If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

Main content
Current time:0:00Total duration:14:58
AP.Chem:
SAP‑4 (EU)
,
SAP‑4.C (LO)
,
SAP‑4.C.1 (EK)
,
SAP‑4.C.2 (EK)

Video transcript

let's use VSEPR theory to predict the structure of this molecule so phosphorus pentachloride so the first thing we need to do is draw a dot structure to show our valence electrons we find phosphorus in group five so five valence electrons chlorine in group seven so seven valence electrons and I have five of them so seven times five is thirty five plus five gives us a total of 40 valence electrons that we need to show in our dot structure so phosphorus goes in the center because it is not as electronegative as chlorine and we have five chlorines so we go ahead and put our five chlorines around our central phosphorous atom like that if we see how many valence electrons we've drawn so far this would be two four six eight and ten so 40 minus ten gives us 30 valence electrons left over and remember you start putting those leftover electrons on your terminal atom so we're going to put those on the chlorines and the chlorine is each one is going to follow the octet rule so that means each chlorine needs six more electrons and so now each chlorine is surrounded by eight valence electrons like that so if I'm adding six more electrons to five atoms six times five is thirty and so I have now represented all of my valence electrons on my dot structure notice that phosphorous is exceeding the octet rule here right there are ten valence electrons around phosphorous and it's okay for phosphorous to do that because it's in period three on the periodic table I like to think about formal charge and so if you assign a formal charge to phosphorous you'll see it has a formal charge of zero and that helps to explain for me anyway the resulting dot structure so now step two we're going to count the number of electron clouds that surround our central atom remember an electron cloud is just a region of electron density all right so I could think about these bonding electrons in here as a region of electron density around my central atom I could think about these bonding electrons too so here's another electron cloud and you can see we have a total of five electron clouds around our central atom the next step is to predict the geometry of the electron clouds and those electrons those valence and shell electrons are going to repel each other right so the VSEPR theory valence shell electron pair repulsion since they're all negatively charged they're going to repel and get try to get as far away from each other as they possibly can in space and so when you have five electron pairs it turns out the furthest they can get away from the other in space is a shape called a trigonal by pyramidal shape so let me see if I can draw our molecule in that shape we're going to have our phosphorous in the center and we're going to have three chlorines on the same plane so let me attempt to show three chlorines on the same plane here these are called the equatorial the equatorial positions because they're kind of like along the equator if you will so three chlorines in the same plane one chlorine above the plane and one chlorine below the plane those are called axial positions all right so there's there's a quick sketch let me see if I can draw a slightly better shape of a trigonal by pyramidal shape here so let me see if I can draw one over here so you can see what it looks like a little bit better so we could have we could have one pyramid looking something like that and then down here let's see if we can draw another pyramid in here like that so that's a rough drawing but we're trying to go for a trigonal trigonal by pyramidal shape here so let's focus in on those chlorines that are the same plane first so if I'm looking at these three chlorines all right and I go over here to my to my trigonal bipyramidal shape you could think about those three chlorines as being at these corners here so it's a little bit easier to see they're in the same plane so those are the equatorial chlorines when I think about the bond angle for those alright so those chlorines being in the same plane right you have these you have these three bond angles here and so when we did trigonal planar right we talked about 360 degrees divided by 3 giving us a bond angle of 120 degrees so you could think about that as being a bond angle of 120 all right so same idea those bonding electrons are going to repel each other when we focus in on our axial chlorines right so this one up here and this one down here you could think about those as being here and here on your on your trigonal by pyramidal shape like that and if you draw if you draw the axis right if you draw like a line down this way all right connecting those it's easy to see those are a hundred and eighty degrees from each other so you could think about a bond angle of 180 degrees between your chlorines like that and then finally if we think about the bond angle between let's say this axial chlorine up here at the top right and then one of these one of these green chlorines right here I think I think it's a little bit easier to see that's ninety degrees here so this bond angle right here would be 90 degrees and so those are your those are your three I those are your ideal bond angles for a trigonal bi-pyramidal situation here it's important to understand this trigonal by pyramidal shape because all of the five electron cloud drawings though that we're going to do right are going to have the electron clouds want to take the shape so it's important to understand those positions all right so for step four ignore any lone pairs and predict the geometry of the molecule well there are no lone pairs on our central phosphorus so the electron clouds take a trigonal bipyramidal shape and so does the molecule let's go ahead and do another example so sulfur tetrafluoride here so we're going to start by drawing the dot structure and we need to count our valence electrons of course so sulfur is in group six so six valence electrons fluorine is in group seven so seven valence electrons I have four of them 7 times 4 is 28 28 plus 6 is 34 valence electrons we know sulfur is going to go in the center because fluorine is much more electronegative so we put sulfur in the center here we know sulfur is bonded to four fluorines so we put our fluorines around like that and let's see how many valence electrons we've shown so far 2 4 6 and 8 so 34 minus 8 gives us 26 valence electrons we still need to account for on our dot structure we're going to start by putting those leftover electrons on our terminal atoms which are our fluorines fluorine is going to have an octet of electrons around it therefore each fluorine needs six since each fluorine already has two around it so we go ahead and put six valence electrons around each one of our fluorine atoms alright so we are showing six more valence electrons on four atoms six times four is twenty-four so 26-24 gives us two leftover valence electrons and remember your rules for drawing dot structures when you get some leftover electrons you're going to go ahead and put them on your central atom now so we have a lone pair of electrons on our sulfur and by adding that lone pair of electrons to our sulfur the sulfur now exceeds the octet rule but once again it's okay for sulfur to have an expanded valence shell it's in period three on our periodic table and once again I like to think about formal charge and if you assign a formal charge that sulfur it has a formal charge of zero so that just helps me understand these dot structures a little bit better so we've drawn our dot structure let's