- Dot Structures Questions
- Drawing dot structures
- Formal charge and dot structures
- Resonance and dot structures
- VSEPR for 2 electron clouds
- VSEPR for 3 electron clouds
- VSEPR for 4 electron clouds
- VSEPR for 5 electron clouds (part 1)
- VSEPR for 5 electron clouds (part 2)
- VSEPR for 6 electron clouds
Definition of formal charge, and how minimization of formal charge can help choose the more stable dot structure. Created by Jay.
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- It's ok for Sulfur to have 6 bonds although it isn't in the 4th orbital or higher?(32 votes)
- The octet rule can be broken by elements starting in the 3rd period and below. It is attributed to the not so large energy difference between 3p and 3d orbitals, which allow for additional shared pairs beyond the octet. Examples of this include SF6, PCl5, SO4(2-), etc.(56 votes)
- Can't we use the formula of Formal Charge as :
FC = ""Total Number of Valence Electrons in free atom - Total Number of electrons in Lone Pairs - 1/2 Total Number of Bonding Electron""(15 votes)
- Yes this formula works out to give you the correct answer.
Total number of electrons in lone pair + 1/2 Total number of bonding electrons = Number of valence electrons in bonded atom(13 votes)
- 1) at03:15, how come the formal charge be +? does that mean the molecule is cation?
(#V.E. in free atom - # V.E in bonded atom) does that mean the molecule has one more electron that is not bonded? then doesn't it make the NH4+ anion? as it has one more electron?
2) at05:08, is there any reason why S is attached to 4 Os first and Hs are attached at each end? why can't Hs be attached to S directly? Thank you(13 votes)
- NH3 is a neutral atom, N has 3- charge and H +1, in NH4 the N forms a dative covalent bond with the H (since H will probs lose it´s electron, +1), nay, the overall charge will be+1(1 vote)
- I dont get how to draw a dot structure when calculating a formal charge....As it is a bit different with respect to the usual dot structure! Please help me out!(5 votes)
- We draw the dot structure in the exact same manner, and then calculate the formal charges for the atoms in the molecule.
Remember that formal charge is calculated by taking the # of valence electrons, minus the lone electrons and the bonds, and we show that charge next to the molecule.
Take ::O=C=O:: for example. Each O's formal charge would be calculated by: 6 (valence) - 4 (lone electrons) - 2 (bonds) = 0. C's formal charge is 4 (valence) - 4 (bonds) = 0.
Hopefully this helps you.(16 votes)
- Why is the Nitrogen in the center if you mention that the least electronegative atom in the center (wouldn't it be Hydrogen in this case)?(3 votes)
- H can't go in the centre because it can bond to only one other atom. H must always be a terminal atom.(16 votes)
- Why is a lower formal charge preferred?(8 votes)
- It takes energy to remove electrons and create a positive charge, so a lower formal charge usually indicates a lower energy level.(5 votes)
Q: Is there a different method to calculate the formal charges?
I think my teacher has a different method.
( V electrons) - (the bonds)- (the number of free electrons).
Is it correct or did I make a mistake ?(3 votes)
- Both methods are correct.
The method your teacher used is best when you know the electron-dot structure of the molecule ( e.g, H-O-SO₂-OH).
The methods used in the video are best for when you don't know the electron dot structure (e.g. H₂SO₄). They are just summaries of the results obtained from your teacher's formula.(14 votes)
- What is the difference between formal charge and partial charge?(5 votes)
- The formal charge is the charge that an atom appears to have when we count the electrons according to certain arbitrary rules.
The partial charge is the charge that the atom really has.(6 votes)
- one question - why do we need formal charges?(3 votes)
- Formal charges help us keep track of the electrons in a molecule. The formal charge tells you whether an atom has more electrons (negative charge) or protons (positive charge) associated with it. This can be helpful when predicting how chemicals react, since areas with excess electrons (negative charges) often react with areas lacking electrons (positive charges).(5 votes)
- But what is formal charge? Not the formula, the meaning, in lay man's terms.(3 votes)
- It’s the hypothetical charge an atom would have if every covalent bond was broken.
