lewis structures that's going to be covered in this very first lesson in the very first chapter of my new organic chemistry playlist now in this first chapter it's pretty much going to be a review of general chemistry we'll start off with lewis structures move on to formal charges uh hybridization and valence bond theory molecular orbital theory which i call it again a review from gen camera that's often a place where we skimp a little bit in gen chem so might be a little less review in that particular section uh and then we'll finish this chapter off with a discussion of polarity and intermolecular forces now because this is the first lesson in the entire playlist let me give you just a little bit of an introduction to the organic chemistry playlist here so this is going to be very similar to my old playlist but i'm going to be using a whiteboard instead of powerpoint and a green screen so a lot of you have said you much prefer the whiteboard instructions so ask and ye shall receive uh so i'll be releasing these weekly throughout the 2020 2021 school year so if you don't want to miss one of these uh subscribe to the channel and click the bell notifications and you'll be notified every time i put one of these up now because this is the first lesson in the first chapter of the entire course i do want to give you just a little bit of an introduction to organic chemistry so organic chemistry uh as contrasted to inorganic chemistry so we look at organic chemistry it's just the study of molecules you find in living systems now they used to think that there was something special called the vital force and things are sort of i don't really care about that but the big thing you know that this is going to be the study of carbon containing compounds that's the way we're going to look at organic chemistry we're not going to talk about like it's the study of growing things without pesticides and natural growing you know whatever it's all about carbon so whereas inorganic chemistry is going to deal with all the rest of the elements here we're going to deal with primarily things involving carbon so this is going to be a beautiful thing we're because we're pretty much not going to talk much about transition metals and this is going to actually simplify and focus on what we really need to study uh in some of this review of gen chem here now one other quick note about organic chemistry so it is not taught exactly the same from textbook to textbook professor professor university to university so i'm going to be shooting for kind of middle of the road organic chemistry course so and where i know there's some variability in what's presented i'm going to let you guys know and i'm like hey some of you guys might see this this way and some of you might see it this way so and things of a sort it's not that one's right and one's wrong so but it turns out that organic chemistry and life in general reality is very complex and so sometimes instead of giving you the exact you know level of detail that we know and understandable organic chemistry we often just give you some trends and then kind of try to make some absolute rules out of them and then don't tell you that that's what we're doing so when i know that something's often commonly presented in more than one way i will make particular note of that and i'll let you know and i'll say hey by the way pay particular attention to how your professor presents this topic so this is not going to happen super often but it will happen in a few key places throughout the course all right so let's talk about lewis structures now in a lewis structure we put valence electrons on there so you got to know how to remember you know how you identify how many valence electrons an element has based on where it is on the periodic table so in the group one elements hydrogen on down they've got one valence electron group two elements beryllium on down have two valence electrons borons column three valence electrons carbon of particular node four valence electrons nitrogen five valence electron oxygen six valence electrons uh the halogens seven valence electrons and then the noble gas is eight valence electrons which we really don't care about them they're not going to be involved in any bonding that we'll see in this chapter but uh that's kind of how you identify that and also of note is you can usually figure out based on that how many bonds an atom is typically going to make now it's not absolute sometimes you'll see carbon with a different number of bonds than its normal number of four so but we can predict that carbon can make for four bonds for for example from this info so the way this works is you look at how many electrons is an atom short from getting a filled octet that's usually its typical number of bonds so like the halogens here they have seven valence electrons well that's one short of a filled octet and that's why you typically see them making just one bond so oxygen's column on the other hand so they've got six valence electrons and six is too short of a field octet so they're typically going to be involved in making two bonds so nitrogen's column typically gonna make three bonds carbons call them typically four bonds and then we'll go back to boron here and you would say five bonds except that would be a violation of the octet rule and you find out it drops back down to three bonds so but let's talk about that octet rule so the octet rule talks about you know most atoms are going uh struggling in their existence to have eight valence electrons and we should talk about where