welcome to the second video for chapter four section six on molecular geometry and polarity in this video we'll be focusing on practicing predicting geometry using vesper the learning objective is to predict the structures of small molecules using valence shell electron pair repulsion theory which is also known as vesper theory in a previous video we talked about some steps to predict the geometry of some molecules and the first step of this is to draw the lewis structure for the molecule i'll be working with these molecules down here um so what i recommend you do is pause the video at this point and spend some time drawing the lewis structures for these molecules um and in one ion and just make sure that you're comfortable drawing that lewis structure alanine is a more complex molecule which we will be talking about at the end so um if you want to look at the structure for that molecule you can or just wait until the end and we'll talk about it all right so once you've got your lewis structures drawn we will go ahead and walk through how to predict the geometry of each molecule our first molecule is water i have drawn all the structures here with basically no attempt at representing the shape and i've done that on purpose so that you can kind of see how you might go about drawing these with some sort of attempted shape to represent the actual reality of the bond angles etc all right so our first step after we've drawn our lewis structure is to count the regions of electron density my favorite way of counting regions of electron density is trying to draw a circle around the region and if i can do that without too much trouble i'm trying not to you know go out of my way to include extra electrons or have to be really focused to avoid certain electrons then that's probably one region so single bonds are pretty clearly a region of electron density so so far we have two regions of density and then each lone pair is also a region of electron density so all together we've probably got four regions of density here the next step is to use that that count of regions of density to identify the electron pair geometry for this molecule i can do this either by having memorized um what that means how many how many regions and what shape corresponds to that number of regions which you should do if you haven't memorized it yet then we'll just go ahead and use our chart um so in your textbook there is this handy chart and if you can't see it well enough it's in your textbook you can go ahead and zoom in on it there the way we're going to use this chart is by first identifying the number of electron pairs what this means is the number of regions of electron density where there are electron pairs and we're going to come on down to this row where we've got four regions of electron density this means that our electron pair geometry is tetrahedral so we'll go ahead and write down tetrahedral i'm just going to write down tetra as an abbreviation the next thing we're going to do is use our lone pair count to see if the molecular geometry is different from the electron pair geometry which i should note here this is the electron pair geometry so we're going to see um if we have any lone pairs that our molecular geometry will be different from our uh our molecular geometry will be different than our electron pair geometry and again that's because uh when you are describing the shape for the molecular geometry you don't include the lone pairs in the description of that shape so we've got two lone pairs here so we'll come back over to our chart we'll come back down to our our row of four regions of density and then we'll go across until we find the spot where there's two lone pairs so this tells us that our shape is bent or angular and that our bond angle is less than 109 degrees so we'll go ahead and write that down our molecular geometry is bent and our bond angle is less than 109 degrees all right so now let's go ahead and try to draw this guy in some kind of realistic fashion so we'll start off with our central atom and then we're going to just go ahead and try to draw our hydrogens with some sort of angle of around 109 degrees um that's an estimation and then we'll just put our two lone pairs um also at something estimating 109.5 degrees those guys are actually turned um they are coming into and out of the page but since they're lone pairs we don't care too much about representing them you can draw them at the end of dashes and wedges if you if you do want to represent that so i'm just going to note here that this angle is less than 109 degrees all right so we'll move on to our next molecule which is uh nitrite the anion one of your polyatomics so we talked about this one in a previous video when we were discussing resonance and uh so this this molecule or this ion has resonance you can draw it in two different ways which means that neither of these is actually the real representation of the shape of this molecule but rather um somewhere in the middle right to an average of these two so the double bond isn't here or here it's somewhere in between it's kind of both at the same time um this is one of the reasons that we like to draw resonance forms or resonance contributors because it helps us understand um the shape without having to kind of go through the process of understanding that partial double bond thing that thing that's not quite a single bond but it's not really a double bond either instead we can just consider the shape of each of these resonance contributors and um they will actually be the same and it turns out that is actually the shape of this molecule so we're going to start out the exact same way we just count up our regions of density so a double bond is one region um i would have to try pretty darn hard to circle only some of those electrons so this whole double bond is one region uh our single bond is a second region and then our lone pair on that central nitrogen is a third region so since we've got three regions of density we will go over to our little chart and find out that that means our electron pair geometry is trig planar or trigonal planar but i'm going to abbreviate it trig whoa oh geez pen pen malfunction trig planar all right now we're going to look at our lone pairs and and uh use that to understand if and how the molecular geometry is different than the electron pair geometry we have a lone pair so it's going to be different