welcome to the first video for chapter 4 section 4 lewis symbols and structures in this video we'll be working on writing and drawing lewis symbols and structures the learning objectives are given here let's begin by uh defining what a lewis symbol is so a lewis symbol is a depiction of the valence electron configuration of atoms and monatomic ions what that means is it's a symbol that shows what element you're dealing with and how many electrons are in its valence shell and that looks like what we've got here in this table we have the elemental symbol so for sodium that's n a surrounded by dots that represent the valence electrons the dots tend to show up in singles up until we have to pair them they show up in sort of four pairs four regions around the uh around the symbol and there can be up to eight in general around the elemental symbol we can use them not only to just show uh elements or ions but we can show the formation of ions with these um because we can show the electrons that are changing in their configuration around the elementary ion so for example we can start off with sodium which has one electron and if we're forming the ion it'll lose that electron and become sodium plus which is different than sodium because it no longer has this electron plus we will generate one electron on its own we can also use the same notation to show the formation of anions so for example sulfur which starts out as we have it in the table it starts out with uh six valence electrons so two pairs and then two electrons hanging out over on the sides and if we're going to form an ion with sulfur that's going to form the sulfide anion which has a two minus charge um because that's going to be most similar to the noble gas configuration so we'll add two electrons here and then we can see that sulfur will have eight electrons so let me just switch colors really quick and i'll just show you the two new electrons are going to come in here and it's gonna we need to show it with the two minus charge so that's how we can use this notation to show the formation of ions we can also use this same notation to show electron transfer during the formation of ionic compounds so for example if we're going to form sodium chloride we're going to start off with sodium and we're going to add some chlorine and here we'll just show one chlorine atom even though we know that chlorine is generally a diatomic and we as we form sodium chloride we know that the sodium element is going to become sodium uh plus the ion so what happens is that this electron needs to get transferred over to this chlorine in some way and we can show this just by using our lewis notation so we show that sodium has become an ion and then we'll show the formation of the chloride ion as well um and we'll just go ahead and put this in brackets just to denote that it's different and here we can see that we've got the sodium chloride compound formed by the attraction between the sodium cation and the chloride anion so we'll move on to thinking about lewis structures instead of lewis symbols and lewis structures are quite similar to lewis symbols they shared a lot of the same notation but what's different is that instead of looking at individual atoms or elements now we're going to be looking at molecules and polyatomic ions that are sharing electrons and lewis structures essentially are drawings that describe that bonding in molecules and polyatomic ions and allow us to kind of look at what electrons are being shared so i'm going to show you what this might look like with two chlorine atoms that are going to form a a a single covalent bond um when they form their uh their diatomic molecule so i'm drawing the two chlorine atoms in different colors just so we can keep the electrons straight and what this looks like when they are have formed that bond and are in their diatomic molecule is that there's just two electrons that are going to be shared between these two chlorine atoms so we can see that these two electrons here are being shared uh each of these electrons or each of these chlorine atoms now has eight electrons around it which mimics the noble gas configuration they have a full p subshell and um so this is the this is the sort of more stable version of chlorine this is why chlorine exists as a diatomic molecule um we will often represent this instead of two dots that are just kind of sandwiched between these two atoms just for clarification we'll often represent this structure with a line between those two atoms and we'll still draw all those valence electrons as dots around them but we will use this line to represent the single bond between these two atoms um the single bond of course is the the shared single pair of electrons between those two atoms we're going to talk about how to draw lewis structures in the next video but uh for now there's a few rules that we should learn and the first important rule is the octet rule and the octet rule is the tendency of main group atoms to form enough bonds to obtain eight valence electrons and so of course eight is where the octet comes from and this is again similar because that's uh they tend to do this because that is a full p subshell uh so that mimics the noble gas configuration uh i've bolded tendency here because it's not a hard and fast rule although we do call it the octet rule it is hard and fast for some elements but we are going to talk about exceptions to the octet rule a little bit later down one thing that's pretty cool about the octet rule is that the number of electrons that's needed to reach an octet is predictive of how many bonds an atom can form so what that means is if you're looking at an element in the periodic table so for example carbon you'll note that carbon has four valence electrons so in order to form an octet this this atom is missing four electrons so uh in order to get to eight electrons it would need four more and so