8.2 How to Draw Lewis Dot Structures | Complete Guide | General Chemistry

how to draw lewis dot structure is going to be the topic of this lesson and it is going to be a comprehensive lesson it'll be a little bit long i'm going to warn you uh because we're going to make sure that you get a chance to look at every little nuance of every little rule possible for drawing these lewis structures we're going to talk about the octet rule we're going to talk about the three major exceptions to the octet rule including expanded octets uh we're going to talk about resonance structures for those that have resonance we'll talk about formal charges and how we can use formal charge to distinguish between some of these resonant structures we are going to work a ton of examples and this is super important this is foundational for the entire next chapter as well which is going to deal with like molecular geometry so you really got to have a good handle on these lewis dot structures for any of that to really make sense my name is chad and welcome to chad's prep where my goal is to take the stress out of learning science now in addition to high school and college science prep we also do mcat dat and oat prep as well you can find those courses at chadsprep.com now this lesson's part of my new general chemistry playlist i'm releasing several lessons a week throughout the school year covers an entire year of general chemistry and if you want to be notified every time i post a new lesson then subscribe to the channel click the bell notification all right so we got to talk about lewis dot structures and uh lewis dot structure just a way of representing the valence electrons and if you recall the valence electrons of the outermost shell uh and the reason they get their own special name so everything closer than the valence is just called chord you know regardless what shell they're in but the outermost shell gets its own special name distinguishing it from everything else the valence because they are the ones involved in chemical reactions in the making and breaking of bonds and so we're going to see that sometimes these electrons are shared uh and things of this sort and we will draw that as a covalent bond now in the last lesson we dealt with ionic bonding and we are going to deal just a smidge with ionic bonding in this lesson but it's mostly going to be about covalent bonding and representing molecules with these lewis dot structures so just as a reminder for the number of valence electrons in an atom so on the periodic table here your group one metals have one valence electron your group two metals have two valence electrons your group three elements here uh or group 13 depending on which convention you're using have three valence electrons carbons group four valence electrons nitrogens group five valence electrons oxygen's group six valence electrons the halogens seven valence electrons and the noble gases eight valence electrons except for helium which just has two now one thing to note because we are really going to be heavily focusing on molecular compounds which are made up of mostly non-metals we are mostly going to be dealing with these elements right up here so yeah i said that the group one metals the alkali metals have one valence electron and that'll be relevant for like two minutes at the beginning of this lesson and then we probably just won't talk about metals much at all through the rest so i'm not saying they're completely unimportant but this lesson is heavily going to focus on the non-metals as we make molecular compounds here so but we will start with something that does have sodium and sodium is a group one metal and it has one valence electron the way this works since atoms can typically have up to eight valence electrons in most cases we're going to split those up that you can draw them as dots on four different sides of an atom so and you put one on each side before you start pairing them up well in sodium's case it just got one and it doesn't matter if you put it on the top or the bottom or the left or the right so and i'm just going to put it on the right here and there's his one valence electron now chlorine on the other hand being a halogen has got seven valence electrons and so in this case we're going to put one on each side and then start pairing them up as well and so these three sides each have a pair and then this one's just got that unpaired electron right there and so these dots represent the electrons cool now in the case of a metal and a non-metal so we'll find out that uh uh the electrons are going to organize themselves in such a way that everybody gets what's called a filled octet and we call this the octet rule and so in one way shape or form atoms are either going to transfer electrons in the case of a metal and a non-metal which would be ionic bonding or they're going to share electrons as we see two non-metals doing over here in an attempt to get these filled octets to feel like it has eight valence electrons in its outermost shell so in the case again of a metal and a non-metal you've got a metal with a low ionization energy you've got a non-metal with a rather high ie very negative electron affinity and so the transfer of electrons make sense here so in this case so we are just going to transfer an electron over from sodium over to chlorine in order to satisfy this octet rule and so effectively chlorine has just stolen sodium's electron and that's going to leave chlorine with a negative one charge and sodium with a positive one charge and sodium is not exactly upset with chlorine though that that chlorine stole this electron because now sodium is positive and chlorine is negative and sodium's like well chlorine is kind of cute actually and so they're going to hang out together having opposite charges that's kind of the nature of an ionic bond and so uh for ionic bonds we don't actually draw a line between the metal and the non-metal or anything like this we'll find out that drawing a line like that specifically designates a covalent bond so ionic bond is just by nature of having a plus charge and a minus charge a cation and an anion all right so what if we've just got two non-metals now and once again these have got seven valence electrons each and so they both have a problem they're both one short of having a filled octet and in this case you know the chlorine on the left here tells the other chlorine hey give me your electron and this chlorine says no you give me your electron and so the problem is that neither one of them has a low ionization energy like sodium and so neither one is going to relatively easily lose an electron they in fact both want to gain an electron and that's kind of the predicament you're in when you've got two non-metals is you've got nobody who can easily lose an electron nobody with a low ionization energy and so with the non-metals that's why we're going to share electrons in an attempt to get a filled octet and so what we're going to do is share each of these guys one of these electrons in the middle and so we can write this a couple of different ways we can put those shared electrons right in the middle so and now we've got a couple of different designations here the ones that are being shared are called bonding electrons so we've got two bonding electrons and then the ones out here are called non-bonding electrons they're not being shared in any way shape or form so these non-bonding electrons are often called lone pairs as well so these lone pairs of electrons this chlorine on the left has three lone pairs of electrons or three lone pairs of non-bonding electrons and the one on the right's got three lone pairs of non-bonding electrons as well and generally bonding electrons are going to be lower in energy than non-bonding electrons and that's the whole nature of why these atoms are going to bond together so by having bonding electrons whereas before they just had other non-bonding electrons they weren't being used to make a bond before so in coming together though it lowers the energy of their electrons and that is what is the driving factor in forming a bond it's an exothermic process for uh making a bond as we'll find out it would be an endothermic process to break the bonds but it's energetically favorable to make bonds all right so another way we can represent those shared electrons though and quite commonly what we'll do is we'll draw a line to represent those bonding electrons and so that line represents a covalent bond just like the two dots in between the two chlorines also represented a covalent bond so but here uh just another way of drawing it but you're supposed to see that every line right there every bond in this case every covalent bond specifically represents two electrons and we'll find out that in some cases we'll have double bonds and triple bonds where we'll end up having more than just two shed electrons we could have four shared electrons which would be a double bond or