all right so today we're going to talk about the hybridization of atomic orbitals and in order to understand why this is important we have to kind of build on um our lewis dot structure ideas okay so let's look at the structure of beryllium hydride okay so here's beryllium hydride this is the formula and one of the ways that we learn to draw lewis dot structures is so we first we look at the number of valence electrons so we we know that beryllium has two valence electrons and we know that hydrogen has one valence electron okay now we have two hydrogens here and so we have two valence electrons from that so one of the things that we will do is we'll take um beryllium and we will draw the lewis dot structure of the beryllium atom right which looks like this and then we'll draw the lewis dot structure for a hydrogen atom so that would be one hydrogen atom and this would be another hydrogen atom and then what we do is we kind of piece this together right so we would say okay well the the bond is going to be between that electron that valence electron from beryllium and that valence electron from hydrogen and on the other side we see the other okay and so here we have our structure of beryllium hydride now we don't we don't usually write it like this in fact we're not supposed to we're supposed to write it where we have a bond right instead of two electrons we want to draw it as a line indicating that there's a bond there so the question that comes up when we do it this way or should come up is how many how many lone electrons does beryllium need in order to form this type of bond okay and so hopefully we can see that it needs two right we need one for one hydrogen on the right we need one for the hydrogen on the left so requires two unpaired electrons okay so then the question is well can this happen does beryllium have two unpaired electrons so a question that we would be asked maybe earlier on in in the semester is this how many unpaired electrons does beryllium have okay so if we're asked this question what we would do is we would write out the electron energy diagram and i'm going to draw the this 2p orbital or the 2p sublevel even though beryllium doesn't have any electrons in there so when we look at this the answer to the question is zero beryllium has no unpaired electrons so then how does beryllium form these bonds right so how does it actually do this okay so this is where this idea of hybridization comes in and we've seen this before when we talk about you know averaging things together we get a hybrid okay so first thing to note is what i've drawn here is a beryllium atom okay so the orbitals we have the 2s orbitals which which are the valence electrons right those are the the ones that are going to be used in this process and then we have these empty p orbitals so this is what takes place the s orbital and one of the two p orbitals will do what's called hybridizing they will combine and they will form what are called sp hybrid orbitals now they're called sp because they're made from an s and they're made from one of the p orbitals so they're called sp orbitals and then notice that i didn't grab those other two p orbitals so they're still there okay so i have these two electrons these two valence electrons that are going to be placed into this new arrangement okay and this new arrangement is is comes about by hybridizing an snmp orbital so the electron the valence electrons one will be placed into one sp orbital and one will be placed into another sp orbital now once this hybridization takes place now we have two unpaired electrons and we can form a bond with one hydrogen and we can form a bond with another hydrogen okay so this over here this is what happens when beryllium hybridizes okay so one of the things that we can say so the question is like why does this happen so there's two reasons that i'm going to provide that are going to tell us why this happens this the excuse me the number one reason or the first reason not the number one but the the first one and i'll write this down on the next slide it allows for the proper number of electrons to become unpaired to facilitate bonding right before this happens beryllium doesn't have two unpaired electrons so something has to change okay now what we're going to do is we're going to take this picture with us because this is going to be important for the next slide and instead of redrawing it i'm just going to take it with us okay so this is what we have okay um and again all i'll write down the the reason for this hybridization in a little bit okay but um it it occurs so that we can get these on these unpaired electrons or these paired electrons to get them unpaired so that we can form the bonds that we need okay now when a bond forms and we've drawn pictures like this but we've usually kept it simpler with like s orbitals overlapping or something like that but when a bond forms it is orbitals overlapping and that's what a bond is okay now keep in mind that an s orbital would look something like this right we learned that s orbitals look like this they're spherical in shape and then p orbitals look something like this and there's three of those and these are p orbitals so what does an sp orbital look like well it looks fifty percent like an s and fifty percent like a p and so when we draw these out a lot of times what they'll look like is something like this so this would be an sp hybrid orbital okay and the way that i've done this in my picture is that these are red okay just so we know which ones i'm talking about okay and there are two of them okay they're they're the there's an sp that holds one electron