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Current time:0:00Total duration:11:30

Video transcript

in the previous video we saw some steps for drawing dot structures in this video we're going to use those same steps to draw a few more dot structures but we're also going to talk about how formal charge relates to dot structures so we'll get back to this definition in a minute for right now let's draw a quick dot structure for the ammonium cation so NH 4 plus the first thing you do is find the total number of valence electrons and so to do that you look at a periodic table and find nitrogen which is in group 5 therefore nitrogen has 5 valence electrons right each hydrogen has 1 and we have 4 of them so we have 5 plus 4 giving us 9 electrons however there's a plus one charge meaning this is a cation meaning we're going to lose an electron here so instead of representing 9 in our dot structure going to represent 8 electrons and so let's go ahead and put the nitrogen at the center remember we put the least electronegative at the atom at the center except for hydrogen so nitrogen is going to go in the center here and we know it's going to have bonds to 4 hydrogen's so we go ahead and put in those hydrogen's right here and let's see how many valence electrons we've used up in our dot structure 2 4 6 and 8 so that takes care of all 8 valence electrons that we were supposed to represent so this is the dot structure and I can go ahead and put some brackets around it here and also a plus 1 charge to indicate that this is an ion and so that's the dot structure for it for the ammonium cation here let's see if we can assign a formal charges to the nitrogen and the hydrogen's and I'm going to go ahead and redraw our dot structure here and I'm also going to draw in the electrons right we we know that each of those covalent bonds consists of two electrons and go ahead and put in those two electrons right here and if I wanted to find a formal charge for let's say the central nitrogen alright what I would do is think about the number of valence electrons in the free atom so if you had a nitrogen all by itself right look at the periodic table it's in group 5 and so therefore we're talking about 5 valence electrons in the free atom 4 nitrogen from that number we're going to subtract the number of valence electrons in the bonded atom and the way to approach that is to look at your dot structure here and think about those two electrons in those covalent bonds right one of them one of them we're going to assign to the hydrogen and one of them were going to assign to the nitrogen and so we go around we do that for each one of our covalent bonds like that and so now we can see that nitrogen is surrounded by four valence electrons in the bonded atom so let me go ahead and write that so it's five minus four and so five minus four is of course plus one so we have a plus one formal charge on the nitrogen so this nitrogen has a plus one formal charge now let's do it for hydrogen here so hydrogen is in Group one on the periodic table so let me just point this out this is for nitrogen and then for hydrogen all right it's in Group one so one valence electron in the free atom and from that we're going to subtract the number of valence electrons in the bonded atom so if we look here all right we assigned one valence electron here to the two each hydrogen so therefore it's just 1 minus 1 or 0 so there's zero formal charge for all of the hydrogen's in the ammonium cation and so that's how to assign formal charges let's let's see how that applies to actually affecting our final dot structure and an example of that would be something like sulfuric acid here so the first step of course is to calculate the total number of valence electrons we need to worry about in our dot structure so once again each hydrogen is one and I have two of them sulfur is in group six on the on the periodic table so therefore it has six valence electrons oxygen is also in group six and so we have six and we have four of them right here so 6 times 4 is 24 all right so we have 24 plus 6 is 30 plus 2 is 32 so we need to worry about 32 valence electrons in our dot structure for sulfuric acid all right next thing we do is choose the central atom and once again you ignore hydrogen so it's between sulfur and oxygen and if you look at a periodic table you will see that oxygen is higher in group 6 than sulfur is therefore oxygen is more electronegative and so therefore we're going to put sulfur at the center all right so we're going to put sulfur right here and once again look at the rules from the previous video if that if that didn't make quite sense to you here so we have sulfur attached to four oxygens I'm gonna go ahead and put my four oxygens in there like that and then I have two hydrogen's and by experience you're talking about an acid here you're going to put your hydrogen's on oxygens and so we're going to go ahead and put our hydrogen's here and let's see how many valence electrons we've used up drawing this skeleton here so we have 2 4 6 8 10 and 12 so we've used up 12 valence electrons so 32 minus 12 gives us 20 valence electrons left to worry about and so remember the next step is to assign some of those leftover electrons to some of the terminal atoms but again we're not going to assign those electrons to hydrogen because hydrogen is always already surrounded by two electrons and so we're going to try to assign some electrons to oxygen and oxygen is going to follow the octet rule so let's examine let's say the top oxygen here and we can see the top oxygen is surrounded by two electrons already right there in green and so if it's we're going to give it an octet it needs six more so we