six shared electrons would be a triple bond so when we'll find out you know how do you know when you need a double bond or a triple bond well we'll have some rules for drawing lewis structures that'll help us figure exactly that out now we talked about this octet rule and now we've seen its utility in predicting structures either the transfer of electrons for an ionic compound to you know have both getting a filled octet notice chlorine here has got 2468 around it and sodium by giving one away it's got its previous second shell still full and so he's got a filled octet as well and same thing here with these chlorines both of them get to count the shared electrons so the chlorine on the left says i've got two four six eight and this chlorine on the right says i've got two four six eight and everybody is happy i.e everybody's got a filled octet now you should know that there are three major exceptions to the octet ruling so one not everybody follows the octet rule some are going to go under the octet rule and if you recall we talked about helium at the beginning of this lesson that uh is the only noble gas here that does not have eight valence electrons it just has two so and so if you drew the lewis structure for helium he would not be getting a filled octet he would be full in his case the first shell only has that one s orbital and so it can only hold two electrons but it's kind of the exception of the rule right so but hydrogen's therefore going to be the same case then hydrogen in order to be full wants to look like it's nearest noble gas which is helium and so as a result hydrogen really only wants to have two electrons around him not eight as the case may be so whereas if you take a look at something like you know carbon nitrogen oxygen fluorine generally each of those are going to be found following the octet rule most of the time so but not hydrogen hydrogen's never going to follow the octet rule when hydrogen is full it just has two electrons now one other thing we should realize is that uh we're going to run some issues here with beryllium and boron and aluminum as well and one thing you realize first of all is that beryllium's somewhat metallic and aluminum is fairly metallic and we're like um they don't really form molecules yet they're not they don't form molecular compounds because metals and non-metals form ionic compounds and again in the last lesson we kind of looked at some example in the first lesson this chapter we look at some examples where the electronegativity difference really is a better measure of whether we have ionic or covalent bonding going on it turns out aluminum and beryllium actually can be involved in covalent bonding now it's going to be polar covalent bonding but it's still going to be covalent bonding and so it is appropriate for us to talk about them a little bit in this context and so it turns out beryllium here has two valence electrons and the way this works is typically he's going to share one with one atom and then share one with another atom in the process make two bonds so in the case of sharing here like the chlorines here both atoms are typically going to contribute one electron each to share and that's what forms a bond and so with beryllium having two valence electrons he shares one with one atom over here he shares the other with another atom over here and then he's out of electrons and he can't make any more bonds and so at that point beryllium would just be bonded to adam you know two other atoms and i'll just call him adam x at this point on either side and end up with only a total of four electrons around him and so he's not going to get be able to get a filled octet he just doesn't have enough electrons to share and so in the case of beryllium he's only going to want four electrons to be full in a typical sense and then boron and aluminum will follow a similar thing so boron and aluminum both have three valence electrons and so they can share with three different atoms and so in this case we might share with atom x there and atom x there and atom x there and at that point so boron's kicking in one electron atom x whatever it is kicking an electron and boron's out of electrons he can't make a fourth bond he has nothing left to donate or contribute to make that bond and so boron's gonna be left with two four six electrons around him to be full and so it's pretty typical for boron and aluminum to therefore have six electrons rather than a filled octet now it turns out that born aluminum can be found with a filled octet it's just they don't have to be much more commonly going to be found with six so in fact we can take this a step further we can actually predict the normal number of bonds an atom is going to make and usually it's about how many electrons that are short from being full which again usually means a filled octet so but again we've got this octet rule and it says that again hydrogen might only get two not eight beryllium is typically only going to get four to be full not eight and then boron aluminum typically six to be full and not eight as well so that's your first exception to the octet rule your second exception to the octet rule is what we call an expanded octet it's atoms that go over the octet rule and you you kinda gotta realize where the octet rule comes from first and the octet rule comes from the fact that a typical shell of electrons is going to have a single s orbital and then three p orbitals and you can put a maximum of eight electrons in that outermost shell well as you guys well know from chapter six there's not just s orbitals and p orbitals there's also d orbitals and f orbitals but you only start getting those d orbitals in shell number three and you only start getting those f orbitals in shell number four and so it turns out only for elements that are you know have their valence in shell number three or higher so which is going to end up meaning uh in the third row of the periodic table or lower only those can potentially throw some electrons in the d orbitals as well and so you've got room to fit 10 more potentially and stuff like that for these d orbitals but again there's no such thing as a 1d orbital or a 2d orbital the first time we ever see the d orbitals is in shell number three and so it's only for these elements in shell number three here aluminum silicon phosphorus sulfur chlorine argon or below that can exceed this octet rule and so you should know they don't have to exceed the octet rule have this expanded octet but they can and so we'll find out that fairly commonly sulfur can be found making two bonds just like oxygen and we'll explain why that is in a minute and with two bonds and two lone pairs it would follow the octet rule however it's not uncommon to see sulfur making like six bonds in which case that would be 12 electrons and he's definitely going over the octet rule so whereas oxygen doesn't have that capability in the second period here boron carbon nitrogen oxygen fluorine there are no two d orbitals there's only s's and p's and so they are limited at eight electrons in that second period so it's only when they get into that third period and below that you can have this expanded octet now finally the third exception and the rarest of the three for that octet rule is if you just simply have an odd number of electrons and the most common example is no so and if we look here nitrogen's got five valence electrons right here whereas oxygen's got 6 valence electrons and so it turns out that no would have a total of 11 valence electrons well the problem is is you can't have these all existing in pairs when you have an odd number of electrons and so there's no way for both atoms to say they both have eight an even number if you have an odd number total and so there's no way this is going to satisfy the octet rule and we'll we'll see an example of that at some point in time so but that's the least common example having an odd number of electrons and typically what you'll end up doing is preferentially giving electrons to the more electronegative atom first as we'll see so now that we've got the octet rule and all three notable exceptions uh down we're just going to start following the rules for drawing lewis structures and we're just going to work our way through examples best way to learn these rules is to actually apply them in a variety of examples okay so i hinted at this earlier but the octet rule also allows us to kind of imply how many bonds an atom is normally going to make to get that filled octet so if we take a look at the halogens here the halogens here have seven valence electrons and so they're only one shy of having a filled octet and so what they're going to really need to do is just share one bond with an atom that way they chip an electron the other atom will chip in an electron and they'll get to credit that other electron as being theirs now now they'd have eight and so you can kind of look at how many electrons are they short from being