go back up and remind ourselves of the next step here so once you complete step one next is the electron cloud step right so how many electron clouds do you have surrounding your central atom so we go back down and we look at our electron clouds that surround our central atom here so regions of electron density so we know that these bonding electrons here right that would be that would be one electron cloud same with these bonding electrons same with these bonding electrons and same with these bonding electrons and then we have a lone pair of electrons on our sulfur well that's also a region of electron density surrounding our central atom so that lone pair you could think of as being an electron cloud as well and so we have five electron clouds so just like in the previous example and when you have five electron clouds those electron clouds are going to try to adopt a trigonal by pyramidal shape just like we saw again in the previous example so let's go ahead and draw two possible versions of the dot structure for this molecule alright so I'm going to I want to draw one right here and for this first version I'm going to show I'm going to show the lone pair of electrons on the sulfur in the equatorial position here someone put the lone pair of electrons right here right equatorial here and so that means there are two fluorines also equatorial and then that means that there's one fluorine here axial another fluorine here axial so that is one possible dot structure the other possibility would be of course to put the lone pair of electrons in the axial position so if we do that we would have a sulfur bonded to three fluorines those would be the equatorial fluorines and then we would have a lone pair of electrons let's just put it right here in the axial position and then another fluorine in the axial position so here are our two possibilities so let's see if we can analyze this structure now when you have lone pairs of electrons in your dot structure lone pairs take up more space or nonbonding electrons I should say take up more space than bonding electrons and so since they take up more space they're going to repel a little bit more and so that means that when you're trying to figure out valence shell electron pair repulsion the lone pairs of electrons are more important to focus in in terms of where you're going to put them and so let's focus in on those lone pairs of electrons and let's think about how they're going to repel the other electrons in these two dot structures and so let's say let's look at the left here where we have the lone pair of electrons in the equatorial position here and if you're thinking about how they're interacting with let's say these bonding electrons in in in the same plane here this is about 120 degrees between these between the bonding electrons and the nonbonding electrons and it turns out that 120 degrees is not as important in terms in terms of repelling as say something like 9 degrees so you tend to ignore the 120 degree interactions when you're analyzing these structures however a 90-degree angle between between a bonding pair and a nonbonding pair and we have that example for let me go ahead and show you right here so let's think about this lone pair of electrons alright repelling these bonding electrons so in the axial position well these two are only 90 degrees away so remember 9 degrees of course being closer you're going to get more repulsion from this interaction than in the previous interaction so we're going to focus in on the 90 degree interactions here so those bonding electrons and nonbonding electrons repel each other and you have one possibility right with the axial fluorine you also have another possibility with with this axial fluorine you have essentially essentially you have you have a lone pair of electrons nine degrees from two pairs of bonding electrons from the example on the left and of course that's going to destabilize it somewhat but let's compare this dot structure with the one on the right now so the one on the right all right so we have our lone pair of electrons in the axial position this time and you can see that we have we have three fluorines in the equatorial position so you have these bonding electrons in the equatorial position which means that that lone pair of electrons is ninety degrees to all three of those and so that of course is going to cause some serious repulsion so nine degrees two three and so you have an example on the right you have these three interactions nine degrees and example the left you have only two of these and so the goal of course is to minimize electron pair repulsion so VSEPR theory actually predicts that this dot structure on the left is the correct one and so you're going to see in the next video that nonbonding electrons are placed in equatorial positions in trigonal bipyramid to minimize electron pair repulsion so just think about putting your lone pairs of electrons in the equatorial position so the structure on the Left wins so let's go ahead and redraw that so we can analyze it a little bit better all right so I have my sulfur in the center here let's go ahead and change colors so I'm going to put the sulfur in the center all right I have my fluorine in the plane in a plane another point in the plane my lone pair of electrons in a plane alright so those are my equatorial ones I have a fluorine it this way and I have a fluorine that way so when you're looking at bond angles right of course between this fluorine a sulfur fluorine bond angle the ideal bond angle anyway would be 120 degrees all right so we can say and we would expect it to be 120 degrees if you're talking about this axial fluorine and this equatorial one we know those are we would expect that to be 90 degrees and then finally between the two axial fluorines right so this bond angle back here would of course be 180 degrees okay so we've done a lot of a lot of talking and we still haven't even talked about the final name for the shape of this molecule right let's go back up here and look at our rules really fast all right so we've done a lot of work to predict the geometry of the electron clouds around the central atom and draw it and finally we get to predict the shape of the molecule and we do that by ignoring any lone pairs and so let's go ahead and do that so we're going to ignore the lone pair of electrons on the sulfur when we're talking about the shape so if we ignore the lone pair and we actually turn this molecule on its side so let's go ahead and do that we're going to put our sulfur here if we turn outside the axial fluorines would now be horizontal so those are my so now it's horizontal like that so I'll go ahead and put in my so these are what the two fluorines that used to be axial there and my two fluorines that were equatorial it would look something like this and it helps if you actually build this molecule with a MOLLE mod set so that that would be what the molecule kind of looks like here and we call this a see-saw shape all right so this is a see-saw shape or geometry and let's think about why right so if you've ever been on a playground and used a seesaw and draw a little kid here on one side of our seesaw like that and so if the little kid puts his weight on this side of course this side of the seesaw would go down right and then this side of the seesaw would go up so just a little bit of intuition as to why you would call this a seesaw shape all right so I think we've we've will have to stop there and the next video we'll do we'll do two more examples of molecules and ions that have five electron clouds
AP® is a registered trademark of the College Board, which has not reviewed this resource.