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.