full which again usually means filled octet and that's usually how many covalent bonds you'll find it making so the halogens being just one electron shiva filled octet you'll normally find them making just one bond so the chalcogens here oxygen sulfur selenium tellurium being two electrons short of a filled octet of a noble gas configuration are typically going to be found making two bonds so nitrogen's column here with five valence electrons that's three electrons short of a filled octet and so they're generally going to be made found making three bonds uh to get a filled octet carbon and silicon here germanium they are have four valence electrons that means they're four electrons short of a field octet and will be most often found making four bonds to get that filled octet now if you recall boron and aluminum were exceptions so they have three valence electrons but they're not trying to get a filled octet in fact you can't make more than four bonds in the case of boron here unless you're going to have again an expanded octet again the period two elements can't have that expanded octet so but for boron aluminum they're usually trying to get just six electrons to be full not eight and as a result with three valence electrons they're just three short of being full and are going to be typically found making three bonds over at beryllium here so brilliant's got two valence electrons and again he's not trying to get a filled octet he's usually just trying to be full at four total electrons and so having two he's just too short and therefore you can predict that he'd typically be found making two bonds and that's true now the reason this is important is it's got some explanatory power so we are typically going to have a central atom in most of the molecules we're going to look at and with the atom in the middle is going to be the one that's bonded to everything else around it and so typically the one in the middle is going to be the one that can generally make the most bonds and now that we know how many bonds things make we can kind of you know got some predictive power here so the way this generally works and often the way it's presented is that the least electronegative element and again reminder that fluorine is the most electronegative element on the pivot table generally the least electronegative element is the one that can make the most bonds and would therefore go in the middle of the molecule now you need to realize that hydrogen is going to be an exception here because hydrogen's less electronegative than say carbon but hydrogen's only going to make one bond it's got one valence electron it only wants two total so it's only going to make one bond and so don't put hydrogen is never going to be a central atom so usually the least electronegative atom but not hydrogen is going to be in the center now it also works that way as you go up a group and stuff like that as well it turns out so if you've got more than one halogen in a molecule it's usually the bigger one that's less electronegative that's going to be in the middle not the smaller one so just something to keep in mind here so if we say the least electronegative element goes in the middle except not hydrogen that usually will suffice to get you there and the idea is that the one in the middle needs to make the most bonds we need to put the one in the middle that can make the most bonds all right so if we take a look at ccl4 here the chlorines can each just make one bond each usually so is what we'll find again they have a chance of having an expanded octet so we can't you know rule that out completely but it's typically not gonna be the case we'll find out to have an expanded octet chlorine needs to be that central atom and he's not going to be in this case but the chlorines usually make one bond carbon can make four bonds with four valence electrons carbon is also much less electronegative and so we're going to put carbon in the middle and so the first thing you want to do is set up a skeleton for your molecule and so i'm going to put carbon in the medical in the middle in the middle and we're going to connect him to those four chlorines so he's the central atom and these guys are on the outside we've now drawn in eight shared electrons ie4 bonds so and then we've got our skeleton set up so next thing you want to do is fill up the octets of all the outside atoms first and so assuming they want to fill the octet and in this case the chlorines do they're not any of the exceptions that go under the octet rule like hydrogen beryllium or boron and so we're just going to fill them up you'll always have enough electrons to fill them up so we set up our skeleton put the least electronegative atom in the middle so it filled up the outside atoms octets and from this time we now need to take an accounting and usually what we'll actually do is count up the valence electrons that are going to be in our structure before we actually even get started so but in this case uh the only electrons that show up in your lewis dot structure again are those valence electrons so pretty good idea if you just figure out how many you've got to begin with so we've got four for carbon we've got seven for each of the four chlorines four times seven is 28 plus the four for carbon gets us 32 valence electrons and so once you fill up those outside atoms you've got to ask yourself a question is do you have any electrons valence electrons specifically left to go well in our case we got 2 4 6 8 2 4 6 8 2 4 6 8 2 4 6 8. so that's 8 times 4 is indeed 32 and we are out of electrons now because we worry about filling up the octets of the outside atoms first we're going to end up in situations where the central atom did not get a filled octet well in our case this carbon says well i got 2 there 2 there 2 there 2 there that's a total of 8 around me and he's happy and so the next thing you want to do after you fill up the outside atoms is you want to check to see if that central atom is happy which really just means full and normally it's going to be filled octet if they're not one of the exceptions in this case he is and because that is the case we are done this is the structure for ccl4 and in this case there are eight again shared electrons and there are a whole boatload here so we got six here six here six here six here not so uh 24 total non-bonding electrons on those chlorines so but that is our lewis dot structure we're going to start off simple here we are going to ramp this up and make this harder as we go so next we're going to take a look at is nf3 so i will also tell you that normally you're going to find out that the first element is going to be the one that's less electronegative and is usually the one that's going to go in the middle uh in this case that's definitely the case here as well notice nitrogen is definitely less electronegative than fluorine so we'll put the nitrogen in the middle and in any three of the four directions we've got to draw the fluorines i could have put one of them up i just randomly chose to go left right and down my choice all right so before we get any further into this let's just add up our valence electrons and again nitrogen based on where he's out on the periodic table's got five and then each of the three fluorines has seven each and seven times three is 21 plus the five is going to get us up to 26 valence electrons all right so we got 26 to work with notice we've already used 2 4 6 and once you've got your skeleton set up then fill up the outside atoms first that's we're going to do fill up those fluorines again you will always have enough to fill up your outside atoms notice you could also look at this as worrying about filling up your more electronegative atoms first as well because they're the outside ones so electronegativity is something to do with you know how how much they like to pull electrons towards them and so the ones that like the electrons the most so to speak are the ones we're giving the electrons to first if you prefer to think of it that way all right so now we've got to take an accounting we filled up those outside atoms and the question is do we have any electrons left so 2 4 6 8 2 4 6 8 2 4 6 8. and 8 times 3 is 24 and we've got electrons left and if you fill up the outside atoms and if you ever have any electrons left they always go on the central atom and you put them on in pairs well we've got two electrons left out of those 26 and so we'll put them on the only side there's no electrons on that nitrogen we'll put them on as a pair and so in this case we're now out of electrons we've used all 26 and again once you run out of electrons you have to ask yourself is the central atom happy ie is the central atom full and in this case 2 4 6 8 that central atom is indeed full he's got three pairs of bonding electrons and one lone pair of non-bonding electrons so and this is the structure of nf3 so nice filled octet for every atom all right next one on the list here is