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Video transcript

if I want to draw a dot structure for boron trifluoride I need to think about VSEPR theory so valence shell electron pair repulsion the valence electrons are going to repel each other and force the molecule into a particular shape or geometry and so we're first start off by drawing the dot structure so for boron trifluoride I find boron the periodic table it has three it's in Group three so it has three valence electrons fluorine is in group seven so it has seven valence electrons I have three of them so 7 times 3 gives me 21 21 plus 3 gives me 24 valence electrons that we need to account for in our dot structure boron is less electronegative than fluorine so bronze going to go in the center like that and the boron is bonded to three fluorine atoms so go ahead and put in those three fluorine atoms like that so we just represented six valence electrons so here's 2 4 and then 6 so 24 minus 6 gives me 18 valence electrons left that we need to worry about we're going to put those leftover electrons on the terminal atoms which are our fluorines in this case fluorine Falls the octet rule so each fluorine is now surrounded by two valence electrons so each fluorine needs six more to have an octet of electrons so I'll go ahead and put six more valence electrons on each of the three fluorines and six times three right is 18 so we just represented 18 more valence electrons and so now we are all set we've represented all 24 valence electrons in our dot structure and some of you might think well boron is not following the octet rule here and that is true it's okay for for boron not to follow the octet rule and to think about why let's assign a formal charge to our boron atom here and so remember each covalent bond consists of two electrons I'm going to go ahead and draw in those electrons in blue here and when you're assigning formal charge remember how to do that you take the number of valence electrons in the free atom so which is 3 from that number you subtract the number of electrons in the bonded atom and so when you look at the bonds between boron and fluorine one of those electrons goes to fluorine and one of those electrons goes to boron here we can see that's the same for all three of these bonds and so now boron is surrounded by three valence electrons in the bonded atom three minus three gives us a formal charge of zero so remember the the goal is to is to minimize a formal charge when you're drawing your dot structures and so this is a completely acceptable dot structure here even though boron isn't following an octet now boron can be surrounded by eight electrons and so you'll even see some textbooks say well this one of these lone pairs of electrons on one of these fluorines could actually move in here to surround the bond with with eight electrons giving it an octet and that's fine that would give the boron a formal charge and that's okay and that might actually contribute to the overall structure of this molecule but for us for our purposes we're just going to stick with this as being our dot structure here so let me go ahead and redraw that and I'm going to draw in a slightly different way when we talk about our next step here for predicting the shape so let me go ahead and put in our lone pairs of electrons here on the fluorine and let's think about let's think about step two we're going to count the number of electron clouds that surround the central atom here so remember electron clouds are either the bonding electrons or nonbonding electrons the valence electrons and bonds or the lone pairs just regions of electron density that can repel each other alright so if I'm looking at my central atom which is my boron I can see that here here are some electrons alright so that's an electron cloud these electrons right here occupy an electron cloud as well and then I have another electron cloud here so I have three electron clouds that are going to repel each other and that allows us to predict the geometry of those electron clouds around that central atom they're going to try to get as far away from each other as they possibly can and it turns out that happens when those electron clouds are on the same plane so I'm going to draw a sheet of paper here and write in a plane and we're going to put our boron atom in the center and here's one of our electronic clouds and then here's another one and here's our third one here so they're going to get as far away from each other as they possibly can we call this shape trigonal planar so let me go ahead and write that here so this is a trigonal planar geometry of my electron cloud outs surrounding my central atom and since we don't have any lone pairs of electrons to worry about on our central atom we can go ahead and predict the geometry of the molecule as being as being the same here as the geometry of the electron cloud so the molecule has a trigonal planar shape as well now when we think about bond angles right so the easiest way to think about what's the bond angle for trigonal planar or what would we predict the bond angle to be right you think about a circle right and since the electrons repel each other equally you want to divide your circle into three equivalent angles so 360 degrees divided by 3 is 120 so you can think about all these bond angles here as being 120 degrees and so you can think about those electron clouds repelling each other equally so we have a we have a trigonal planar geometry with a bond angle of 120 degrees so three electron clouds let's do another example of a molecule that has three electron clouds so in this case sulphur dioxide right so sulfur dioxide let's count up the number of valence electrons sulfur is in group six so six valence electrons oxygen is also in group six six times two is of course 12 12 plus six is 18 so we have 18 valence electrons for our dot structure sulfur is less electronegative right if you look at a periodic table so sulfur is below oxygen so sulfur is going to go in the center sulfur is bonded to two oxygens like that so that is represented four valence electrons right here's two and here's two more so 18 minus 4 is 14 valence electrons left all right so we're going to start assigning some of those leftover electrons to our terminal atoms which are our oxygens