have one two three four five six same thing for this oxygen down here it needs an octet so we go ahead and give it six more electrons like that alright so we also have these other oxygens over here to worry about so let me go ahead and use green again so let's let's say this oxygen over here on the left the one bonded to this hydrogen here's two electrons and here's another two for four so for that oxygen to have an octet it needs four more so that means we're only put two lone pairs of electrons on this oxygen and then we're going to do the same thing for this oxygen as well so let's see how many how many electrons did we just represent there well we had we had six on the top oxygen 6 in the bottom oxygen that's 12 and then we had we had 4 on the left and 4 more on the right so that's 8 so 12 plus 8 is 20 so we've now represented all of the valence electrons that we needed to show and let's think about this as possibly being the final dot structure right so we have an octet around sulfur we have an octet around oxygen hydrogen's fine so you might think that we are done here however let's go ahead and assign some formal charges and let's see what that does so let me go ahead and draw in some electrons here so we know that each bond consists of two electrons I'm going to go ahead and make them read here like that and let's assign a formal charge to the top oxygen here alright so this top oxygen alright I'm going to in the bond between oxygen and sulfur I'm going to give one of the electrons to oxygen and one of the electrons to the sulfur right and so I can see that oxygen is being surrounded by seven electrons in the free atom right we would expect off shouldn't have six valence electrons six minus seven in this case gives us a formal charge of negative one and so this top oxygen has a formal charge of negative one it's the same situation for this bottom oxygen here so that one has a formal charge of negative one as well let's look at the sulfur alright so if we if we examine the sulfur here and and we know right sulfur is in group 6 on the periodic table so normally six electrons six valence electrons for the free atom and in this bonding situation alright let's go ahead and we know this one oxygen goes this this one electron I should say goes to sulfur and then in this bond between oxygen and sulfur sulfur is going to get one electron and so on all the way around here so sulfur stranded by four electrons in the bonded in the bonded atom here and so it's six minus four which is a formal charge of plus two so this dot structure might look like we're done but we have a lot of formal charges right we have negative one plus two and negative one and usually molecules like to have like to minimize the formal charge and so if there's any way to get this formal charge as close to zero as possible that would be the preferred dot structure and so let's let's go ahead and redraw this really quickly and let's see if we can move some electrons around to get to minimize our formal charges now we can't add any more electrons because we've already represented all 32 valence electrons that we were supposed to so the only thing that we can do is to share some more electrons and so I took if I took two electrons from this top oxygen here so if I took these two electrons and I move them into here and if I took these two electrons right here and I move them into here to form double bonds let's go ahead and look and see what our dot structure would look like and assign some formal charges so now I would have sulfur double bonded to this top oxygen and double bonded to this bottom oxygen top oxygen has only two lone pairs electrons around it same thing for this bottom oxygen and then these oxygens are the same with the O H on the left and then we have the O H on the right right here okay so now let's look at our formal charges so we'll put in our electrons so let's go ahead and do that and once again we're going to do the same thing that we did before so assigning formal charges we'll start with the top oxygen here so the top oxygen right six valence electrons in the free atom and then if I go like that you can see there are six here so 6 minus 6 gives us a formal charge of 0 so the top oxygen 0 now same thing with this with this bottom oxygen this one's 0 let's look at the sulfur now right so normally we're talking about 6 for the sulfur and let's see how many valence electrons around the bond the bonded atom here so we do the same the same assigning of electrons that we've done before and so now we can see that that sulfur is surrounded by six so we would go 6 minus 6 gives us a formal charge of 0 and if you assign a formal charge to all these other oxygens so let's go ahead and do that really quickly so one of these other oxygens right here so I'm doing the oxygen on the right okay so let's see how many valence electrons are surrounding this atom so we have a total of 6 as well so 6 minus 6 is a formal charge of 0 so this dot structure actually gives us formal charges of 0 for everything and so this would be the preferred one and let's finally talk about octets right so we know that hydrogen we know the hydrogen is happy surrounded by two electrons we know that oxygen is happy with an octet and so you can see that all of these oxygens have an octet sulfur in this case is not by night 2 as an expanded outer shell right it's an okay for sulfur to have an expanded valence shell because it's in the third period on the periodic table and so we talked about why that's okay in the previous video and so this is just one more thing to think of when you're telling your drawing dot structures sometimes formal charge will affect the final structure of your molecule or ion