hcn and when you don't have a binary uh molecular compound but with more than two elements here usually they actually put the one in the middle that goes in the middle in that case this is carbon you could all say the least electronegative but not hydrogen is going to go in the middle and that's still carbon and so we're going to set up our skeleton here like so and we should count up our valence electrons so hydrogen has just the one carbon's got the four and nitrogen's got five and five plus four plus one is ten valence electrons and what we're gonna do is fill up the outside atoms first and you might want to start putting some electrons around the hydrogen you might be inclined to start doing this and then hopefully you realize pretty quickly that that's a bad idea hyden does not want to fill the octet if you recall hydrogen just wants to look like helium and have two electrons well before i drew any of those other dots we already had two electrons they're both bonding but we had two electrons hydrogen's already full no electrons to give to him so we'll go to the other outside atom the nitrogen and we'll fill him up as well because he's not yet full so now he's full and the question is do we have any electrons left well 2 4 6 8 10. we've used them all and so we're out of electrons and the moment you run out of electrons the question you have to ask yourself is is the central atom happy and in this case carbon is not happy he does not have a filled octet he's not full he's only got two four electrons around him he wants to have eight and so but again we can't just put lone pairs on them because we've already used all 10 electrons so we have to work with the electrons we have and the way this works is somebody next to him is going to have to share some more electrons and we're going to start forming this is kind of what happens when you get into this situation this is when you have double and triple bonds and so in this case hydrogen doesn't have anything to share but the nitrogen has these lone pairs that he could instead of leaving them unshared he could share them and so we're going to take one of those there and turn it into a shared pair of electrons and in this case now the carbon says well now it's better but it's still not perfect i only got two four six electrons around me so nitrogen says okay i'll share another pair since hydrogen's not got anything to share and now the carbon's like much better because now he's got two four six eight electrons around him nitrogen still has two four six eight electrons around him a lot more of them are shared than unshared in the original structure we had drawn so uh but we're out of electrons everybody's now going to fill the octet or is full as full as they want to be and this is the structure of hcn so when you run out of electrons and your central atom's not happy that's the evidence you're going to start adding additional bonds double or triple bonds to that central atom from atoms that had lone pairs to share okay so moving on to co2 here let's just start with those valence electrons so carbon's got four each of the two oxygens has six each 6 times 2 is 12 plus 4 is 16 valence electrons carbon's less electronegative so we'll definitely put him in the middle he can make more bonds and then we'll fill up the octets of the outside atoms so 2 4 6 8 10 12. and then 14 16. so at this point once we fill up the outside atoms we've used all 16 of our valence electrons we're out and the question we have to ask ourselves when we're out of electrons is is the central atom happy i.e is that central atom full and in this case he is not so he's only got a total of four electrons he'd like to have eight we need somebody next to him to share and in this case it might not be intuitively obvious here but there are three ways the sharing could happen here so we could have so to get him two more pairs of electrons we could have the oxygen on the left do all the sharing so we'll take away his two lone pairs and put a triple bond there and so for those of you that uh uh might be a little ocd and be like chad it's not symmetrical don't do that so again we'll find out that symmetry is not the governing principle here it'll turns out it'll work in this example but it is not the governing principle though so we'll go back to this structure yet again because the other option would have been to have the auction on the right do all the sharing and put the triple bond over there well finally the other option would have been to have each of the oxygens sharing one pair each and so we could take and erase a pair here put in a double bond so erase a pair here and put a double bond on the other side now when you can draw multiple structures that get every atom a satisfied octet rule and stuff like this we refer to this as resonance now normally we don't actually think of co2 as exhibiting much resonance and in truth it really doesn't exhibit much at all but if you can draw multiple structures that are not the same you do have a case that's called resonance and if you've got resonance it turns out you've got what we referred to as d localized electrons electrons that are in more than one location at the same time and so if i asked you you know what kind of bond is between carbon and oxygen so well it turns out when you've got these resonant structures it's some average of all these structures but some of the structures contribute more than others and it turns out to find out how well they contribute we have to worry about something called formal charge so informal charge is different um and then oxidation state that we learned earlier in the semester so uh and the way this works uh there's a couple different ways you can learn the real formula and use it but you're not likely to remember that six months from now or you can kind of use a different formula but the real formula says take the number of valence electrons for an atom and subtract half of its bonding electrons and all of its non-bonding electrons so we apply this to say oxygen over here in this structure oxygen's normal number of valence electrons is six based on where he is on the pivot table so we'll start with that six and then we'll subtract half of his bonding electrons well he's got one bond that's two bonding electrons half of that would be just a one so half of two is one and then also all of his non-bonding electrons and he's got six of those non-bonding electrons and so we'll do six minus a total of one plus six which is going to get us to negative 1. so again it's the normal number of valence electrons minus half the bonding electrons and the non-bonding electrons so the sum of all those cool the way i like to look at this instead is i just take the normal number valence electron six minus the number of dots one two three four five six end lines instead of looking at as two bonding electrons i'm now not looking at these as electrons at all just dots and lines and i have six dots and one line six dots in one line and so the normal valence of six minus the number of dots and lines is going to get me that same negative one formal charge and so oftentimes when you've got a formal charge like this you'll write it up right next to the atom like so if we do the same thing for carbon in the middle so carbon's normal number of valence electrons is four and carbon in this case has got four lines and no dots and so four minus four one two three four is zero and he's got no formal charge so we don't write anything in just no formal charge and then finally for the other option over here again the normal valence number valence electrons for action is six six minus two dots and three lines is gonna be plus one so six minus one two three four five and now we got plus one and when it's just plus one or minus one sometimes i'll just write plus and minus and not actually say plus one or minus one but if you've got a plus 2 or minus 2 you actually have to write the 2 in and stuff like that so but you could have also seen this simply written plus and minus oftentimes circle doesn't have to be circled cool and that's one of the resonant structures had we looked at it for this one we'd have found out it had been the auction on the left that was minus and the one on the right was plus and the carbon in the middle still no formal charge but what we'll find on this last structure is that there are no formal charges they're all zeros all the way across so here for options six minus one two three four five six so four dots two lines six minus six is zero same thing for the auction on the right and for the carbon in the middle four minus one two three four lines is zero and there's no formal charges and so formal charges there are some rules here but when you've got different resonance structures for drawing out the lewis dot structure of a molecule the best one has the best formal charges and by best formal charges i put it on your hand out there we want fewer formal charges well this molecule