that comes from first and so in any typical subshell or i should say really say shell there's going to be an s subshell and a p subshell for where electrons can live and in the s subshell there's just one orbital and so you can only put two electrons there so but in the p sub shell there's three orbitals and so you can fit a grand total of six electrons there so two per any orbital in the universe here so we'll put one two there and then one two three four five six for a grand total of eight so this is where the octet rule comes from it's just how many electrons can you fit in a typical shell is the way this works so for like carbon he's shooting for eight things of his sort now boron here we'll find out is an exception so he goes under the octet rule and then we'll find out some of these bigger guys can go over the octet rule and so you really need to understand these exceptions as well so first we'll start with what's called an expanded octet now the truth is this is not the most important one for organic chemistry because for organic chemistry we're really going to be focusing on these guys and it turns out they're not allowed to go over the octet rule so what is special about say sulfur that lets him go over the octet rule well first off it's not just sulfur it's anybody in this third period of the periodic table and lower and what's special about that is starting in the third shell in addition to the s and p subshells you're also going to have a d subshell so and the d sub shell gets lets you put more electrons in there up to 10 more if you want to so here's the deal though so that only you only start having a d sub shell in the third shell or above so a third shell and above corresponds to the elements in the third period and below on the periodic table that's why they're allowed to go over the octet rule now they don't have to they're just allowed to cool but we're not going to be dealing a lot with these because most of the chemistry we're going to study is going to involve like carbon nitrogen oxygen and then hydrogen quite a bit so now you might see this in an inorganic reagent like sulfuric acid we might see it in him in fact we'll stick him up on the board real quick sulfuric acid so i won't draw in all the lone pairs and everything but you can see sulfur ends up with 2 4 6 8 10 12 electrons around him way over the octet rule now again sulfur is not always going to violate the octave he doesn't have to violate the octet rule but he does have the option of going over the octet and again we call that an expanded octet now your second example and this one we'll come across a little more frequently is going under the octet rule so and hydrogen definitely goes under the octet rule so oftentimes we say that when an atom satisfies octet it looks at least from its electron configuration standpoint like a noble gas well hydrogen is just trying to look like helium and in the first shell there's no p or d there's just the s and so if you're in the first shell like hydrogen you only want two electrons not eight and so oftentimes you know we'll see hydrogen always actually going under the octet rule so but turns out beryllium also now beryllium is more of a metal than a non-metal so you're not going to see it making covalent bonds too often but maybe you'll come across this in this first chapter but beryllium only has two valence electrons and he's like i'll share one with you and i'll share one with you and so with beryllium we often see beryllium only making two bonds like we do here in beh2 and as a result he only ends up with four electrons around him total so it goes under the octet rule and that's pretty typical for beryllium so we'll see the same thing actually with boron and aluminum so if we look at boron here in bh3 at least in this particular version of drawing the lewis structure bh3 we see that boron with three valence electrons can only share with three different atoms to make three bonds and ends up with only six now there are cases where boron will end up with four bonds and end up with a filled octet but what i'm saying is that it's normal it's typical for him to only end up with six not eight and therefore he goes under the octet rule and again both boron and aluminum often only get six and go under the octet rule cool so third row and lower can go over but we see in hydrogen beryllium boron aluminum they all have an option of going under and it's pretty typical for them now the last exception to the octet rule is pretty much not going to be relevant to organic chemistry if you've got a species like no which overall has an odd number of valence electrons then somebody's going to have to have an odd number which is not 8 as a result but we're really just not going to encounter this in organic chemistry so i've put it you know on your hand out there and i put it on the list and i'll put it up on the screen here but the truth is you're probably got about a 0.001 chance of ever encountering it in organic chemistry all right so there's our exceptions to the octet rule so now we need to work on like you know how do we accomplish getting a field octet well there's one of a couple different ways and so one of them involves having a metal and a non-metal and the way we draw lewis structures for elements as you just look at how many valence i've got and sodium's got one and so we'll put one electron here for sodium and then for chlorine we've got seven and so you put them on the four sides you might recall so one two three four and then start pairing them up and so there's our seven and