there's an sp that holds another electron okay so when these orbitals overlap with the hydrogen they're going to overlap with the 1s of hydrogen because hydrogen has one unpaired electron and it's 1s it doesn't need to hybridize okay now one other thing i want to mention before we continue on is just to make sure that we're clear this and this and this represent the nucleus right that's where the nucleus is when we think about these shapes of orbitals okay so that's going to be important so that we can draw some pictures okay so now what we're going to do is we're going to draw um an overlap showing that these bonds so let's go back and and write out what we have we had our beryllium and it was bonded to a hydrogen on one side and bonded to a hydrogen on another on the other side okay and if i were to ask for that angle we would say that it's 180 degrees right okay so we have these these orbitals and so from beryllium so this is the orbital that we're going to use beryllium looks something like this it has this little nub i call it and then it has that now remember that the the lobes that we see these are probabilities so what this is telling me is that you know the electron relative to the nucleus of beryllium the one that i've drawn is going to spend more time on the right hand side of beryllium so in this orbital there's one electron in there and i've drawn it spinning up because any time we draw one electron in somewhere we always put it spinning up so that is one of the sp orbitals that we have for beryllium right now we have two of them now the other one is going to face as far away from this one as possible so the other one is going to look like this and in there there's an electron so what i've drawn are the two sp hybrid orbitals for beryllium one is facing to the right it's going to bond to that hydrogen on the right and one is facing left it's going to bond to the hydrogen on the left okay now hydrogen is going to use its 1s and so the 1s is going to overlap with the 2p and those two electrons are going to spend most of their time in between the two nuclei so this right here that's the bond that we're seeing that is the bond right it's obviously nicer sometimes to draw it in a straight line because it doesn't get as cluttered but that's the bond that's that's forming and then on the other side we have the same thing right we have this overlap and then there's the electron from the 1s of hydrogen so if we label all this stuff right this is the 1s of hydrogen this is the 1s of the other hydrogen and then we have the s p from the beryllium and then just to make sure that we're clear because this is probably a new idea for a lot of us that's another sp from brilliant so those are my four orbitals that are doing all the overlapping right one of the sps overlaps with the 1s to form the bond on the right one of them overlaps with the hydrogen on the left to form the bond on the left okay so this drawing that i've done i refer to as an orbital diagram all right so if we're ever asked to draw a picture of an over orbital overlap diagram this is what we're being asked to do okay now there's a couple nice things about this and there's probably a couple not so nice things about this but this helps me kind of see now keep in mind so i'm gonna i'm going to kind of go over this but keep in mind that that this this is one kind of big electron cloud area and this is one big electron cloud area right i mean that's where the electrons are so it makes sense that they're going to be as far away from one another as possible and sometimes seeing these pictures and seeing how bulky the area that the electrons are in helps me understand why one bond or one arm is trying to get away from the other one right because they're actually these big bulky negative regions they're not just single lines okay so keep that in mind it's one of the one of the benefits of drawing these pictures now the other maybe not non-benefit is that things get a little cluttered things you know tend to um too many things on one page but these are all trying to help us understand this idea okay now going back to the shapes of the orbitals before we move on to the next structure okay and as i mentioned i'll i'm probably going to sum up at the end the um the reasons why hybridization occurs but the first reason we found was that it allows the proper number of atom of electrons so that they become unpaired and then we can facilitate bonding okay now going back and thinking about what we learned about s orbitals they get their shape right so looking at this shape they get their shape based on the energy that the electron has and the energy the electron has is based on its attraction to the nucleus and potentially repulsion of any other electrons right that's what defines that shape and the pr rule gets its shape based on the energy the electron has its interaction with other electrons the attraction the nucleus same thing okay so when we form these sp orbitals they only form when something's going to bond okay so that's that's important so it's not as if beryllium hybridizes and then goes and forms a bond okay it's it's more like it all happens at the same time now the reason why that's important is because when this hybridization occurs the orbitals are changing their shape and the reason why they change their shape is because now the things that influence the energy of the electron have changed it's no longer being attracted to only its own nucleus it's now feeling the attraction of another nucleus and so there's new forces that are coming about which changes the energy of the electron which also changes the shape