has two different atoms with formal charges this and i should molecule this resonant structure has two different atoms with formal charges whereas this one doesn't have any atoms with formal charges that's what makes him the best structure so in fact i think it's worth your time to memorize the lewis structure for co2 it shows up so commonly you should still know how to drive this but it shows up so commonly that if you've got it memorized it'll save you some time on your exam because it's very likely to show up in one way shape or form on your exam cool but that's the best resonance structure so if you were asked to draw the lewis structure for co2 you really shouldn't draw these they're going to ask you to draw this one now we'll find out in other cases that when you've got multiple structures you can draw and none is better than the others something special is going to be implied about that and i'll worry about that a little bit later but we'll talk about resonance and delocalize electrons a little more when we get to that point all right so i told you this is going to get progressively harder and harder and if i wanted to ask you a question on exam that i thought was sure to kind of divide the a's from the b's uh because it's tricky and a lot of students get it wrong i might ask you something like draw the lewis structure for n2o and it's tricky for two different reasons as we'll see so but definitely you just want to follow the rules and as long as you follow the rules in order correctly you're still going to get the right answer there's just some key places where students really want to not follow the rules so but we'll start with the number of valence electrons the two nitrogens each have 5 valence electrons the oxygen's got 6. so 2 times 5 is 10 plus 6 is 16 valence electrons so and a lot of students here's where the first error they're like well there's two nitrogens there's only one oxygen so we'll put that oxygen in the middle well if you recall we said that the less electronegative atom goes in the middle and nitrogen is less electronegative than oxygen we also said that usually it's the one that's written first in a binary compound that goes in the middle and so it turns out one of the nitrogens goes in the middle not the oxygen so there's the first common mistake students make is they don't follow the rules and the rules say put the less electronegative element in the middle again not hydrogen or again that's the one that can make the most bonds and it's usually the first one listed in the compound but again it's just so tempting with only one oxygen two nitrogens to put the oxy in the middle so because that's the pattern that's kind of been happening and so the pattern we're gonna take back up right over here but it's not the case here uh from here we will then fill up the octets of the outside atoms and again you'll always have enough to fill them up in this case we've used 2 4 6 8 10 12 14 16 we've used all 16 of our electrons we don't have any left and the moment you run out of electrons you have to ask yourself is is the central atom happy and is our central atom full not at all in this case he's only got four electrons just like we saw with co2 in similar fashion and so we need him to get four more electrons we're going to need to make two additional bonds and so just like last time there's really three ways we can pull this off we can have the nitrogen on the left do all the sharing so we'll have him form a triple bond we could have had the oxygen on the right doing all the sharing as well so we'll erase two of his lone pairs and make two additional bonds a total of a triple bond to the oxygen or once again we could have had well let's try and get that right that's the nitrogen we could have had both outer atoms do some sharing which just seems fair and equitable right so we could have the nitrogen share one and form a double bond we could have the oxygen share a pair and also form a double bond and so the question is which of these is the best and again if you're all about symmetry and ocd like i tend to be then this is probably the one you want to choose and unfortunately it is not the best resonance structure and i say resonance and again if you can draw multiple structures wherever it is getting this filled octet you've got resonance going on so but this is again going to be an example where there is one best resonance structure based again on formal charges again when you've got multiple resonance structures it's formal charge that will help you distinguish between them all right so if we look at this first one so nitrogen's got five valence electrons and five minus two dots and three lines five minus one two three four five is zero for the central nitrogen five minus one two three four lines is plus one and for auction his normal valence is six six minus one line and six dots so six minus one two three four five six seven is negative one okay so we got formal charges in the next one for nitrogen here so we've got five is as normal valence and so five minus one two three four five six dots in one line so five minus seven is negative two and that's horrible it turns out we'll keep going though uh for the central nitrogen five minus one two three four lines is plus one and then for the oxygen six minus one two three lines at four five two dots so six minus five also plus one now one thing you should realize that the formal charges are always going to add up to the overall charge on your compound or ion now these there's a neutral compound notice there's no charge listed right over here and so the formal charges have to add up to zero just like they do here and just like they do here now we talked about some rules for getting the best formal charges so and the first rule is you want the fewest atoms with formal charges well here i got two atoms of formal charges here i got three this one's going to be better we still got to worry about this last one so we can already rule this one out it's not it's not as good now the second rule deals with having formal charges closer to zero like negative one is better than negative two so this is going to get ruled out again by that rule and then we'll find out if you still haven't distinguished between your resonant structures then if you've got a difference of where the which atom gets the negative charge always put it on the more electronegative atom to get to the major resonance contributor we call it the best resonance structure so but so far this one's out the top one's winning so far and if we go assign formal charges to this bottom one now that again symmetrically just looks so good uh nitrogen's only got five valence electrons minus one two three four five six five minus six is negative one for the central nitrogen here five again minus one two three four is plus one and then oxygen's normal number of valence electrons is six six minus one two three four five six is zero and he's got no formal charge and so now we see these two and the first rule for distinguishing resonance structures based on formal charges the fewer formal charges well they both have two atoms with formal charges each no difference so then you want to say uh you want formal charges that are closer to zero well we have plus one and minus one plus one and minus one and there's still no difference and this is what makes this last one doubly tricky is now it's distinguishing between these resonant structures again using formal charge not using symmetry it's using formal charge and the last rule says that if the first two rules didn't work then it's going to come down to which atom gets the negative charge and in this case the the first one the negative charge is on the oxygen whereas in the last one the negative charge is on the nitrogen on the left notice in both structures that central nitrogen gets a positive formal charge there's no difference in that regard it's all about where the negative formal charge goes and do we want it on the oxygen or do we want it on the nitrogen well again the best place to put a negative formal charge is on the more electronegative atom and in this case oxygen's more electronegative than nitrogen and so it turns out this first one is actually your major resonance contributor and so when you've got these resonant structures it turns out that none of these structures is perfectly correct usually the way this works what the molecule would really look like would be some average of the structures but it would look more like this one than it would look like either one of these and it would probably look more like this one than it would like this one because it's a little better than this one and stuff but what that means then in this case is that you're probably going to have a significant amount of negative charge on this oxygen since this is the best resonance structure notice in the in this resonance structure there's no charge and this one's positive charge but in the best one it's a