in this case when you have a metal and a non-metal well it turns out non-metals love to gain electrons and metals i won't say they love to lose electrons but if somebody's got to lose an electron the metals have rather low ionization energies and they're the easiest ones to remove one from so here chlorine says hey sodium give me your electron and sodium says okay and that electron is going to get transferred over and as a result we're going to end up with a sodium no longer having that valence electron and a chlorine now having eight valence electrons and it turns out with chlorine now having eight he's got what's called a negative formal charge we'll deal with that a little more specifically of how you calculate that here later and sodium's gonna have a positive charge and they now would potentially be ionically bonded together now we're not going to deal with ionic bonds very often in this course they will come up about time and again but most of them will be dealing with covalent bonds so so with an ionic bond here we get the complete transfer of electrons and the result is that now chlorine's got a filled octet and you might be like sodium's got nothing well sort of but sodium actually still has a complete second shell that's full and in that sense he's still got a filled octet as well and so now by a transfer electrons when you have a metal and a non-metal that's how you achieve your filled octet now what if you've got two nonmetals so say i've got two chlorine atoms so one two three four five six seven and one two three four five six seven cool and this chlorine says hey glory give me your electron and this chlorine says no give me your electron and the problem is now we only have two elements that both want to gain electrons but neither one so is going to lose one very easily they have rather high ionization energies so in this case there's no transfer of electrons it's not possible so we have nobody who's willing to lose one but they can share and so they're going to put their electrons together and share them and the results gonna look like so cool and there's our two shared electrons and much more commonly you might recall that we represent those shared electrons as a line to represent that covalent bond and so those two shared electrons are a covalent bond and now the chlorine on the left gets to say hey i've got 2 4 6 8 around me and the chlorine on the right says i've got 2 4 6 8 around me they both get to count those shared electrons as part of their octet and so sharing now accomplishes getting a field octet in this case so if you've got two non-metals and for organic chemistry we're gonna almost exclusively be dealing with non-metals most of the time that's what we're gonna be dealing much more commonly with sharing electrons rather than a transfer of electrons like we saw in the first example so let's draw some lewis structures so we're going to draw five lewis structures here the ones we got outlined here in this case we're going to start easy and then we're going to work our way to make it a little harder and then we're going to see it finish this off with a couple examples of very common organic molecules you might encounter both in this chapter as well as later on in this course now we'll start with ch4 we're going to start off easy here and first thing you want to do is just figure out how many valence electrons you got to work with and so again carbon's got four and each of the hydrogen's got one each for four more for a grand total of eight valence electrons second thing you want to do is figure out who's your central atom in some cases you might have more than one but who's your central atom and then set up a skeleton for your molecule with connecting all the atoms by single bonds and so in this case your central atom is typically the one that can make the most bonds and you might recall that the halogens being one short of a failed octet typically make one bond oxygen columns two bonds nitrogens call them three bonds carbons call them four bonds borons call them back to three bonds and then hydrogen just a single bond and so in this case the way that usually works then is usually you can make more bonds as you move this way up to a point and so usually we like to say that there's a pattern where typically the least electronegative atom becomes the central atom except just not hydrogen because hydrogen actually is less electronegative than carbon but carbon makes way more bonds so if kind of a general rule it actually works most of the time so the least electronegative atom will be the central atom just not hydrogen all right so in this case we'll put carbon in the center and we'll set up that skeleton so we'll connect carbon to all of the hydrogens with single bonds cool and next thing you'll do is you'll fill up the outside atoms well in this case the hydrogens don't want to fill the octet they just want two in order to look like helium and so in this case they've all got two around them and they're full and so at that point you'll always have enough electrons to fill up the outside atoms so then you move to the central atom and any electrons left over at this point go on that central atom and sometimes you'll have some left over and sometimes you won't so this is your first point you've got to take an accounting in this case we had eight electrons to work with and we've used all two four six eight so there's none left to give the central atom and once you're out of electrons then you ask is the central atom satisfied and usually that means this essential atom have a