of the orbitals so thinking that the orbitals just kind of recombine and make these new shapes and then they and then they form a bond it's all really kind of this one kind of overall process that takes place at once okay sometimes it's easier to think about it hybridizing and then doing it but it wouldn't hybridize if the bonds weren't forming as it did this okay so let's now look at something that's got three bonds to it instead of this that has the two bonds okay so let's look at um boron trihydride and then the question i would ask we'll just start off with this question but actually let's draw the lewis dot structure first just so we don't think that there's anything kind of hidden from us this is the lewis dot structure and then my question is how many unpaired electrons does boron need okay well it's going to form one bond with hydrogen one bond with another hydrogen one bond with another hydrogen so the answer to this question is three okay so let's see if boron needs to hybridize so how are we going to check that so let's draw the electron energy diagram for boron which looks like this so if we were asked how many unpaired electrons does it have the answer would be one so according to this boron can only form one bond so something has to change because we know that boron trihydride does form so we know that in this case hybridization is going to take place okay now the way this is going to work is we need three unpaired electrons so we can see which ones are going to be they're going to be the valence electrons now the thing that's important when we grab these if i just grabbed the if i i'm going to grab and then i'm going to undo but if i just grab these i would have three electrons but when i combine an s and a p orbital i can only make two new orbitals so i'm still going to run into the same problem where i'm going to have two orbitals and i'm going to have three electrons which means that some of them are going to be paired so what i'm going to do instead is i'm going to grab an empty orbital okay all right now when we do this what we end up getting are three new hybrid orbitals and we're going to refer to these as sp2 hybrid orbitals and the reason we call these sp2 is because they are made from an s and two p orbitals okay so now we take we put our electrons in there and now you can see that we have three unpaired electrons you can imagine one of them is going to form a bond with one hydrogen one is going to form bond with another and so on okay now these orbitals let's go back and let's let's kind of draw the picture that we drew before where we said okay well this was what an sp orbital looked like something like that this is an sp well these are these orbitals that we just formed these new ones these are going to be more they're going to we the way we say it is they have more p character so the s p orbital is 50 s and it's 50 p these are 67 p and only 33 s so they're gonna be more p like okay so when i do this um keep in mind that s orbitals are more spherical in shape and sp orbitals are more elongated p orbitals are more elongated so a lot of textbooks will draw on sp2 to look something like this where it's extended out it's actually a longer orbital what ends up happening when you get to organic chemistry this actually plays an incredibly important role because what you might imagine is that an an sp2 hybrid orbital is going to be involved in a longer bond because it's a longer orbital and there's very small subtle differences within organic molecules where shorter bonds and longer bonds change how reactive something is so this is something that's going to come back to us for those of us who go on to organic chemistry okay now the other thing i'm going to also note is that i referred to this little part here as as a kind of a nub earlier most textbooks and i'm going to start doing it from here on out most textbooks actually leave that off it is there it's officially there that's there's represents a small probability that the electrons can be there but what we're going to do instead is we're going to draw it as if it just looks like this okay um and so these are for the most part these mean the same thing okay and and you'll see in my picture here why we we tend to do that or why we want to do that okay so i have my and again let's we'll orient ourselves just to make sure as we go through this but this represents the nucleus of the atom we're talking about and this is the nucleus and then this is the nucleus here okay and again when we leave that nub off we're doing it because it's going to make our pictures look a little cleaner okay so now i have this boron atom that's been hybridized it's got three orbitals three sp2 orbitals and we know that those are going to want to be spaced 120 degrees away from one another because that's what a trigonal planar shape will be so one of them is going to face up one of them is going to face out to the left and one's going to face out to the right and each of these contains a single electron so here i have my three sp2 orbitals that i've drawn so this is my boron and again just for emphasis this is the nucleus of the boron okay all right now these are each overlapping with the 1s of hydrogen so there's the 1s of 1 hydrogen there's the 1s of another hydrogen and there's the 1s of another hydrogen so here's my orbital overlap diagram okay and for completion we'll put this here's the nucleus of the the top hydrogen here's the nucleus of the hydrogen down there here's the nucleus of the other hydrogen and if we were to go and look at the bond angles which again these orbitals help me understand a little better we