we'll go back over to our chart um and we come across from our in our third row or our row of three regions of density across to one lone pair and we find that this is also bent or angular but the bond angle here is actually less than 120 instead of 109.5 as it was above so here we actually have two different kinds of bent molecules here we have um tetrahedral bent with the bond angle less than 109 and down here we're actually going to have trig planar bent with the bond angle less than 120. so this is one of the reasons why it's super critical to start with your electron pair geometry and then think about your molecular geometry because the electron pair geometry is what sort of defines your shape to begin with and then how many lone pairs just kind of modifies that shape it just modifies the bond angles and the name of the shape so if we're going to draw this guy um with some sort of reality i'll just pick one of these resonance structures to draw although we could draw both or we could draw the resonance hybrid that would be the same thing but i will just try to draw this guy with some sort of representation of the angle between these two oxygens as something in the ballpark of less than 120. all right so we'll move on to our next molecule which is carbon tetrachloride so we'll start off the same way we count up our regions of density one two three four and we can check on our chart but we've already done one of these so when we have four regions we know that that means our electron pair geometry is tetrahedral the next thing we're going to do is see if our molecular geometry is different than our electron pair geometry here we have no lone pairs so our molecular geometry is not different it is also tetrahedral and therefore we know that these angles are 109. okay so this guy's a little bit challenging to draw with some sort of semblance to reality um and this is where our dashes and wedges are going to come in really handy so i've got this little molecule that may help you visualize what's happening here so what you're going to do is draw your first atom as your central atom and then you're going to pick two two of the surrounding atoms here are chlorines to to be in the plane of the paper or the screen here i'm going to draw this guy and this guy in the plane of the screen just because that is uh easy to me um you can draw any two but whatever you do they're going to have an angle between them that's approximately 109 i'll draw the lone pairs on at the end the next thing you're going to do is think about that there are two molecules or two two atoms at the end of the bond um these guys are essentially in a plane that's 90 degrees off from these guys they're exactly perpendicular to the screen except they're also tilted a little bit so we're going to use our dashes and wedges to um to sort of understand that so i'm going to draw a wedge here for this chlorine and then i will draw a dash going back for that chlorine and then i'll just go ahead and add in my lone pairs around my chlorines to represent this structure and all of these bond angles are 109. all right so then we'll move on to our next molecule which is iodine pentafluoride this is a hypervalent molecule it's clearly got more than eight electrons around our central iodine but luckily for us that doesn't actually change anything about the way that we assign geometry we're gonna start off the same way just counting up the regions so we have one two three four five regions um that are bonding and then we have a sixth region that's a lone pair so we've got six regions and we can go to our chart and discover that that means our electron pair geometry is octahedral and that all of these species should be something in the ballpark of 90 degrees off from each other the next thing we're going to do is use our lone pair count to determine if our molecular geometry is different and since we have a lone pair it will be different and that single lone pair means that our actual geometry or sorry our molecular geometry is square pyramid square pyramidal square pyramidal or square pyramid um and i'm just going to abbreviate that square here okay and our bond angles are actually less than 90 degrees because that lone pair squishes everybody down a little bit okay so here's another one that's a little bit challenging to draw um but we'll give it a shot so we're going to go ahead and start off same thing with our central iodine atom and then the easiest way to do this is to pick one guy to be axial and the other four will be equatorial and then you can draw the axial one either up or down it's the same thing so i will go ahead and draw this with my axial fluorine going up and then the other four fluorines are in a plane that's perpendicular to the screen so they're coming like straight in and straight out so i will go ahead and draw two of my fluorines going straight back but at an angle into the plane so i guess for you guys it's back coming towards me and then um the other two are coming straight out of the screen with wedges and then i'll go ahead and just add in my lone pairs real quick and then of course i've got a big lone pair hanging out in the other axial position on the iodine so all of these angles are less than 90 degrees both the equatorial and the between the axial and the equatorial all right so we'll move on to our next molecule and that is carbon dioxide so hopefully this guy is familiar to you at this point we've talked about it a few times so we're going to start off exactly the same way count our regions two regions so two regions means that our electron pair geometry is linear um we can't have uh yeah when you have only two regions of electron density you can't have one of those be a lone pair because then you can't it's not a central atom anymore um so this is linear and and you can actually see this on your chart there is no such thing as a lone pair for a linear central atom so um so we're done and the bond uh bond angle is 180 so in fact we don't have to rewrite this one because it's linear and we tend to write it as linear so that's it all right our next molecule is phosphorus pentachloride so again another hypervalent guy so we'll just start off the same way as normal and we're going to count up our regions of electron density one two three four five so we've got five regions so our electron pair geometry is going to be we can check our chart trigonal bipyramidal