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• - [Instructor] Let's see if we can draw the Lewis diagram

• for a nitrate anion.

• So a nitrate anion has one nitrogen and three oxygens,

• and it has a negative charge.

• I'll do that in another color.

• It has a negative charge.

• So pause this video and see if you can draw that,

• the Lewis structure for a nitrate anion.

• All right, well we've done this many times.

• The first step is to just account

• for the valence electrons.

• Nitrogen has one, two, three, four, five valence electrons

• in its outer shell, and in that second shell,

• if it's a neutral, free nitrogen atom.

• So we have five valence electrons there.

• Oxygen has one, two, three, four, five,

• six valence electrons.

• But if you have three oxygens,

• you're going to have six times three.

• And so if you just add up the valence electrons,

• if these were free, neutral atoms,

• you would get five plus 18, which is 23 valence electrons.

• Now the next thing we have to keep

• in mind is this is an anion.

• This has a negative one charge right over here.

• So it's going to have one more extra electron,

• one more extra valence electron than you would expect

• if these were just free atoms that were neutral.

• So let's add one valance electron here.

• So that gets us to 24 valence electrons.

• And then the next step is let's try

• to actually draw this structure.

• And the way we do it is we try

• to pick the least electronegative atom

• that is not hydrogen to be the central atom.

• And in this case it is nitrogen.

• It's to the left of oxygen in that second period.

• So let's put nitrogen in the center,

• right over there.

• And around it let's put three oxygens.

• So one, two, three oxygens.

• Let's put a single bond between them.

• And so so far we, and let me do that

• in another color, so we can account for it better.

• So I'll do them in purple.

• So so far we have accounted for two, four,

• six valence electrons.

• So minus six valence electrons gets us

• to 18 valence electrons.

• The next step is we would try to allocate

• as many of these as possible to our terminal atoms,

• the oxygens over here.

• Try to get them to a full octet.

• So let's do that.

• This, each of these oxygens,

• they're already participating in one

• of these covalent bonds,

• so they already have two valence electrons hanging out.

• So let's see if we can give them each another six,

• to get to eight.

• So two, four, six.

• Two, four, six.

• And then two, four, and six.

• And so just like that

• we have allocated 18 valence electrons,

• six, 12, 18.

• So minus 18 valence electrons.

• And we are now left with no further valence electrons

• to allocate.

• But let's see how our atoms are doing.

• We know the oxygens have a full octet,

• but the nitrogen only has two, four,

• six valance electrons hanging around.

• It would be great if there was a Lewis structure

• where we could have eight valence electrons

• for that nitrogen.

• Well one way to do that is to take one

• of these lone pairs from one of the oxygens

• and turn that into another covalent bond.

• So let's do that.

• So let me just erase this pair right over here,

• and I'm just going to turn that

• into another covalent bond.

• And this is looking pretty good.

• We have eight valence electrons around each

• of the oxygens.

• And now we have eight for the nitrogen,

• two, four, six, eight.

• And we have to remind ourselves

• that this is an anion.

• It has a negative one charge.

• So to finish the Lewis diagram

• we would just put that negative charge there.

• And this is all well and good,

• but if this was the only way that nitrate existed

• when we observed nitrate anions in the world,

• we would expect to see one shorter bond

• and two longer bonds, and we would expect one

• of the bonds to have a different energy

• than the other two.

• But in the real world we don't see that.

• We see that all of the bonds actually have the same length,

• and they actually have the same energy.

• And so an interesting question is why is that?

• And one thing that you might appreciate is,

• when I took that lone pair

• to create this covalent bond,

• I could have done it with that top oxygen.

• I could have done it with this bottom-left oxygen.

• Or I could have done it with that bottom-right oxygen.

• And so there's actually three valid Lewis structures

• that we could have had.

• Not only could we have had this Lewis structure,

• we could have had this one,

• and I'll draw it all in yellow to save us some time,

• where you have this nitrogen.

• It has a single bond with that top oxygen.

• And so that top oxygen still has six electrons

• in lone pairs.

• And maybe it forms a double bond

• with the bottom-left oxygen.

• So this bottom-left oxygen only has two lone pairs.

• One of them would have gone to form the double bond.

• And then this oxygen would look the same.

• So what I am drawing here is another valid Lewis structure.

• Or the double bond might have formed

• with this bottom-right oxygen,

• so let me draw that.

• So another valid Lewis structure could look like this.

• So nitrogen bonded to that oxygen has three lone pairs.

• This oxygen also has three lone pairs.

• And now this one has the double bond

• and only has two lone pairs.

• And whenever we see a situation

• where we have three valid Lewis structures,

• we call this resonance.

• Resonance.

• Resonance.

• And we'll put an arrow,

• these two-way arrows between these structures.

• And when you hear the word resonance,

• it sometimes conjures up this image

• that you're bouncing back, you're resonating

• between these structures.

• But that's actually not right.

• What the right way to think about it is,

• these different ways of visualizing the nitrate,

• these contribute to a resonance hybrid,

• which is really the true way

• that the nitrate exists.

• And so, if we wanted to draw a resonance hybrid,

• it would look like this.

• You have the nitrogen in the center.

• You have your oxygens, one, two, three.

• I can draw our first covalent bond like that.

• And then you would show the bond

• between nitrogen and each

• of these oxygens are a hybrid between someplace

• between a single bond and a double bond.

• And so instead of just one of them having the double bond

• and the other two having single bonds,

• they're all somewhere in between.

• So maybe you draw a dotted line,

• something like that, to show what the reality is,

• is that you actually have three bonds

• that are someplace in between a single

• and a double bond, because the electrons

• in this molecule are delocalized throughout.

• And of course you wanna make sure,

• you always wanna make sure that people recognize

• that this is a anion.

• So this is the idea of resonance.

• You have multiple valid Lewis structures.

• They all contribute to a resonance hybrid,

• which is actually what we observe.

• We're not just bouncing between

• these different structures.

• The actual observation will be a hybrid of the three.

• Now what we just drew here,

• these three are all equivalent.

• But in certain cases, we'll see this

• in future videos, you don't have equivalent structures,

• and some of them might contribute more

• to the resonance hybrid than others.

• But we'll see that in future videos.

- [Instructor] Let's see if we can draw the Lewis diagram

B2 中上級