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  • - [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.

- [Instructor] We're told

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工作實例。使用形式電荷來評估非等效共振結構|可汗學院。 (Worked example: Using formal charges to evaluate nonequivalent resonance structures | Khan Academy)

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    林宜悉 發佈於 2021 年 01 月 14 日
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