um so i'm going to just abbreviate that as trig by here whoops pure all right and then we check to see if we have any lone pairs uh we don't so that electron pair geometry is the same thing as our molecular geometry and here we actually have a bit of interesting geometry the um we have some axial positions and we have two equatorial positions as well and um and they're going to have different bond angles from each other luckily this is a little bit easy because we don't have any lone pairs our electron pair geometry is the sort of simple model version so we'll go ahead and try to draw this before we describe the bond angles all right so we'll start off with our phosphorus in the center and then we're going to have um kind of the same way that we had our we had an axial and we had equatorial positions for the uh the octahedral guy with with six regions of density around it we're going to have axial positions here as well and then we'll have equatorial positions so i'm going to go ahead and start off by putting my axial positions in the plane of the board and then my equatorial positions that's where the other three chlorines are going to go is directly around this in kind of a again a flat plain sort of around um and so i'm going to choose one of these guys to be in the plane of the screen or the paper and go ahead and pop in that chlorine it doesn't matter which way it goes and then i will put the other guys on the other side one of them is going to be going back into the screen at an angle or i guess sorry back into the screen at an angle and then the other one would be coming straight out of the screen at an angle so i'll go ahead and uh pop my lone pairs on really quickly and then we'll talk about our bond angles so between the um axial position and the equatorial plane is 90 degrees but between all the positions in the equatorial plane it's 120. so here's where you have this uneven distribution of ankles all right and then the last thing we're going to do is talk about alanine and so essentially the reason i'm talking about alanine is because it's going to help us kind of understand what to do when you don't have one single central atom when you've got more than more than one single central atom and essentially the answer is you're going to consider every single central atom individually so we're only looking at local geometry about each of these atoms um there is especially with proteins if you go on to biology biochemistry especially you'll find that we'll talk about larger structure um but we're not going to cover that at all in gen chem we're going to be focusing on local structure so i'm going to focus on just a couple of atoms which i will circle we're going to focus on this carbon this nitrogen um and then we'll do this carbon as well okay um yep okay so uh we're going to focus on this carbon over here first we'll go ahead and figure out what its geometry is we've got one two three four regions of density around it hey that's tetrahedral it doesn't have any lone pairs so its molecular geometry is identical this is just a tetrahedral carbon it turns out that we can go ahead and just see that this guy is also tetrahedral since we've just done that one this guy also has four bonds around it all right now let's focus on this nitrogen this nitrogen has one two three bonds and one lone pair so this nitrogen has four regions of density so it's electron pair geometry is tetrahedral tetrahedral but its molecular geometry is not tetrahedral since it has that one lone pair so if we come over here we can see that when we have four uh regions of density but one lone pair that's trigonal pure middle and so that lone pair sort of pushes everybody down a little bit more and the bond angle is less than 109. so i will abbreviate this as trig pier okay and then we'll come and look at this carbon over here we've got one two three regions of density so this is trig planer all right so if i'm going to try to redraw this it gets a little bit complicated so what i'm going to go ahead and do is draw the carbons in the plane of the board so i'll start with this guy and then i'm going to go ahead and say this hydrogen and then this carbon-carbon bond is in the plane of the board this carbon um has another thing that's in the plane of or the screen which is the other bond to this carbon and then this carbon is trig planar so since i've drawn it with one bond in the plane of the screen or the paper or the board or whatever um then everything is in that plane and then this guy is actually bent as it turns out so we'll just go ahead and draw him in as well okay so now i'm going to come back to this tetrahedral carbon and um sometimes it can help to number things to just not lose tracks we'll go ahead and do that one two three carbons so i'll just go ahead and one two three okay so this carbon here is carbon number one it's tetrahedral um i've drawn in one of its hydrogens and so the other two hydrogens need to be tetrahedral so coming out and going into the board okay this guy i have uh drawn with two bonds in the plane already and they're going this way so the other two need to be going sort of the opposite direction of them maybe it's better if i use this so right now i i've got this and this in the plane of the screen so i have one guy going this way and the other guy going this way and i'm just going to go ahead and decide that the nitrogen is the guy that's coming out at me and the hydrogen i will try really hard to draw going back in some sort of way okay so that's carbon number two and then i will go ahead and draw my nitrogen um so my nitrogen center is uh um is tetrahedral but it's actually trig pyramidal so i've drawn one bond kind of this way and i'm going to go ahead and uh just sort of draw another hydrogen kind of coming out and then one hydrogen there and then my last lone pair will be sort of up and this is really hard to imagine um and then that's my alanine molecule so you can see this gets a little bit complicated the more things you add to it the harder it is to draw in three dimensions which is why sometimes we don't try and we represent them with these sort of 90 degree angles that don't represent reality all right so hopefully that's given you some practice on assigning geometries um if you still need help there are plenty more problems in the back of the book or just google geometry practice and there are plenty of tools online