we can think that carbon might tend to form four bonds and in fact that is the case this is especially this is especially true for the top row uh p elements so carbon nitrogen oxygen and fluorine this is really predictive carbon tends to make four bonds nitrogen tends to make three bonds oxygen tends to make two bonds and fluorine tends to make one bond so and again that's just because of the electrons that are needed to reach the octet based on the the valence electrons that the atom has in its neutral form and we will get some practice drawing uh these as i said in the next video but real quick i just wanted to show you a couple of really common compounds that follow this pattern so um for example for example carbon tetrachloride so a carbon with four chlorines around it um and we'll talk about how you determine to form these structures so this is a really common uh example of a carbon atom with four bonds around it each one to a different atom nitrogen a really common molecule that we see is ammonia which is nh3 and that's three hydrogens around essential nitrogen and again this nitrogen tends to make three bonds uh oxygen you have probably seen water uh so water is h2o so there's two hydrogens around the central oxygen the oxygen tends to make two bonds and then with fluorine a common example of fluorine is hf hydrofluoric acid and again this fluorine uh tends to make one bond um because of uh in the in its neutral state it has seven uh valence electrons which means it has one vacancy if it needs to make if it needs to reach nothing again this is predictive this is not always the true this is not always the case but if you are looking at a lewis structure this is a good guess to begin with and then um this is predictive it is not a rule you can uh you can break this trend but it's especially true it tends to be true for these four uh atoms so we've talked about single bonds where we've got one shared pair of electrons but what happens if a molecule needs to uh you can't you can't uh you can't fill the octets with just single bonds um in other words you uh a a pair of atoms or a molecule needs to share more than one pair of electrons to reach the full octet and essentially what happens then is you have double and triple bonds so a double bond will represent with two lines and that means uh two lone pairs so essentially two pairs of electrons and a triple bond will represent with three lines and that indicates three shared pairs of electrons between the two um between the two atoms and so i'm just going to show you two examples of what this looks like so formaldehyde which is c h2o and essentially the lewis structure of this molecule looks like this where we've got a central carbon we've got two hydrogens and then in order to reach the octet of both the carbon and the oxygen we have to share two lone pairs of or two pairs of electrons between um between those atoms and that's a double bond another example of a molecule with a triple bond is carbon monoxide which you are probably at least familiar with um it's co and essentially this one has a triple bond between the c and the o and then both atoms have a lone pair on the outside so again this is three pairs of electrons being shared between these two atoms in order to make uh to fulfill the octet structure um i will just point out here that you'll note that this where carbon is making three bonds and so is oxygen breaks this trend of the prediction that we just talked about above again that's a prediction not a rule and that's okay um if you have to do that that's okay all right so the last thing we're going to talk about in this video is exceptions to the octet rule so sometimes you can't fulfill an octet for one reason or another and there are a few cases when one that's going to be especially common and one is if you have an odd number of valence electrons so nitrogen for example has five valence electrons uh and so it's a problem if you have an odd number of valence electrons you can't fulfill an octet because you have a lone electron hanging out so a common molecule that has this is no and we will talk about the structure of no in the next video but when you have this uh basically the structure looks like this where you have a double bond oxygen has a full octet nitrogen doesn't it has a lone electron hanging out and we call these free radicals when there's an unpaired electron there's also electron deficient molecules which is a case when the central atom has fewer electrons than needed for the noble gas configuration or the full octet and this is often in group uh 2 and 13 so that's usually beryllium and boron are the are the the general players here so a molecule that we see this with is bf3 and it actually has boron with three fluorines around it um and the boron actually only has six electrons the fluorines will have the full octet i'm sorry this is a little messy but the boron actually only has six electrons it does not have a full octet um this means it's quite reactive but this is in fact the structure of uh the boron or of bf3 that we'll find experimentally and then the last exception to the octet rule that we'll see is hypervalent molecules and that is when you have a central atom with more electrons than needed for the noble gas configuration that is to say it has more than eight electrons it breaks the octet rule in the other direction this you cannot do until the third row um only things in the third row and below can exceed the octet the top row carbon nitrogen oxygen fluorine cannot exceed the octet and we'll talk a little bit about that but it has to do with the d orbitals a common example of this is pcl5 and that's going to be a phosphorus with five chlorine atoms around it and the chlorine atoms all have their valence their uh their octet and the phosphorus has five chlorines attached to it um and it actually has therefore ten electrons around it instead of 8.