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Video transcript

in this video we're going to start talking about exceptions to the octet rule which we've talked about in many other videos the octet rule is this notion that atoms tend to react in ways that they are able to have a full outer shell they're able to have eight valence electrons now we've already talked us about some exceptions things like hydrogen its outer shell is that first shell which gets full with two electrons so it's trying to get to that duet rule but as we'll see there are other exceptions boron and aluminum for example they can form stable molecules where the boron or the aluminum only have six valence electrons not eight and there are exceptions in the other direction as you get to the third period and beyond well actually see atoms that can maintain more than eight valence electrons and we're actually going to see an example of that with xenon so let's just go into a few examples given what I've told you see if you can come up with the Lewis diagram for aluminum hydride so aluminum hydride has one aluminum and three hydrogen's see if you can draw the Lewis diagram for that alright now let's do this together so the first thing you want to do is account for all of the valence electrons aluminum's outer shell is the third shells in the third period here and it has one two three valence electrons and then we have three hydrogen's and each hydrogen has one valence electron and so you add all of this up together three plus three is equal to six valence electrons in aluminum hydride now the next step after that is to try to draw the structure with some covalent bonds we don't want to make hydrogen our central atom that would be very atypical and so let's put aluminum in the center and then we're going to have three hydrogen's so one two and three and then let's put some covalent bonds in here and so let's see how many valence electrons have we now accounted for this is two in this covalent bond another two gets us two for another two gets us to six so we have just accounted for all six valence electrons so we have no more valence electrons to play with let's think about how the various atoms are doing so the hydrogen's are all meeting their duet rule these two electrons in this bond are hanging around hydrogen and around the aluminum but from hydrogen's point of view and as a full duet and that hydrogen as well on that hydrogen as well but notice the aluminum over here it has two four six electrons valence electrons around it and so it's not a full octet but aluminum hydride is actually something that has been observed let's think about another example let's think about xenon pentafluoride xenon pentafluoride cation it's a positively charged ion here pause this video and see if you could draw the Lewis diagram for this all right now let's do this together if this seems unfamiliar I entered I encourage you to watch the video on introduction to drawing Lewis diagrams but what we'd want to do here is first think about our valence electrons so xenon right over here it's actually a noble gas it already has a full octet in its outer shell so it has eight valence electrons so xenon has eight valence electrons and then fluorine we've seen this multiple times has one two three four five six seven valence electrons but there's five of them so five times seven I'm gonna be drawing a lot of electrons in this so this gives us a total of eight plus 35 which is 43 valence electrons but we have to be careful this is a cation it is a positively charged molecule has a positive one charge so we have to take one electron away because of that so let's take away one valence electron to get that cation and so we are left with 42 42 valence electrons so the next step is to try to draw its structure with some basic single covalent bonds and xenon would be our preferred central atom because fluorine is more electronegative it's actually the most electronegative element so let's put xenon in the middle and then let's put some fluorines around it five of them to be specific so one two three four I'm having trouble writing an F four and then five fluorines and now let me make five covalent bonds one two three four five so just like that I have accounted for ten valence electrons because you have two valence electrons in each of these covalent bonds two four six eight ten so let me subtract ten valence electrons and then we are left with 32 valence electrons now the next step is to try to allocate some more of these valence electrons to the terminal atoms so that they get to a full octet so let me do that to the fluorines each of these fluorines already are participating in a covalent bond so they already have two valence electrons hanging out with them so let's give them each six more so let's give that fluorine six and that fluorine get six and that fluorine gets six valence electrons and that fluorine gets six valence electrons and then last but not least this fluorine gets six valence electrons so I've just given away six valence electrons to each of five fluorine atoms so that is 30 valence electrons that I have just allocated and then what does that leave me with and that leaves me with two valence electrons that have gone unallocated and the only place to now put them is on the xenon and as I said things that are lower down in that periodic table of elements especially as we get below the third period these can defy the octet rule xenon already has 10 valence electrons and I'm going to just I'm about to allocate it two more to it just like that so you allocate those two more and then we have allocated all of our valence electrons and make sure I remind myself and everyone that this is a cation so I have to put that plus charge just like this but this is something that has been observed where you can actually have an central atom like this that goes beyond an octet number of valence electrons in this case it has two four six eight ten twelve valence electrons now an interesting question is how do these atoms that are in the third period or beyond handle more than eight valence electrons and it is a matter of debate but some chemists believe that it's possible because they're able to place their electrons in their empty valence D orbitals but once again this is controversial in the chemistry community