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Worked example: Using formal charges to evaluate nonequivalent resonance structures

AP Chem: SAP‑4 (EU), SAP‑4.B (LO), SAP‑4.B.2 (EK)

Video transcript

- [Instructor] We're told that three possible resonance structures for the thiocyanate ion are shown below. All right, there we have them. Based on formal charges, which of the three structures contributes most to the resonance hybrid of thiocyanate? And they have given us some extra information. They've given us the various elements in these resonance structures, and they've told us their Pauling scale electronegativity, so maybe that is going to be useful for thinking about basing on the, based on the formal charges, which of the three structures contributes most to the resonance hybrid of thiocyanate? So pause this video and see if you can figure that out. All right, now let's work through this together. So there's really two things we want to optimize for when we're thinking about which of these resonance structures contributes most to the resonance hybrid? One, we want to figure out the resonance structures where individual atoms have formal charges as close to zero as possible. So let me write that down. Individual, individual atoms have formal charge as close to zero as possible. As close to zero as possible. We're not talking about the charge of the entire ion. We're talking about individual atoms' formal charges, close to zero as possible. And then the electronegativity is useful because we also want to see if there's any negative formal charge on an individual atom that ideally, that would be on the most electronegative of the atoms. So any formal charge, so once again, we're not talking about the charge of the entire ion. Any formal charge, any negative, any negative formal charge on individual atom, individual atom, ideally, ideally on most electronegative ones, or most electronegative one. Electronegative. All right, now with these two principles, let's figure out which of these resonance structures get closest to these ideals. So to do that, let's just calculate the formal charges in each of these resonance structures. So the way that we do that is for each of these elements, if you had just a free atom of it that was neutral, how many valence electrons would it have? And actually, let me make another column right over here, which is just the valence electrons. You can look it up on a periodic table of elements or you might already know that carbon has four valence electrons, six total, but four in that second shell. Nitrogen has five valence electrons, a neutral nitrogen, seven overall electrons, but it has five in its outer shell, and sulfur has six valence electrons. And the way that we calculate formal charge of the individual atoms in each of these resonance structures is we say, all right, how many valence electrons would say, sulfur, a neutral, free sulfur atom typically have? And we know that that is six. And then we say, well, how many outer electrons are hanging out around the sulfur in this resonance structure? And the outer electrons that we see here, it's really from this Lewis diagram, we can see one, two, three, four, five. So five electrons versus six valence electrons in a typically neutral sulfur free atom, and so it's one less electron. So you would expect a plus one formal charge here. Another way you could think about it is typically, six valence electrons and, but we are only seeing five hanging out in this Lewis structure, so that's where we get our plus one from. Now we can do the same exercise for the carbon here. Carbon typically has four valence electrons when it's neutral, and this Lewis structure, in this resonance structure, we can see that four outer electrons are hanging out, the same as you would expect for a neutral carbon atom. And so four minus four, you have zero formal charge here. And then for the nitrogen, we have one, two, three, four, five, six, seven. We can say outer electrons hanging out. Neutral nitrogen would have five valence electrons, so five valence electrons, we have two more than that. Five minus seven is negative two. So since we have two more outer electrons hanging out than we would typically have for a neutral nitrogen, we have a negative two formal charge. Now let's go to this resonance structure here. So same idea. Here, we have one, two, three, four, five, six outer electrons hanging out, the sulfur. Now that's the same as a neutral sulfur valence electrons. So here, we have no formal charge. You could think about it, six minus six is equal to zero. Carbon, we have four outer electrons hanging around from this Lewis diagram, and that's typical of the valence electrons of a neutral carbon, so once again, four minus four, we have no formal charge there, and then we move onto the nitrogen. We have one, two, three, four, five, six outer electrons hanging out. Nitrogen would typically have five. Five minus six, we have one extra electron hanging out, which gives us a negative one formal charge, the nitrogen right over there in this resonance structure, and then last, but not least, in this resonance structure, we have one, two, three, four, five, six, seven electrons hanging around, outer electrons hanging out around the sulfur. Neutral sulfur would have six valence electrons. Six minus this seven, we have one extra electron. That's what gives us this negative one formal charge for the sulfur in that resonance structure. The carbon is still having four hanging out, which is typical of carbon and neutral carbon's valence electrons, so no formal charge there, and then the nitrogen has one, two, three, four, five outer electrons hanging out, which is equivalent to a neutral nitrogen's valence electrons, and so five minus five, you have no formal charge. So there you have it. We've looked at the formal charges on all of these, and now let's look at these ideals. So individual atoms have formal charges close to zero as possible. In this first resonance structure, we have two individual atoms whose formal charges are not zero, and in fact, nitrogen is quite far from zero, while in these other two resonance structures, we only have one atom whose formal charge is not zero. So I'm liking, just based on this first principle, I'm liking these second two resonance structures as contributing more to the resonance hybrid than this first one. So I will rule that one out, and then if we had to pick between these two, we could go to the second principle. Any negative formal charge on an individual atom, ideally on the most electronegative. So in this resonance structure here, I guess the second resonance structure, the negative formal charge is on nitrogen. While on this third one, the negative formal charge is on sulfur. And we can see from this table that nitrogen is more electronegative than sulfur. So it's in the second resonance structure, you have the negative formal charge on an atom that is more electronegative than nitrogen than in this third resonance structure, and so this is the one that I believe contributes most to the resonance hybrid of thiocyanate for these two reasons.
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