would see that this is 120 degrees okay so again the reason why we would argue that boron goes through this process and hybridizes is so that it allows for three unpaired electrons to form three bonds if it doesn't do this it's only going to form one bond so we have to have to come up with another explanation as to how it forms three bonds and this is the explanation there is hybridization that takes place okay and so what i'm going to do here is i'm going to say that these are the sp 2 orbitals from boron right there's three of them and then just so we're clear and keeping ourselves oriented this is the oneness of the hydrogen okay all right so now let's look at something that has four bonds to it so we'll draw the lewis dot structure and then i would ask the question how many unpaired electrons does carbon need in order to do this and the answer would be four right it needs one for each of these so this requires four unpaired electrons so now the question is how many does it have let's test it out let's see so what we would do is we would draw the atomic lewis dot structure sorry the the electron energy diagram for the atom and according to this carbon can form two bonds so something has to change right so we need to have four unpaired electrons so i can see the four electrons there are four valence electrons and when i grab them remember that i have to grab an empty orbital so that i have somewhere to put the fourth electron right if i grab four atomic orbitals i can make four hybrid orbitals so what we're going to see is this and maybe some of us can kind of forecast what these are going to be called they're going to be called sp3s because we have one s and three p's okay all right so this is what we have and now i can see one of those electrons is going to bond with one hydrogen one's going to bond with the other and so on okay so when i draw this picture this is where it starts to get a little bit hard to draw these because this molecule is a tetrahedral right so i've got things kind of coming in and out of the board at me so just a reminder this is what this is how we would draw this perspective drawing okay now if i want to draw the orbital overlap version of this here's my best shot at that so an sp3 is going to be like it's going to be a little bit more elongated than sp2 right because it's got more p character now so we're going to have one of these p orbitals going up to the hydrogen on top we're going to have one that's coming out to the hydrogen here and then we're going to have one that's going to kind of come out of the board at us and then we have one that's behind the board all right so this is my best way of drawing these and this is pretty conventional and so and again just so we're all clear this is the nucleus of the carbon here okay and this represents an sp3 of the carbon and there are four of them right so it's not one big thing there's four different orbitals here and to kind of go along with the way we've been doing this i'll keep my hydrogens in blue and so here i have the that's the bond to this one that's the bond to this one that's the bond to this one and then i have the one that's kind of going away from us and that would be the bond to that one so that would be my orbital overlap diagram and then these are the 1s's of hydrogen okay so sometimes pictures are prettier than other times um i'm sure when you guys draw these some of you are going to have lovely pictures um so but this is the idea right so in order to form the bonds that we need in order or to form four bonds with carbon we need four unpaired electrons carbon's normal situation is that it has two unpaired electrons so there's a hybridization that takes place we form these four identical energy orbitals these sp3 hybrid orbitals they each hold an electron and they can each form a bond okay and so far again every time we've seen this the goal was for the atom to get the correct number of unpaired electrons right and it needed to do this this hybridization ordered for this to happen okay um the other thing i will say since i didn't say it earlier but this is definitely one of the more conceptual ideas that we're gonna go over okay obviously there's no math at all in here okay um now in general the questions that were asked are oftentimes easier to answer than truly understanding the the the concept um this is one of those things that takes a while for somebody to really really understand okay all right but it's important for many reasons chem 1b is going to rely on this heavily okay so let's look at this one now nh3 so this is where so i might say something like what is the molecular geometry of nh3 and some of us are going to have a gut instinct to say that it's trigonal planar because i have three things on there but we should never answer the question about the shape until we have figured out the lewis dot structure so when we do this and again i assume we're all experts at lewis dot structures at this point when we do this what we find is that the nitrogen has a lone pair on it so there's actually four things that are connected right there's three bonds and there's three paired sets of electrons and there's one unpaired okay so my question same question that i've been asking is how many unpaired electrons does this nitrogen need the answer is three right it needs one to bond to the hydrogen needs one to bond to the other hydrogen one to bond to the other hydrogen the lone pair those that's a pair of electrons it doesn't need to be unpaired to do that so this requires three unpaired electrons okay and then the question is how many does it have well the way we've been figuring