negative charge and so there's probably going to be a fair amount of negative on that oxygen so and in all the resonance structures there's a positive charge on the central nitrogen so there's just going to be a positive charge on that nitrogen and then finally in in the the nitrogen on the left it's either neutral or it's negative 2 or it's negative 1. well in the best structure it's neutral so it's probably going to have a very small amount of partial negative due to the contributions of these guys but it's going to be closer to neutral than it is going to be either negative 1 or negative 2. so that's the way this works uh in terms of uh delocalization here and so delocalization means the electrons can be in multiple locations at the same time and so if you notice the extra bonds that we have well they're mostly between the two nitrogens but they could be a little bit between the nitro and the oxygen according to these two structures as well and so it turns out they're partially in both the locations at the same time and that's really frustrating but again in this example again it mostly the the resonance hybrid we say is going to mostly resemble this particular structure rather than the other two we'll get to an example with equivalent resonance structures in a little bit and again we'll have to do something a little special for that okay next one on the list here is sf4 and we're going to find out this is our example our first example of an expanded octet and there's nothing you can look at this that's just going to be like it's an expanded octet now it turns out the least electronegative atom is the sulfur and sulfur is in the third row of the periodic table or lower so it's allowed to go over the octet rule but that doesn't mean that it's going to have to go over the octet rule well we'll find out that when you've got one of these atoms uh in the third row or lower the only chance it has of going over the octet rules if it's the central atom but again there's no guarantee there either even when it's a central atom it doesn't have to go over the octet rule but if you follow the rules it will naturally work out to either go over the octet rule or not as we'll see all right so sulfur's got six valence electrons being right below oxygen and then each of the fluorines has seven each so four times seven is 28 plus six is 34 valence electrons and again sulfur's less electronegative so we'll put him in the middle and surround him with four fluorines cool next thing you do is fill up the outside atoms which are more electronegative and so and again you'll always have enough electrons to just fill them up so don't even have to count you just start filling this up and now we have to take an accounting though any electrons left over at this point would always go on the central atom and in this case we've used up two four six eight eight around this one eight around each of the four fluorines if you look at it so 8 16 24 32 we've got two electrons left and if you have filled up the outside atoms and you still have electrons left over they always go on the central atom and in pairs as well and so in this case we got two electrons left it has to go on the central atom that's where they go that's what the rules say and they go on as a parent so you have to actually put them on one of the diagonals here and you see that this means that sulfur now has not just 8 electrons but 2 4 6 8 10 electrons around him and we're out of electrons we've we've used all 34 and usually this would be the point where we'd say is the central atom happy well in this case sulfur is more than happy he is drunk with electrons he's not just full he's more than full he's got ten and that's okay because he's in the third row of the periodic table or lower he's allowed to exceed that octet rule and it doesn't have to but if you follow the rules and he does no problem with that now one thing you should do when you've got an expanded octet like this is you should just check your formal charges real quick and most of the time you're probably not going to end up with any formal charges but what you'll find out is that if you've got some of the outside atoms being negative and the central atom being positive you could actually probably reduce your formal charges by forming a double bond and possibly more than one from the outside atoms but again it's only gonna be the case if you have a negative formal charge atom on the outside bonded to the central atom with a positive formal charge but in this case if we look at those fluorines uh normal valence is seven minus one two three four five six seven is no formal charge on any of the fluorines and then for sulfur his normal valence is six and so six minus one two three four five six is zero as well and so no formal charges there's no better structure than this there's no other structure than this it turns out as well this is the correct lewis structure for sf4 our first example of an expanded octet all right next one on the list is special a little unusual it has a noble gas in it xenon in this case and recall the noble gases are known for being chemically inert they hardly ever do anything chemically and the only chance we talked about we briefly mentioned that uh they can when they're bonded to something very electronegative for some of the larger noble gases and fluorine is electronegative as they come and so this is one of the few compounds uh that we can form using one of these noble gases well if you've got a noble gas in your compound i'll tell you right off the bat it's going to be the central atom and so in this case xenon is going to go in the middle and be bonded to four fluorines so and if you've got a a noble gas that already had a filled octet and now he's going to make bonds then he's going to exceed the octet rule so this is definitely expanded octet if you've got a noble gas involved in any molecule it will be an expanded octet uh if we count up our electrons xenon's got eight each of the four fluorines has seven four times seven is 28 plus eight is 36 and we'll do same rules though we're just going to follow the rules we set up our skeleton and we are just going to fill up the outside atoms next and again you'll always have enough to just fill them up so just do it indiscriminately cool and at this point we've used 8 16 24 32 we have four electrons left and once you fill up the outside if you have any electrons left over they always go on the central atom in pairs and since we got four left that means we got to add two pairs and you just got to put on two of the diagonals and it's customary to put them on opposite diagonals as if they're like repelling each other and we'll find out in the next chapter when we talk about molecular geometry how that kind of is an ac a little more accurate portrayal of reality than putting them right next to each other but technically there's nothing wrong with putting them right next to each other so but we'll find out that when we do actually try to draw the three-dimensional shape that there really will be 180 degrees apart cool at this point once you run out electrons you always ask yourself is is the central atom happy and again xenon's more than happy here two four six eight ten twelve electrons more than a filled octet drunk with electrons super happy and nothing wrong with it because he's definitely in the third row or in this case lower uh and is allowed to have that expanded octet cool we should really quickly assign formal charges just to make sure there's nothing else special with any any expanded octet and in this case uh the fluorines again seven is the normal valence seven minus one two three four five six seven is zero and for xenon he already had a filled octet eight is his normal valence and so eight minus one two three four five six seven eight also zero with no formal charges there's nothing to do to doctor up this structure this is the correct lewis structure for xc f4 all right so so4 2 minus is the next one and this is the first example where we've seen an ion and when you've got an ion if it's an anion you have to add extra electrons if it's negative one add one extra electron if it's negative two add two extra electrons and so on and so forth if it's a cation you actually remove electrons plus one you remove one electron plus two you lose two electrons so on and so forth so in this case when we go to add up the electrons sulfur's got six valence electrons each of the four oxygens has six each as well and four times six is 24 plus another six for sulfur is 30 and since it's negative two we have to add two additional electrons for a total of 32. cool so that's how we account for that negative charge and we're going to put the less electronegative atom in the middle which is sulfur in this case notice oxygen is the second most electronegative on the periodic table after fluorine so it's definitely sulfur that's less electronegative and first thing we'll do is fill up those outside