oxygens going to follow the octet rule and so therefore each oxygen gets six more electrons to give each oxygen and octet so go ahead and do that so I just represented 12 more valence electrons right six on each oxygen so 14 minus 12 right is two valence electrons so I have two valence electrons left over and remember when you have leftover valence electrons you go ahead put them on your central atom so we can go ahead and put those two valence electrons here on our central atom like that now we're not quite done with our dot structure because sulfur doesn't have an octet as one way of thinking about it you could also think about formal charge sulfur's formal charge is not minimized in this dot structure so we need to we need to we need to share some electrons right so I could take let's say let me go ahead and make these blue here so I could take a lone pair of electrons from either oxygen I'm just going to say that we take a lone pair of electrons from that oxygen and move them in here to form a double bond between the sulfur and the oxygen so if I do that right now I have a double bond between the sulfur and the oxygen the oxygen on the right now has only two lone pairs of electrons around it the oxygen on the Left still has three lone pairs of electrons around it like that and the sulfur still has a lone pair of electrons here in the center all right so if we if we assign formal charges now let's go ahead and do that really fast so we know that we have electrons in these bonds here and so if we assign a formal charge to the oxygen on the Left let's do that one first alright so this oxygen on the left here right oxygen normally has six valence electrons in the free atom and in our dot structure right we give one of these electrons and blue to the oxygen one to the sulfur so you can see the oxygen is surrounded by seven valence electrons this one's a little bit hard to see so let me go into market there so six minus seven gives us a formal charge of negative one on this oxygen and when we do the same formal charge for the sulfur here all right we can see that sulfur is surrounded by five valence electrons sulfur is in group six so in the free atom there's 6 6 minus 5 gives us a formal charge of +1 so the formal charge of +1 to solve for a formal charge of negative 1 on this oxygen and even though um it's it's it's we don't have a formal charge of 0 on these atoms this is about as good as we're going to get in terms of this representation of the molecule here and another way that another thing to think about is the fact that I didn't have to take the lone pair of electrons from this oxygen right I could have taken the lone pair of electrons from over here alright I like on that oxygen and that would just be another resonance structure of this so I don't want to I want to go to in detail about resonance structures in formal charge we talked about those in earlier videos to make sure that you watch them but this does this this is our final dot structure here so let me go ahead and redraw it here this is one of the possible dot structures you can draw for sulfur dioxide right that fulfills our rules for drawing dot structures so let me go ahead and put in these valence electrons here so you can see it a little bit better all right so I'll leave out formal charges now because we're just focusing in on geometry we're concerned about VSEPR theory so we have our dot structure we go back up to check our steps for predicting the shape and now we are going to count the number of electron clouds that surround our central atom so regions of electron density which could be valence electrons in bonds so bonding electrons or nonbonding electrons like lone pairs of electrons so we look at our central atom right which is sulfur and let's see if we can count up our electron cloud so this is an electron cloud over here there's a region of electron density so we have one electron cloud over here on the right this double bond we can consider it as an electron cloud we're not worried about numbers of electrons just regions of them and then for the first time we now have a lone pair of electrons right and this right this we can also think about as occupying an electron cloud alright so we have three electron clouds all right and we saw in the previous example that when you have three electron clouds the electron clouds are going to try to adopt a trigonal planar shape so I could redraw the stop structure and attempt to show it in more of a trigonal planar shape here all right so let's go ahead and show it looking like this all right once again not worried about drawing in formal charges here so something like this for the structure let me go ahead and put those electrons in our orbital here so we can see that that electron cloud a little bit better and so once again our electron clouds are in a trigonal planar geometry so we would expect those bond angles to be approximately 120 degrees right so let me go ahead and put this in here so approximately 120 degrees it's probably slightly less than that but that is what we would predict the geometry of the electron cloud to be so let's go back up and look at our rules here our steps for predicting the shapes of molecules so we've we've done step three right we have predicted the geometry of the electron clouds around our central atom and now we go on to step four here we're going to ignore any lone pairs on our central atom when we predict the geometry of the overall molecule and so that now now pertains to our example here we're going to ignore that lone pair of electrons on the sulfur right we're going to ignore this lone pair of electrons we're talking about the shape of the molecule so even though the electron clouds have a trigonal planar geometry we say that the shape of the molecule has a bent or angular shape and so if you look at that if you just look at the atoms right if you ignore the lone pairs and look at the atoms right you'll see this kind of bent or angular shape here and so that's what we say is the shape of the molecule so you could say bent or you could say angular here alright so that's two examples of molecules with three electron clouds and it's not necessarily conclaves you have to think about you have to ignore lone pairs of electrons to predict the final shape of the molecule