this out is we've drawn the atomic electron energy diagram right so let's do that okay so this is nitrogen so then my next question is how many unpaired electrons does it have and the answer is three so i would ask a follow-up question that does nitrogen need to hybridize and most of us if we're paying attention we would say no and i'd be very happy even though you'd be wrong right so so the idea is this still needs to hybridize but it's not for the reason i gave it's not because it needs to get three unpaired electrons because it already has it so why does it do this okay now in order for us to understand why this happens we're gonna have to think spatially okay so what i'm gonna do is i'm going to draw an orbital that i'm going to color in green now that that orbital let's say looks like that right that's a p orbital it's up and down um this one actually let's use a different color let's use let's use hot pink why not that orbital is got the same shape but it's a different orientation so that'll be what that one will look like and then we'll use this last orbital here and this is the same shape a different orientation this is the one that's kind of coming in and out of the board so the way we draw that is to you know to draw it a little bit of an angle and draw something like that okay so looking at this i'm thinking okay now these look like they should be able to form the bonds because they all have one unpaired electron in them okay so what i'm going to do is i'm going to draw the electron that they have and the electron can be on either side right so it'll spend some time up here and then it'll be below and so on and then we have this i'm going to just draw it in one place okay and then we have that one so these are my three electrons and what we have to do is we have to think what do these electrons actually look like like what are they like they're in these orbitals but what are the what do the orbitals look like relative to one another okay so what i'm going to do is i'm going to take this i'm going to copy it and i'm going to paste it right here and i'm going to change it so remember that this is the nucleus of this nitrogen atom this is the nucleus and this is the nucleus so what that's telling us is that this is really there okay and i'm going to move this one in and then i'm going to move it back out just to kind of keep it more you know a little bit cleaner but this belongs right here now you can see where it starts to get cluttered very quickly right so what i'm going to do is i'm just going to leave this out okay keeping in mind that it's actually in there though okay now if our hydrogens were to come and form a bond so there would be one that would be one hydrogen this would be another hydrogen and then we'd have a third over here but i'm trying to keep it out because i don't want it to clutter things too much so my question is this what would this angle be if this is what happened so keep in mind that the p orbitals are 90 degrees away from another so this angle this bond angle would be 90 degrees so then my question to you is what is the actual angle between a hydrogen the nitrogen and a hydrogen we know the actual angle is somewhere less than 109.5 but very close to 109.5 and actually i'm going to write that it's that it's close to 109.5 instead of less than but we know that it's less than because of our our our rule for the um lone pair so the angle is nowhere near 90. okay it's it's probably i think it's around 104 or something like that so it's it's it's around 109.5 it's much closer to that than 90. so what this is telling me is that hybridization occurs because if it did not occur the angles that we would find would be 90 degrees all of those hydrogens would be attached to a p orbital and all the p orbitals are 90 degrees away from one another so we would expect that the bond angles would be 90 degrees but they're not okay the bond angles are close to 109.5 okay so let's move this up here out of the way and then let's draw our picture so first of all we're going to grab all of these and we're going to form some sp3 orbitals and i have five electrons to put in there one two three four five you'll notice that there are three unpaired electrons fantastic those are going to help me form those bonds there's also an there's a paired set of electrons that's my lone pair so when we draw the pictures of this it's going to look a lot like the ch4 that we just saw we're going to have an orbital facing up one facing out another one coming out at us and one going away from us this is my lone pair this is going to bond to one hydrogen this is going to bond to another and this is going to bond to another and we're going to end up with a tetrahedral electronic geometry and a trigonal pyramidal molecular geometry so one of the things that is important to note here is that this so the question that you would be asked by the way so you know instead of saying draw the the orbital overlap diagram you might be asked the question what is the hybridization around the nitrogen atom and the answer that you would give would be sp3 so if you're asked for the hybridization this is what you're going to give the answer would be sp3 that's the way we answer those questions notice that this has an electronic geometry of tetrahedral it's sp3 hybridized the carbon from the previous example has an electronic geometry of tetrahedral it is also sp3 hybridized if we were to go and draw water which i'll do next it has an electronic geometry of tetrahedral so it is also sp3 hybridized the electronic geometry is going to tell us what the hybridization is and technically that's actually