all right so continuing on with the hybridization lecture um in the first part of the lecture we talked about um how things hybridize why they hybridize and we really focused on things that only had single bonds so that we would um keep the idea as straightforward as possible even though hybridization's not really that straightforward but here we're going to deal with um something with a double bond and something with a triple bond and we're going to see how this looks okay so we're going to start with this molecule so this is th this is pretty much the simplest molecule i can think of that has a double bond between these two carbons that i've highlighted in blue and so let's answer these questions over here so what is the electronic geometry around each carbon okay so the answer to that is trigonal planar okay and what we talked about in the previous lecture was that if you know the electronic geometry then you will know the hybridization and things that are trigonal planar so i see three arms coming off of each of these carbons so that would be sp two okay so this is where it differs now kind of between what we talked about with only single bonds because a lot of time when there's only single bonds the hybridization and the number of unpaired electrons is very closely related how many unpaired electrons does each carbon need and the answer here is four it needs four because if we look at this we have [Music] we have one bond here one bond here we have one here and one here so we have four bonds that we're going to form there's two that are single and there's two that work together to make a double bond but i need four unpaired electrons so when i do the hybridization drawing here it's going to look a little different than what we've seen before so i'm going to i'm going to do this in two ways i'm going to do this once for the blue and once for the red okay so our blue carbon looks like this okay so this is going to be for the blue carbon and again this is the carbon atom prior to hybridization so i need to end up with four unpaired electrons but i also need to grab an s and two of the p's so the way i'm gonna do this is i'm gonna grab the s and i'm gonna grab one of the empty peas and one of the peas that has an electron in it now when this hybridizes we're going to end up with some sp2 hybrid orbitals and then notice that we have this p orbital out here that still has an electron in it okay all right so we've done the hybridization we've made our sp2 hybrid orbitals we've made sure that we have four unpaired electrons but but there's a difference between how these electrons are now these sp2 electrons are lower in energy these are the ones the ones that i've highlighted in blue these are the ones that have hybridized and then this one electron out here that's in the 2p that is left unhybridized okay so i'm going to just move this over to the left a little bit just so i don't run out of space here okay so now we're going to do the same thing for the red carbon okay so the red carbon is is pretty much the same as the blue carbon in terms of its hybridization and shape so i have to when i do this hybridization i have to grab the s and i have to grab two of the p's one without an electron of one with an electron now these p orbitals are all pretty much the same and so what i'm going to do is i'm gonna i'm gonna do this a little different than i did the first one not that it matters which of the p orbitals i grab with an electron so i'm doing it a different way just to show you that it really doesn't matter which one of those i grabbed as long as i grab one empty and one full or one empty than one that has an electron in it so this is what we get from the red carbon okay and so these are our valence electrons and these are the electrons that we are allowed to use when we build the the molecule right or the molecule gets to use we these when it builds itself and so what we're going to do now is we're going to draw in the orbital overlap diagram and the orbitals that we're going to be using are going to be these and these okay so those are the orbitals that we're allowed to use when we do this all right so let's now kind of see how we're gonna we're gonna do this all right so we'll redraw our our structure here so we can see this and we have our blue carbon and we have our red carbon okay so the blue carbon has three sp2 hybrid orbitals so i'm going to put those down so now they're trigonal planar so the way they're going to look is they're going to have one's going to be there and that's and again this is the nucleus of the carbon so that's facing the red carbon we're going to have one that's going to come down here and that's going to overlap with this hydrogen here and then we have one that's going to come up here and that's going to overlap with that hydrogen so these are my three sp2 orbitals that i have for the blue carbon okay and now the red one is going to basically do the same thing so the red carbon is going to one of its orbitals its sp2 orbitals is going to overlap with that and so those are my three sp2 orbitals from the red carbon now remember that in the blue carbon i've got three sp2s and then i have a p orbital and that p orbital doesn't show up in my picture yet okay all right so one of the questions that i would ask i just want to kind of make sure that we're all oriented how many bonds does my picture currently have shown okay the answer that is one okay it's the bond there's there's one bond that is shown between the two carbons okay and now in my picture i'm going to use this a green to represent the 1s of hydrogen so there's a 1s here that's overlapping with that and there's a 1s here and 1s here and a 1s here so if i ask the question how many bonds are we showing now the answer is five okay we have this bond this bond this bond this bond and one of these bonds shown so we're still missing one bond okay so what we have left to make make up that last bond is we have a p orbital from the blue carbon that was not part of the hybridization process and we have a p orbital from the red carbon which was not part of the hybridization process that's those are the orbitals that are going to overlap and give us this last bond or the second bond of the double bond okay now we have to think to ourselves what is the right way to orient these orbitals and it turns out that the best way the lowest energy way to orient these is to make them come in and out of the board so the p orbital which looks like this right so that's what a p orbital looks like that's going to be coming in and out of the board so that it minimizes its interaction with these other orbitals are already laid down so then i'm going to draw it this way so this is behind the board this is in front of the board this is behind the board this is in front of the board okay now i'm going to go in and try to clean up my picture a little bit here and i don't know how easy that's going to be so and then i'm going to put an electron here and i'm going to put an electron here now the way i've drawn those those orbitals are not overlapping right but this is this is a 2p orbital from the blue carbon and this is the 2p orbital from the red carbon they're currently not overlapping now the distance between them looks like it's kind of large just because the way the picture is drawn