atoms all right so once they are full we'll see if we have any electrons left and in this case again we've got eight around this auction eight around this auction eight around this auction eight around this oxygen and that's eight sixteen twenty 16 24 32 we have used all 32 electrons so the question we got to ask ourselves then is once you're out of electrons is the central atom happy and in this case this sulfur has got 2 4 6 8 he's got a filled octet he is happy and again we do have sulfur in the center though and notice sulfur how many bonds would you normally expect to see sulfur making well normally he's just two electrons short of a field octet we normally expected to see him making just two bonds well he's making four so and often times when you've got an element that's third row the periodic table or lower and it's not making its normal number of bonds it also means that an expanded octet might be possible or even probable in your structure and so in this case we're actually going to go through and assign our formal charges as well for the options on the outside again normal valence number is 6 valence electrons so 6 minus 1 2 3 4 5 6 7. so six dots in one line so six minus seven is negative one and they're all going to come out with that same calculation and then for the sulfur in the middle his normal number of valence electrons is also six so six minus one two three four lines so six minus four would be plus two and so not the greatest structure in the world truth be told we got five atoms with formal charges and we even have a plus two which is so far from zero and stuff like this uh and if you recall i also said if you've got an atom who can have an expanded octet in the middle and he's got a positive formal charge of some sort and the ones next to him on the outside have a negative formal charge well you can actually reduce your formal charges by having the negatives donate and share electrons with the positive and so that's what we're going to do here we're going to take say the auction on the right here we're going to grace one of his lone pairs and have them be a shared pair with sulfur and you'll see what we accomplish with this so in this case if we look at this oxygen now now they shared some electrons uh his normal valence number is six again and six minus one two three four five six is now zero he no longer has a formal charge and for the sulfur here his normal valence number is also six and six minus one two three four five is now only plus one and so by having one of those negative atoms on the outside share with the positive central atom we reduced our formal charges one of the auctions went from negative to now having no formal charge and the sulfur went from plus two to plus one both of which are improvements on the structure so since he's still plus one though we could do this again to get him to neutral and so in this case any one of the other three we can choose and it's customary to kind of choose the opposite one but there's nothing incorrect about choosing a different one it turns out but we will share again and that's going to make the auction on the left now have no formal charge and it'll make the sulfur now also have no formal charge again sulfur's normal valence is six six minus one two three four five six is zero and now we've reduced the formal charges now in this case if we had either one of these auctions do any more sharing it would actually start giving the sulfur a negative formal charge and we don't want to go that far so you only want to kind of take this approach when you can reduce your formal charges so in this case if i uh let's say we do it let's just do it and we'll undo it in a second so if i make and have one of this this auction do some sharing as well what we'll find out is that this auction now would no longer have a formal charge but the sulfur in the middle would have a formal charge and this is not a better structure because in the previous structure i had the negative formal charge on the more electronegative oxygen now i have the negative formal charge on the less electronegative sulfur and i'd much rather have it on the more electronegative oxygen so we were done it turns out with this as the best possible resonance contributor here and again two auctions have negative formal charges but the other three items have no formal charges uh when you've got an ion by the way it's also customary to put the entire lewis structure in brackets and then to put its charge in the upper right hand side and so for an ion here we'll put it in brackets and put that charge we'll find out there's one other case where we might use these brackets as well in the next example one other thing to note sometimes this happens when we could have an expanded octet or not and if you recall we had before we got this far we had negative formal charges on all the oxygens and a plus two charge on the sulfur and again this was this superior structure it had fewer former charges yeah fewer formal charges formal charges that are closer to zero uh the whole bit so however sometimes a question might be asked in such a way for something where this expanded octet is possible for that central atom that might just say draw the best structure for so4 2 minus where every atom satisfies the octet rule well notice again our central atom here is totally exceeding the octet rule but back in this structure every atom is satisfying the octet rule it's not the best one based on formal charge which is typically the the most important criteria but if your instructions were to draw the best lewis structure where everybody has a filled octet well then that would be the one and again it wouldn't be the best overall structure but it would be the one where every atom has a filled octet all right so last one we're going to take a look at here is no3 minus and no3 minus first thing we should do is figure out how many valence electrons we've got nitrogen's got five each of the oxygen has six each so it's 3 times 6 is 18 plus 5 is 23 and because we have a negative one charge we add one extra to get to 24. and in this case nitrogen's less electronegative and goes in the middle so we've got our skeleton set up and then we'll fill up the outside atoms and once we got them full we now have to take an accounting when we've used 8 16 24 electrons we only had 24 electrons and so we're out and once we're out we have to ask ourselves is the central atom happy is the central atom full in this case he's not he's just got two four six electrons around him he needs one more pair so he needs somebody next to him who has a lone pair to share it with him and form a double bond well how many options do we have well we have three options and in this case it's not like one's a nitrogen one's an oxygen like we saw with n2o or something like that they're all oxygens they're all equivalent and as a result if i just need any one of them to share and form a double bond well then any one of them could and we actually have three equivalent structures to draw so i could have it be the oxygen on the left and we'd draw that structure we could have had it be the oxygen in the middle and in this case i erased a lone pair essentially effectively that was here made it a double bond or it could have been the oxygen on the right that did the extra sharing and that leaves us with these three structures and again when you've got multiple structures you can draw we call them resonance structures and usually we'll use formal charge to distinguish between them well i'm going to take it for granted that you guys know how to uh go about and do formal charges now so i'm just write some in i want to get lost uh in the trees here so but it turns out for this option six minus one two three four five six seven is negative one same for this one this one's neutral and then for nitrogen five minus one two three four is plus one and what you're going to find is that for all three structures the nitrogen's got a positive formal charge and then two of the oxygens are going to have a negative formal charge and if we look at this then formal charge is not going to help us distinguish between these three resonant structures they all have three atoms with formal charges and they're all uh with a plus one formal charge on nitrogen and two of the oxygens having minus one formal charges they're equivalent in every respect and so instead of having like one major resonance contributor which would be the you know if you're asked to draw a lewis structure if you've got resonance but there's one major resonance contributor that's typically what they want you to draw but in this case they're all equivalent and so you're not going to be able to get away with just drawing the best one in this case they actually want you to draw all of them as a proper lewis structure and so what they want you to do when you've got equivalent resonance structures like this is draw them all out and then put an arrow a double-headed arrow