the opposite where the hybridization happens that defines the electronic geometry but as a student if i can recognize with the electronic geometry is i will be able to figure out what the hybridization is it's usually easier to do it that way okay so this is electronically a tetrahedral it has an sp3 hybridization okay now again when i'm done with this i will write kind of one final note where i will write the the reasons why hybridization takes place but so far there are two of them the first is that it happens so that the atom will get the proper number number of unpaired electrons in order to do the bonding and the second which is what we're seeing here which can arguably be the more important of them is that it allows for the proper spacing of the electrons around a central atom to be spaced as far away from one other as possible right so keep in mind back here if it didn't do it and i look at this i'm going to have angles about 90 degrees so things are going to be really close together they're going to try to push each other apart and they do that by allowing hybridization to take place okay so let's look at the next example which is water so the the question that i would ask here and once once we get used to this answer like i said answering the questions is actually a lot easier than kind of the way i'm presenting it but if we really want to understand it we have to kind of do it this way but the question would be what is the hybridization on the oxygen atom okay and the answer because it is electronically a tetrahedral right so i see one pair of electrons another pair another pair another prey four pairs of electrons around this that is a tetrahedral electronic geometry so this is tetrahedral so the electronic geometry is tetrahedral and anything that is tetrahedral is going to be sp3 hybridized okay so now and again the reason for that if we were to draw this out what we would find is that oxygen does have the correct number of unpaired electrons to form two bonds but the spacing is wrong so it hybridizes in order to form this tetrahedral electronic geometry so we don't actually have to draw out those those electron energy diagrams and show the hybridization it helps us understand things so that's good but if we're going to be asked what is the hybridization all we have to do is really count up how many things meaning bonds and lone pairs are attached to an atom and then we'll be able to answer that question okay all right so then let's look at this what about nf3 so if i draw my lewis dot structure of this and my question is what is the hybridization around that nitrogen well if we're looking at this we see that we have one two three four sets of electrons so it has an electronic geometry of tetrahedral and so that means that we're at sp3 in terms of the hybridization okay now when i look at the fluorines i see one pair two pair three pair and then i see a fourth pair so then this is also electronically a tetrahedral so this is also sp3 hybridized okay and all of the fluorines are the same here so these are going to be sp3 hybridized as well okay now there is um and and and the reason for this by the way the reason why we think of fluorine is is hybridizing so let's let's let's write fluorine out okay let's draw this out we just did the nitrogen in the previous example so or two examples ago so let's look at fluorine okay so this is what fluorine looks like and i'm going to move this down here okay and so when we look at fluorine so there's two things that happen so i'm telling you that that fluorine in this molecule and and they're all the same so i'm only going to talk about one at a time but the fluorine that we're talking about is sp3 hybridized now the reason for that is because those four sets of electrons around it have to be spaced apart just like we saw in nh3 so it's the same reasoning okay so when it hybridizes i know it's sp3 because it's electronically tetrahedral so those are going to combine okay and this is going to be the one that's going to form the bond and the others are going to be approximately 109.5 away from that okay so so here's the thing this is how i'm presenting it to you and it makes sense logically based on specifically based on the nh3 model that we have seen that helps us understand why this that fluorine would also want to be sp3 hybridized because it wants to space everything out if you look online or in textbooks and you are asked what is the hybridization around the fluorine it in this molecule you will come up with two answers and one of the answers is what i'm showing you here which is sp3 and the other answer that you will find it'll be about 50 of the time is that it does not hybridize now there's some complicated reason reasons as to why we don't have a consensus on this okay which i'm not going to go into here okay but some people will make the argument that this electron here is unpaired and it can bond so it doesn't need to hybridize but what we saw with nh3 is that there's there's reasons that things hybridize other than needing to have an unpaired electron right it needs to space them out properly so using that logic we would fully expect that fluorine would be sp3 hybridized now one of the reasons that we can't definitively come to a conclusion on this is because fluorine's only bonded to one thing so there's not ways for us to check bond angles to see if the bond angles are really 109.5 so i don't want to like i said i don't want to get into the reason why we we argue about this but but it's a healthy argument that we're having okay so in this class if you are asked for the hybridization of that fluorine the