but if you think about the distance between two nuclei it's actually not that great and remember that these clouds they represent probabilities and so it turns out that this electron is able to jump back and forth between these two orbitals and this electron is able to jump back and forth between these two orbitals and so that's the second bond that is part of this double bond okay there's two types of bond that make two types of bonds that make up a double bond one is the what we call the side to side overlap of those hybrid orbitals which are here and then we have what we call the sorry that's the end to end overlap and then we have the side to side overlap which is what we're getting from this here where those orbitals are lined up side to side and the electron is jumping back and forth between them so those are two different types of bonds okay um now in a molecule let's see if we can make this picture a little bit larger okay so in a molecule in this molecule in particular there are two types of bonds so we have this bond here this is what is known as a sigma bond okay now this is when we have end to end orbital overlap okay and so in this model i have another sigma bond here and i have another sigma bond here and i have another sigma bond here and then i have that sigma bond there so there are five sigma bonds in this molecule and those were all the ones that we laid down first okay those were the ones that we use we did using hybrid orbitals this other type of bond the one that's between these two p orbitals this is what's known as a pi bond this is when we have side to side orbital overlap okay and this bond is not as strong so in a double bond when you have a double bond it's not twice as strong as a single bond because the bonds themselves aren't the same type right so you have a sigma bond which is stronger and then you have a pi bond which is weaker so again when we looked at the the bond energies a double bond is not twice as strong as a single bond and this is why okay so there have been examples um rare examples that have been discovered relatively recently where you can have a single bond between two atoms and it would be a pi bond nothing that we will see in this class is going to be that way so if there's a single bond in a molecule or between two atoms that is going to be a sigma bond if there's a double bond that means that one of them is a sigma bond and one is a pi bond so when we look at this molecule if we know that idea i could ask how many sigma bonds does this molecule have and the answer would be five and if i said how many pi bonds the answer would be one okay now keep in mind the other thing that kind of comes up is that what i've drawn here kind of in black right what i've drawn between these two p orbitals this is one bond it looks like it's it looks like a double bond the way that we draw the the the lines up here kind of going back and forth and this one down here going back and forth it looks like a double bond but that's just one bond okay the other bond of the double bond is this which is a different type of bond okay so now we're going to look at a molecule that has a triple bond so questions that we'll just work through and kind of logically come up with with our idea here what is the electronic geometry around each carbon okay the answer is linear and again if we know the electronic geometry we will know the hybridization which in this case will be s p how many unpaired electrons does each carbon require so again the answer is four because it's forming four bonds right there each carbon is going to form a single bond with a hydrogen and then a triple bond so three different bonds that all add up to making a triple bond with the other carbon so it needs four unpaired electrons okay so the way that i'm going to draw this is going to be a little different than how we've been doing it so notice that i'm not calling you know one carbon the blue carbon this time and one the red carbon and so on okay so what i'm attempting to do here is i'm just drawing kind of this one this one diagram for carbon and both the carbons are the same in this example so they're both going to be sp hybridized so what i'm going to do is i'm going to grab that because i need to grab the empty one so that i have somewhere to put that other electron that's currently paired okay so now what i have are these s p electrons these sp orbitals that have electrons and then i have a 2p that i'm going to draw in blue and a 2p that i'm going to draw in red okay and these are the two pe orbitals and i have two of this set up right because i have two carbons in this molecule okay so now what i want to do is i want to take these to the next page with us here just so we can remember exactly what we have to work with so i have this for each carbon okay so i'm going to start with my hybrid orbitals so for the carbon on the left there's an sp orbital that has one electron and it is going to be overlapping with the 1s of the hydrogen so what i've just drawn is this bond here then we're going to have the other sp there and what that's going to do is that's going to be overlapping with the sp of the other carbon so the other carbon there's its sp that's overlapping with the carbon on the left and then we have another one of these that's going to be overlapping with the hydrogen on the right okay so now going back to the question i asked how many bonds have i drawn to this point the answer is three i've drawn the bond between the hydrogen the carbon on the left the bond between the hydrogen and the carbon on the right and i've drawn one of these bonds what type of bonds are these are these sigma bonds or these pi bonds well the answer is they're sigma bonds right these are end to end overlap okay so now what i have left are these right these are the ones that i have left they're i have two of them from the carbon on the left and i have the two from the carbon on the right so the way that they're going to be oriented is going to look like this so i'm going to draw the blue 2p orbital of one of them and the blue 2p orbital of the other one and notice that i've drawn them where i put one electron facing up and one electron facing down and that's to indicate that they're interacting in their bonding right so these are jumping back and forth and these are jumping back and forth and that's a pi bond all right so now i've used up the blue 2p orbital from both carbons so now i have another one of these bonds i still have one more to draw the most stable orientation for that other p orbital is to be coming in and out of the board so this is behind the board so i'll draw that way and this is in front of the board so what type of bond is that that i've just drawn in red that would also be a pi bond so we've got a pi bond we have another pi bond and then we have three sigma bonds okay and again we can obviously or we after practice we should be able to figure out how many sigma bonds and pi bonds something has without having to draw this picture but these pictures really do help us understand what's going on they help us visualize this and as i mentioned when you get into certain aspects of chem 1b and more so in organic chemistry these pictures become very very important in terms of how things are overlapping why some molecules react the way they do and so on okay