filled octet for those that are following the octet rule and in this case carbon's following the octet rule and he's got two four six eight and he's got that filled octet and so i like to think he's happy and that makes this your structure for ch4 but this is kind of the general pattern we start off with you know so figure out your valence electrons figure out who goes in the middle and set up your skeleton so fill up the outside atoms first any additional go in the inside atom and then we'll see if we need to take it further we didn't in this example we'll see what to do in the next couple so the next one here is ammonia so nitrogen has got five valence electrons and again each of the hydrogen's got one each so for another three and once again we've got eight electrons now nitrogen being the least electronegative besides hydrogen goes in the middle because it can make the most bonds and so we'll set up our skeleton here and at this point we'll then fill up the outside atoms but these are hydrogens and they're already full and so then we've got to do our accounting again this is the one point where you got to do an accounting and figure out are there any electrons left to give the central atom well in this case we've used two four six we've got eight yeah we've got electrons left over and when you've got electrons left over after filling up the outside atoms you always put them on in pairs on that central atom well we've only got a pair so we're gonna have to just put one pair on that central atom and now that we've used all eight electrons we're out of electrons you ask is the central atom happy or is the central atom satisfied does he have a filled octet in this case two four six eight he's got a filled octet and he is indeed happy and that's your structure for nh3 so let's see another example here so h2co each of the hydrogen's got one valence electron each carbon's got four and oxygen's got six for a grand total of 12 valence electrons now in this case carbon is the least electronegative besides hydrogen and so carbon being able to make the most bonds goes in the center so we'll put carbon in the center surrounded by the others and in this case we say okay fill up the outside atoms first well again the hydrogens are full but oxygen wants a filled octet and he could definitely use some more electrons and then you say are there any electrons left for that central atom well in this case we've used 2 4 6 8 10 12 and there are none left to give carbon and now that we're out of electrons we say is carbon happy or is carbon satisfied or does carbon have a filled octet in this case he doesn't he's only got six around him but we can't put another lone pair there because we don't have another pair of electrons to give and when you get to this point when you're central you run out of electrons your central atom is not happy somebody next to him needs to share and in this case the only atom that's got any electrons to share lone pairs is the oxygen and so the oxygen is going to share a pair we'll erase one of these lone pairs of electrons non-bonding electrons and make them bonding electrons to carbon instead in this case a pi bond and so now we've got a double bond and now carbon's like oh great i feel good now i've got two four six eight electrons around me now everybody's got a filled octet and this is your lewis structure for h2co so these last two are a little more complicated so and these are typical organic molecules and these are what we call condensed formulas so we'll talk about these in chapter two in a little more detail but for now you should realize that they give you a nice way of representing a little more complicated molecules and one thing you should know at this point when we put a carbon followed by hydrogens those hydrogens are going to be bonded to that carbon so in this case when i see ch3 i know that these three h's are bonded to that carbon and suffice to say that's going to help you set up your skeleton so in this case we've got two carbons four valence electrons each that's eight an oxygen that's another six for fourteen plus four hydrogens another four total for a grand total of eighteen valence electrons now with more complicated molecules here we're going to find out that you don't necessarily just have one central atom so but kind again use that condensed formula to kind of piece together what this looks like so we've got this first carbon that's bonded to all three of those hydrogens and then he's also bonded to the next carbon in the chain so and then he's monitored that hydrogen and then he's bonded to that oxygen as well so really both of these can be viewed as being central atoms in this case now we're going to fill up the outside atoms next and we've used 2 4 6 8 10 12 electrons we've still got six left with but you're always gonna have enough to do those outside atoms well the only outside that's not full is oxygen because the hydrogens are all good but that oxygen will give him three lone pairs and now we've used all 18 electrons and we have to ask are all the central atoms happy well there's one on the left he's good two four six eight he's got to fill the octet but this one on the right is not he's only got two four six so in this case how do you give him that filled octet well again once you round electrons if any central atoms are not happy somebody next to them needs to share well the only atom that's got lone pairs to share next to this carbon is the oxygen and so just like on the last example we'll release erase one of these lone pairs and make another bond