Simple Structure, Formal Charge and Resonance Part II

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En provenance du cours de Duke University

Introduction to Chemistry: Structures and Solutions

67 notes

Duke University

Introduction to Chemistry: Structures and Solutions

67 notes

This is an introductory course for students with limited background in chemistry; basic concepts such as atomic and molecular structure, solutions, phases of matter, and quantitative problem solving will be emphasized with the goal of preparing students for further study in chemistry.

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Formal Charge, VSEPR Theory, Hybridization

<p>Welcome to week 3! This week includes introduction to the octet rule and expanded octets, formal chargevalence shell electron pair repulsion (VSEPR) theory, sigma and pi bonds, hybridization of the main group elements, and introduction to molecular shapes.</p> <p>I hope you enjoyed the last two weeks of learning chemistry. Great work on watching videos, posting on the discussion forum and submitting your exercises!</p>

Simple Structure, Formal Charge and Resonance PartI13:20

Simple Structure, Formal Charge and Resonance Part II11:52

Rencontrer les enseignants

Prof. Dorian A. Canelas
Assistant Professor of the Practice
Chemistry

[BLANK_AUDIO]

So if formal charge exists then our resonance structures.

What if both structures we draw have formal charge.

Or what if I can draw three or four

resonance structures and they all have some formal charges.

Then which one's the best.

Well, the most stable, or lowest energy, or best, if you will, Lewis dot structure,

is the one that has the negative charges on the more electronegative elements.

So the lowest lowest energy is if you can draw one that has no formal charge.

But if you have to draw a formal charge,

put the negative formal charge on the more electronegative element.

To have the lowest energy structure.

And that means that positive charges then in contrast on the

flip side, would prefer to be on the less electronegative elements.

Let's do an example. Let's do an example of this molecule.

N2O.

Oops, white is a terrible color choice. Let's do green.

N2O.

Now N2O if you do a calculation to calculate

how many electrons and bonds and lone pairs there are.

N2O has four bonds.

So if I draw N2O first of all, the first thing I need

to recognize is because nitrogen is written first, nitrogen is in the middle.

Oxygen is not in the middle.

So I can arrange the two nitrogen's and the oxygen different ways,

but I'm always going to put one of the nitrogen's in the middle.

[SOUND] All right, so I've drawn the molecule three times there.

I have to place four bonds,

between these atoms.

So I'm going to start by putting a single bond

on the left and a triple bond on the right.

Or I can put a double bond in each bonding space.

Or I could put a triple bond between

the nitrogen's and a single bond between the oxygen.

Then I can go in and fill in the lone pairs.

Let's just add them without even counting first to give everything an octet.

I already have an octet on the middle nitrogen.

I can do the same thing here,

and finally on the bottom structure

that's where the electrons would go.

Let's check and make sure we have the right number of electrons.

How many valence electrons does each atom bring?

Well, we can look on the periodic table and see, nitrogen's in group five.

So I have two nitrogens, they each bring five.

Oxygen's in group six, so it brings xix electrons.

I should have 16 electrons in my picture.

And if I count in all the pictures, two, four, six, eight, ten, 12, 14, 16.

I have 16 in all

of them, so I have the right number of electrons in all of the pictures.

But the pictures are not equivalent.

We need to assign formal charges.

I'm going to fix these long curves aren't very good.

I'm going to make' em look a little better.

How about this nitrogen on the top structure on the left.

What's it's formal charge?

It should have five electrons, because it's nitrogen.

How many does this nitrogen

formally own? One, two, three, four, five, six, seven.

So it's got a minus two formal charge, doesn't it?

How about the nitrogen in the middle?

Nitrogen should have five, and this nitrogen has one,

two, three, four, so 5 minus 4 is plus 1.

How about the oxygen on the right?

Well oxygen should have six electrons because

that's how many valence electrons it has.

This particular oxygen doesn't have six it has five.

One, two, three, four, five.

6 minus 5 equals plus 1, so that also has a positive formal charge.

I can, so that's actually not looking very good.

It's got a lots of charge on all the atoms.

What about the middle structure?

Nitrogen should have five electrons and this one nitrogen

formally has one, two, three, four, five, six so

this one is a minus one formal charge.

That's a little bit better than the one above it in terms of that nitrogen.

The central nitrogen still has a positive formal charge.

It should have five and this one only owns four.

So 5 minus 4 is plus 1 there.

And the oxygen on the right has, should have six electrons.

And this oxygen does have six.

So no formal charge on that one.

If I was going to draw boxes around these structures,

I don't need to draw the box around the calculation,

do I, but I could draw a box around that one.

And now I can draw a box around this one.

Right?

What about the bottom structure? Not quite finished with that one yet.

Nitrogen should have five valence electrons.

And this nitrogen on the left of the bottom structure does have five.

So no formal charge on the nitrogen on the left.

Nitrogen in the center has a plus one formal charge.

And the oxygen here should have six and this oxygen for has only

one, two, three, four, five, six, seven or a minus one formal charge.

So that structure now is finished.

I have three structures to chose from.

