatsprep.com happy studying