and now this carbon's got a filled octet two four six eight and this is the proper lewis structure for ch3cho and we call this an aldehyde so which we'll get to also in chapter two not important for now though but a very common type of organic molecule and you'll want to get to the point of recognizing what this kind of structure is going to look like especially the end of this based on seeing this kind of condensed formula so cool we just kind of derived what this lewis structure looks like but by the end of chapter 2 you're going to want to memorize what that lewis structure looks like all right so this guy right here has acetic acid it's the main non-water component in vinegar so it's what we call a carboxylic acid so in this case if we're going to set up our lewis structure first we're going to do is count our valence electrons we've got two carbons that's eight four hydrogens for another four more and that's 12 and then two oxygen six each for another 12 to get us to 24. now we'll start with that first carbon and set up our skeleton he's bonded those three hydrogens next to him and then he's bonded to the next carbon in the chain and here's where things are not going to be so clear-cut so here you might be like it's c-o-o-h well one we know this carbon is not bonded to the hydrogen or else it would have been written right next to that okay but we might be tempted to try and do this so and it turns out there's not a great way to make this work but there's no way for you to know that up front but we'll work it out and then figure out why it doesn't work and then we'll start again so in this case we got a lot of central atoms here and stuff like that and we'll fill up what are the outside atoms which is a little bit nebulous but in this case i'm just going to fill up the more electronegative atoms first but we've used 2 4 6 8 10 12 14 we got 10 electrons left so we've got plenty of electrons i'll give them to the more electronegative oxygens first and that leaves me to give one pair to the carbon and then we say are all the central atoms happy well no this is the only one that's currently not happy no filled octet only two four six and so we might say well why doesn't one of the atoms next to him share well we can do that and we come up with a problem and we'll find out that this really comes down to a problem in formal charge and that's going to make this not a very likely structure to exist but that's in the next lesson so i don't want to defer to that now but what i will say is this is we've got atoms in this structure that are not making their typical number of bonds we'll find out the next lesson that usually means they're going to have a formal charge which usually leads to greater instability in this case so the carbon here has one two three bonds we learned that carbon having four valence electrons which is four short of a field octet usually means carbon is going to make four bonds like this one is so that's not the normal number of bonds and then oxygen having six valence electrons typically makes two bonds to get a filled octet but this oxygen's making three and so we've got two atoms bonded to each other both of which are not making their typical number of bonds so and this is probably not going to be the best structure we could come up with for this molecule so if we start this over again so what we'll end up doing is having the carbon bonded to both oxygen so this carbon is now a central atom to both of these and that one of those oxygens is then bonded to a hydrogen now if we do the same thing here and once again we've used 2 4 6 8 10 12 14 electrons we've got 10 left and we'll fill up those outer atoms and now we've used all 24. and again the only atom not happy is our central atom here this carbon so who's only got six electrons around them so somebody next to him needs to share well we've got two options for this case the both of these auctions have a pair they could share however it turns out one's going to be better and we usually look at it again one being better than the other from a standpoint of formal charge but again that's in the next lesson so i'm not going to allude to that just yet but we can talk about making the normal or typical number of bonds auction usually likes to make two well this auction right here is already making two but this one here is only making one so we're going to have him be the one that shares we'll erase a pair put a double bond and now all of a sudden every atom in this structure is making the typical number of bonds this carbon's now making four this oxygen is making two this oxygen making two and this carbon's making four life is good and this is the proper lewis structure for acetic acid now i'm not saying that it's impossible for an atom to make anything other than its typical number of bonds it is possible so however if you can make a structure where every atom's got a typical number of bonds and another structure where it doesn't the one where every atom is making its typical number of bonds is better and when you see this cooh by the end of chapter two you're gonna have to get good at recognizing that oh that means this kind of lewis structure right here but cool this is your typical lesson on lewis structures so hopefully this is mostly review from gen chem but if you're looking for some more practice and if you want the uh study guides that go with this lesson by all means check out my premium course on chatsprep.com and if you've benefited from this give me a like give me a share and put a big smile on your face happy studying