It’s a bookkeeping tool for electrons.(4 votes)
The previous video, we saw some steps for drawing dot structures. In this video we're going to use the same steps to draw a few more structures. But we're also going to talk about how formal charge relates to dot structure. So we'll get back to this definition in a minute. Right now, let's draw a quick dot structure for the ammonium cation. So NH4 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 five. Therefore, nitrogen has five valence electrons. Right, each hydrogen has one. And we have four of them. So we have 5 plus 4, giving us 9 electrons. However, there's a plus 1 charge. Meaning this is a cation, meaning we're going to lose an electron here. So instead of representing nine in our dot structure, going to represent eight electrons. And so let's go ahead and put the nitrogen at the center. Remember you put the least electronegative 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 four hydrogens, so we go ahead and put in those hydrogens right here. And let's see how many valence electrons we've used up in our dot structure. Two, four, six, and eight. So that takes care of all eight valence electrons that we were supposed to represent. So this is the 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 the ammonium cation here. Let's see if we can assign formal charges to the nitrogen and the hydrogen. So I'm going to go ahead and redraw our dot structure here. And I'm also going to draw in the electrons, right? We know that each of those covalent bonds consists of two electrons. I'm going to go ahead and put in those two electrons right here. And if I want to find a formal charge for, let's say, the central nitrogen. What I would do is think about the number of valence electrons in the free atoms. So if you had a nitrogen all by itself, right? You look at the periodic table, it's in group five. And so therefore we're talking about five valence electrons in the free atom for 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. One of them we're going to assign to the hydrogen, and one of them are 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 5 minus 4. And so 5 minus 4 is of course plus 1. So we have a plus 1 of formal charge on the nitrogen. So this nitrogen as a plus 1 formal charge. Now let's do it for hydrogen here. So hydrogen's in group one on the periodic table. So let me just point this out. This is for . Nitrogen and then for hydrogen. It's in group one. So one valence electron in the free atom. And from that we're going to subtract a number of valence electrons in the bonded atom. So if we look here, we assigned one valence electron here to each hydrogen. So therefore, it's just 1 minus 1, or 0. So there's 0 formal charge for all the hydrogens in it the ammonium cation. And so that's how to assign formal charges. 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, I have two of them. Sulfur is in group six on the periodic table, so therefore I have 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. 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 the periodic table, you'll see that oxygen is higher in group six than sulfur is. Therefore oxygen is more electronegative. And so therefore, we're going to put sulfur at the center. So we're going to put sulfur right here. Once again, look at the rules from the previous video if that didn't make quite sense to you here. So we have sulfur attached to four oxygens. And I'll go ahead and put my four oxygens in there like that. And then I have two hydrogens. And by experience, you are talking about an acid here. You're going to put your hydrogens on oxygens. And so we're going to go ahead and put our hydrogens here. And let's see how many valence electrons we've used up a drawing this skeleton here. So we have two, four, six, eight, ten, and twelve. 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 left over electrons to some of the terminal atoms. But again, we're not going to assign those electrons to hydrogen because hydrogen's already surrounded by two electrons. And so we're going to try to assign some electrons to oxygen. And oxygen's going to follow the octet rule. So let's examine, let's say the top oxygen here. And we could see the top oxygen is surrounded by two electrons already right there in green. And so if 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. Right so, we also have these other oxygens over here to worry about. So let me go ahead and use green again. So let's say this oxygen over here on the left, the one bonded to this hydrogen here. Here's two electrons and here's another two for four. So for that option to have an octet, it needs four more. So that means we're going to 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 electrons did we just represent there? Well we had six on the top oxygen, six in the bottom oxygen. That's 12. And then we had four on the left and four more on the right. So that's eight. 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. So we have an octet around sulfur, an octet around oxygen, and 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 red here like that. And let's assign a formal charge to the top oxygen here. All right, so this top oxygen. I'm going to, and 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. And so I can see that oxygen is being surrounded by 7 electrons. In the free atom, right? We would expect oxygen to have six valence electrons, 6 minus 7 in this case gives us a formal charge of -1. And so this top oxygen has a formal charge of -1. It's the same situation for this bottom oxygen here, so that one has a formal charge of -1 as well. Let's look at the sulfur. So if we examine the sulfur here, and we know-- right sulfur is in group six on the periodic table. So normally six electrons for the three atom. And in this bonding situation, right let's go ahead, we know this one oxygen-- 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 is surrounded by four electrons in the bonded atom here. And so it's 6 minus 4, which is a formal charge of plus 2. So this dot structure might look like we're done, but we have a lot of formal charges. We have -1, plus 2, and -1. 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 0 as possible, that would be the preferred dot structure. And so let's go ahead and redraw this really quickly. And let's see if we can move some electrons around 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, 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 are on it. Same thing for this bottom oxygen. And then these oxygens are the same, with the OH on the left. And then we have the OH on the right, right here. OK 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 will 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 is 0 now. Same thing with this bottom oxygen, this one's 0. Let's look at the sulfur now. So normally, we're talking about 6 for the sulfur. And let's see how many valence electrons are on the bonded atom here. So we do 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 0. And if you assign a formal charge to one of 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. OK, so let's see how many valence electrons are surrounding this atom. So we have a total of six is well. So 6 minus 6 is a formal charge of 0. So this dot structure actually gives us formal charges of zero 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's 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 surrounded by an octet. It has an expanded outer shell. Right, it's OK for sulfur have an expanded valence shell, because it's in the third period on the periodic table. And so we talked about why that's OK in the previous video. And so this is just one more thing to think of when you're drawing dot structures. Sometimes formal charge will affect the final structure of your molecule or ion.