Right? Which of these structures

is the most stable? What do I

have to look at to determine stability?

Well, first I have to look at minimizing formal charge.

So, the Todd's structure is probably not

very good because it doesn't minimize formal charge.

The next thing I need to look at is this rule that I've given you.

That the

negative charges preferred to be on the more electronegative element.

Well, what's more electronegative?

Nitrogen or oxygen, which one's more electronegative?

[SOUND] Look on the periodic table, remember?

Fluorine's the most electronegative element.

If we look on the periodic table up there where fluorine is.

[SOUND] Here, I'll just draw a little part of it.

Carbon, nitrogen, oxygen, fluorine.

[SOUND]

Right, and under fluorine is chlorine, right?

So oxygen is closer to fluorine, it's more electronegative.

Oxygen's more electonegative.

So which of these two structures then is better, the

one on the bottom or the one in the middle?

The one on the bottom because it's got the formal charge on the oxygen.

So I'm going to put, I'm going to put a slash through this one.

It's not, it's not the lowest energy.

So the one on the bottom is the lowest energy.

Now, all three of these resonance structures do exist.

I wasn't drawing lines through them to imply that they don't exist.

Okay.

They all exist.

But, in fact here, they're actually listed in order of their relative energy.

If energy was on the y axis, I'll just write energy going up.

Right, the lowest energy one is the one that's on the bottom.

The middle energy is the one that's above it, and the highest energy one

is the one with all that charge separation, that has the minus two charge.

We can interconvert between resonance structures, can't we?

How do we do that?

Well, in order to do that, we're going to push electrons away from negative charge.

And toward more positive charge.

The pi electrons and the lone pairs are the things that have been moving.

Have you noticed that the only thing I've been

moving is the pi electrons and the lone pairs.

I have not been breaking any sigma bonds and I have not been moving any atoms.

So if we think back to the, the structures we just had.

Let me just draw a couple of them.

I'll draw the one that was with the lowest energy again it was this one.

[SOUND] If I wanted to convert that to a higher energy structure.

Let's say I wanted to convert it to the one that had the double

bonds between both pairs of atoms what would I do in terms of pushing electrons?

[SOUND]

[BLANK_AUDIO]

Well, I'm going to start where the negative charge is.

That's where I'm going to start my first

arrow, is where I have the negative charge.

And start there.

And I'm going to push away from

the negative charge toward the positive charge.

So here's the negative charge.

I'm going to, in fact let me go back to the other

one, and before I do that, put my formal charges in, right.

Okay.

I'm going to start at the negative charge

on this oxygen push towards the positive charge.

The bonding, the lone pair here is becoming the bonding pair.

It's becoming this pi bond.

That puts more than four electrons around the nitrogen and it cannot handle that.

It is not large enough.

It doesn't have a big enough atomic radius for

that many electrons for more than an octet around it.

So it can't have more than four bonds. And so it has to kick out this pi bond.

And that pi bond becomes a lone pair on the nitrogen.

Becomes that lone pair right here.

So I can draw a little double headed resonance

arrow between those two to show that I have resonance.

But I do have to be careful and realize that this is

a higher energy resonance form and this is the lower energy resonance form.

The ones that's more stable, I've drawn lower around the page.

Let's do another example. This will the last example.

Let's do an organic example. I have tried of these inorganic molecules.

How about these two structures?

Do either of these have another resonance form?

Well, I look at the one on the lower left, and I look

at my rules, and one of my rules says, push electrons towards positive charge.

I don't have a loan pair that I can move, do I?

But I do have a pi bond that I can move.

This pi bond is close

enough to that positive carbon that I can take

that, those two electrons from the pi bond and

move them all the way over here to make

a pi bond between the two carbons on the right.

If I do that, draw a double headed arrow, I can draw another resonance

structure, so I've got two hydrogen still on the carbon that's

on the left, now I have a double bond, there.

[BLANK_AUDIO]

This was the new bond that I just made and that

leaves a positive formal charge on the carbon that's on the left.

But some of you don't like this picture because

you say well Dr. [UNKNOWN] has only four bonds.

And these structures don't have for bonds, that's true.

This is an unstable species.

This is what's called a reactive intermediate.

It's a carbo cation in both cases.

But this carbo cation is a little more stable than your everyday carbo cation

because it can spread the positive charge over two different carbon atoms, can't it?

Some of the positive charge is on this carbon atom,

and some of the positive charge is on that carbon atom.

What about the structure on the right? What about this one?

Does this one have any other resonance structures?

This is the Tert-Butyl carbo cation.

We'll see this

guy a lot. Might as well write the name now.

Tert-butyl carbo cation.

Do I have any lone pairs that I can move? No.

Do I have any pi electrons that I can move?

No.

So this structure on the right does not have any resonance structures.

And I show you this to point out that not everything has resonance.

Sometimes when I cover resonance, students start trying to

draw resonance structures of every single compound in the universe.

When in fact resonance is a fairly special thing that only happens for some

compounds or some species.

So this concludes our brief review of

Lewis Dot Structures formal charge in Resonance.

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