correct answer is sp3 at some point in your future another teacher might tell you that it doesn't hybridize and and you'll and they'll have a reason for their their belief as well okay so but for now in chem 1a sp3 is the answer all right next molecule let's look at p cl5 so when i draw the lewis dot structure of this so question what's the hybridization of the chlorine hopefully we are seeing this and we're thinking okay it's got one two three four that is tetrahedral electronically so this is going to be an sp3 okay great now the question is what's the hybridization on the phosphorus okay now one of the things that some of us may have noticed that there's a pattern when you have a center atom with two things attached is sp when you have a center atom with three things attached is sp2 when you have a center atom with three things attached sorry four things attached it's sp3 so here we have five things attached so the answer that i like to have students tell me is that it's going to be sp4 because that means that they're paying attention but it also means they're wrong so it's not going to be sp4 because in the p sublevel there's only three p orbitals to grab so it doesn't grab a fourth p orbital because there isn't one so the answer is this so let's draw this out and then we'll see we can come up with the name before i say it so for phosphorus let's draw this out all right so this is what phosphorus looks like um it can only form three bonds right the way that it currently is so it's clearly going to hybridize now the other thing i'm going to draw here which i haven't drawn yet is the 3d the 3d now gets involved so i need five unpaired electrons so i need to grab all of these i need to grab those five but i also need to grab an empty orbital to kind of unpair those that are being paired right now so that's what i'm going to do i'm going to grab a d so some of us might be wondering why don't you grab from the 4s which is lower in energy the you when you hybridize you stay within the same energy level so the 4s is not really an option it can it can combine s's and p's and d's but they all have to be in the same energy level so what we end up getting we get five of these and these are going to be called sp3 because i took the s three ps and one d um and then when i draw this out i'm gonna have one electron in each like that okay and then just for completion of my picture i have these three ds that were not involved okay so the answer is that this is sp 3 d hybridized now again once we're used to this we don't have to draw this whole diagram out i'll just see five things and anytime something is trigonal bipyramidal electronically you have sp3d okay now the other thing that this kind of shows so if we go back and think about our our lewis dot structures things that are in row two like um carbon nitrogen oxygen fluorine they can only form four bonds when we get to row three we can start forming more than four bonds and the reason for that is because once we get to energy level three the d orbitals are now available to us and we can form more complicated bonding systems okay so the reason why things that are in row three or lower can form five bonds or six bonds it's because those d orbitals are available okay all right so let's do one more of these so sf6 so my question is going to be what's the hybridization on the sulfur and what's the hybridization on the fluorine okay now in order to really do this we need to draw out our lewis dot structure okay so you really cannot answer these questions without a good lewis dot structure so the hybridization of this we've already seen is sp3 okay now one of the things one of the tools that i use that i've made is this so i'll say okay well i have an s orbital i have three p orbitals and then i have five d orbitals so this sulfur has six things on it so i'm gonna say okay one two three four five six this is sp3 hybridized anything with six things attached is going to be sp3d2 hybridized so one of the challenges that we might have is just basically remembering kind of the notation here but if we draw something like this out we can easily do this right so if i draw this out let's say i have something that has let's just let's make something up let's say that we have a x well when i look at that so what is the hybridization around the the a 1 two three four five so then i say okay well one two three four five the hybridization around this is sp3 d okay so again it's not crucial for us to really understand all of this in order to answer most of the questions that we're going to be given what is the hybridization of this and so on but if we really want to understand it drawing these diagrams out is very helpful okay and so i'm not one for just memorizing quick ways to answer questions but this is one of those categories where you can probably pull that off okay so let's write down the two reasons we have seen so allows for the proper number of unpaired electrons to facilitate bonding okay and that's what we saw in the first few examples and then the second reason which again i may argue is kind of the main reason um allows for the proper spacing of electron pairs around a central atom to minimize repulsion so this is an example of like why nitrogen hybridizes even though it has its three unpaired electrons to form three bonds it hybridizes so that it can space everything out and be more stable okay so this is the first of two lectures on hybridization and in the next one we're going to talk about what do double bonds and triple bonds look like which you may have noticed did not appear in this lecture at all so that will be the next lecture