like this notice don't do this that is got a different meaning in chemistry so it's got to be a double-headed arrow like that guy and put it in between each of the structures and then put them all in brackets now in this case we would have put these in brackets anyways because it's an ion with a charge so it turns out there's two reasons to use brackets either because you're doing the lewis structure for an ion that has a charge or if you've got multiple equivalent resonance structures like we do here well in this case either reason would have led us to use brackets here and we still need to put that negative one charge right here you need to write negative or negative one either one is correct okay so in this case we have resonance and this is where things get a little bit weird so and you need to understand what resonance means and again resonance always implies delocalized electrons and i mentioned that earlier but i didn't explain it because it's going to make the most sense in this example now we say it means delocalized electrons and here's the deal none not one of these three structures is an accurate depiction of reality none of them really exist the real molecule it turns out as an average some sort of average of all these structures at the same time but we try to confuse you by putting these double-headed arrows and using the word resonance because if we use that word resonance then you might also try to use the word resonating and then we put this double arrow and it makes you think well maybe it's resonating back and forth maybe if i took a photo of this molecule with an electron microscope or something at one moment in time maybe it looks like this and so you had a double bond over here and two single bonds over here and then maybe if you take you know a photograph of it again a couple seconds later now the double bonds down over here and these are the single bonds and then if you took a another photograph you know a second later now the double one's on the right and these are single bonds and one thing you should realize that a double bond is stronger than a single bond and a triple bond is stronger than a double bond when it's between the same two atoms and when you have a stronger bond between the same two atoms it actually ends up being a shorter bond they're pulled closer together and so by saying we have a double bond here and two single bonds here we're actually saying that we have one short bond and two longer bonds so again one short bond and two longer bonds one short bond and two longer bonds so but unfortunately when we actually look at the structure of no3 minus here we actually find that all three bonds are exactly the same length all the time it's not like you look at it and which one's the short one right now no it turns out they're all the same length all the time and that is super frustrating so it turns out that none of these structures actually give an accurate depiction of what the molecule looks like at any moment ever what the molecule looks like all the time though is the average of all these structures and so we draw what's called the resonance hybrid and so it turns out that you know here we've got a double bond but here it's a single and here it's a single what we end up having is an extra partial bond so it's not a double bond but it's more than a single bond it's somewhere in between a single and a double bond and it turns out it's if we actually measure its length and stuff like that it's a little bit shorter than a single bond but a little bit longer than a double bond and so we find out it's intermediate in both strength and length between a single and a double bond and same thing for this one right here it's a double in this structure but it's a single in the other two and so it's a partial double bond right there as well and then same thing for the one on right single here single here but double here and so it's a partial and so it turns out that all three bonds are exactly the same length exactly the same strength and they're stronger than a single bond but weaker than a double bond and if you actually look at any one structure you can say well between all three options you have a total of one two three four bonds spread across three locations and four bonds over three locations is four thirds of a bond and so these are really like one and one third bonds and so they're like 33 percent stronger a third stronger than a single bond essentially is how that works and therefore a third shorter on average-ish than a single bond as well cool so you should understand that when we've got these equivalent resonance structures a proper way to draw the lewis structure out is draw all of them in brackets with these lovely double-headed arrows between the structures you should know that none of these structures is correct that what we really have again is this resonance hybrid i'll draw that out so you might get a multiple choice question that just shows which of these you know represents the resonance hybrid or something like that so you might also have to figure out you know how strong is that bond like it's a four third bond a one and one third bond in this example and again we just figured that out by taking you know our structures and saying there's one two three four bonds shared across three locations evenly okay you could have also looked at it this way and said double here single here single here and if you take you know the average of double single and single well you're like double single and single is four and then divide by three structures you still get four thirds of a bond you can get it that way as well technically all right notice also in this resonant hybrid here i have refrained from drawing charges but what we can do is say that that the nitrogen in the middle is fully positively charged so but it turns out the oxygens in this structure this auction has no formal charge in these structures it's minus one what's the average of zero minus one and minus one well we'd say minus one plus minus one plus zero is minus two and then that's averaged across three structures so it's minus two-thirds well we don't usually draw the fractions in so what we usually do is use this symbol to mean partial and it's got a partial negative charge and they all essentially have this partial negative charge which we just figured out was a negative two-thirds of a charge and we've seen this once before but this is uh the lowercase greek letter delta and we use it for like partial derivatives like d y d x uh uh but instead of the d with partial derivatives we use the delta instead but it just means partial and so partially negative on all three oxygens and that is our resonance hybrid this is what the molecule looks like all the time again if you're asked to draw a proper lewis structure you draw all three of these if you're asked what does the resonance hybrid look like well it looks like this and so this extra partial bond it's one extra bond shared in all three locations it is simultaneously in all three locations and that's what we mean by delocalized electrons so again delocalized just means in multiple locations at the same time we don't have you know anything on the macro world that really look works like this but for these electrons they can be in three different bonding locations at the same time so and that's what that dashed line represents cool we have pretty much covered just about every nuance of drawing lewis structures you could possibly see and i recommend you get some good practice at this and you get fast at it before your exam it's going to be super important usually this is taught in the same unit as molecular shapes which will come in the next chapter and unfortunately to be able to properly assign molecular shapes you have to be able to properly draw lewis structures first so this is going to be important for itself but also important because it's going to be foundational for the next chapter as well now if you found this lesson helpful would you consider hitting that like button best thing you can do to make sure youtube shares this lesson with other students as well and if you're looking for practice on lewis dot structures or anything general chemistry related check out my general chemistry master course i'll leave a link in the description free trial is available happy studying

Channel: Chad's Prep

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Keywords: lewis structure, lewis structures, lewis dot structure, how to draw lewis structures, lewis structure practice problems, lewis dot structure tutorial, drawing lewis structures, lewis diagrams, lewis dot structure practice, lewis dot diagrams, lewis structure chemistry, lewis structure formal charge, lewis structure for co2, lewis structure of hcn, lewis structure of no3-, expanded octet lewis structure, lewis structure for no3-, drawing lewis structures practice

Id: UrY5LTNdDdw

Length: 67min 5sec (4025 seconds)

Published: Wed Oct 27 2021