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CATHERINE DRENNAN:
Lewis structures.
So I tell students who have sort
of no background in chemistry
before they come
into this class,
and there always are some, that
there are some topics in 5.111
that having no
experience with the topic
is actually a good thing,
and Lewis structures is
one of these things.
I think if you've seen it
before, you're like, oh yeah,
that's easy.
I know how to do that.
You don't practice
and you get on an exam
and you're like,
oh wait a minute.
I forgot how to do this.
And the students who
haven't seen it before
know all the rules and just
get brilliant, perfect scores
on the exam.
And the other students
are like, man.
I forgot how to do
my Lewis structures.
So here are Lewis structures.
We're going to go over them.
So I really like
Lewis structures
because they're relatively
simple, and they work.
Like 90% of the time, they
are the correct structure.
And I'm a big fan of
simple things working.
If you wanted to get
from 90% to 100%,
you'd have to use
Schrodinger's equation,
but you can get 90% just with
the simple Lewis structures
I'm a big fan of that.
I love it when simple things
really work pretty well.
OK.
So when Lewis structures
the key to this
is thinking about the electrons
being shared so that you
get a full valence shell.
And having the electrons
distributed in such a way
that all of the atoms have
the number of electrons
that make them happy,
which is usually
eight electrons, which is
an octet, that noble gas
configuration.
So every dot in
a Lewis structure
represents a valence electron.
And we can then
look at some atoms
and put dots around them
to indicate the number
of valence electrons.
So we also have to know how many
valence electrons atoms have.
And so why didn't
you just practice
with a clicker question.
And here's part of the periodic
table up here if you need it.
All right.
I'm told 10 seconds.
Everyone was crazy fast.
Yes.
So seven is the correct answer.
You could look at
the periodic table
and sometimes with these
it's a counting thing.
So this is one where you
want to always double check
if things don't make sense.
All right.
So we can put seven
electrons around fluorine,
and we'll have two
fluorines here.
They'll both have seven
electrons around them.
And now I'm going to
jump to another slide,
but I'm going to show you
the seven again in case
you haven't written them down.
If you don't want to,
they're not in your handout,
but that's probably OK.
So when you bring
them together, you
can bring them together in such
a way that they can all share.
And so if we put in green, then,
one of the fluorine's seven.
And then we put in blue
the other fluorine's seven,
you can see that they can
share two in the middle
and both are very, very happy.
So just thinking about
this really simple idea,
how many electrons will give you
an octet, will give you eight?
And how can you
put things together
in such a way that allows
for that to happen?
Now there are a
few elements that
do not want eight in
their valence shell,
and hydrogen is one of them.
It just has that
one S. So it only
wants two, that's
all it can handle.
So this is an
exception, hydrogen
is going to want two electrons.
Hydrogen loves to interact
with things, though.
It interacts with lots
and lots of things.
And here hydrogen with
its one valence electron
is interacting with
chlorine with its seven
valence electrons, and they are
sharing two electrons forming
a bond together.
So when we're talking
about Lewis structures,
we're talking about
different kinds of electrons.
So we're talking about bonding
electrons, the electrons that
are involved in the bond,
and also lone pair electrons.
So chloride in HCl is going
to have two bonding electrons,
one was its, and one
came from hydrogen.
And it's also going to have
six lone pair electrons,
or we could say
three lone pairs.
So when we say a lone pair, that
indicates two electrons there.
So it has one, two, three,
four, five, six or one pair,
two pairs, three pairs.
Now there are rules
to Lewis structures,
and here is the complete rule.
In your handout, this
wouldn't fit on one page.
It's on two pages.
And these rules, if you
do work these problems,
you will remember these rules,
and they become pretty easy.
But it's important to work
Lewis structure problems
so that the rules become
really familiar to you.
And it takes time to work
Lewis structure problems,
so don't wait to the last minute
to start this problem set.
There's a lot of Lewis
structure problems on it,
which means it's not
difficult, but it's
going to take some time.
All right.
So let's briefly go
over these rules.
First what you want to do is
draw in the skeleton structure.
Just put the atoms down.
Hydrogen and fluorine are always
going to be terminal atoms.
Don't put them in the
middle of a molecule.
That gets chemistry
professors really upset
to see hydrogen in the
middle with lots of bonds
to things, so don't do it.
And typically the element with
the lowest ionization energy
goes in the middle, and
there are some exceptions
and we'll see some
of those exceptions.
But that should be
your first guess.
You want to count the
number of valence electrons.
If there's a negative charge,
you need to count that in
or if there's a positive charge
you need to subtract that
from the total.
Then you want to figure out
the total number of electrons
needed, so everyone has
their full valence shell.
You need to subtract these two
to get the number of bonding
electrons.
And here are some of
the things that it's
really easy to make
math mistakes here,
so if your structure makes
zero sense at the end,
go back and check your math.
Assign to bonding
electrons to each bond.
If any remain, you
want to think about
whether you have
double or triple bonds.
And there's only certain
kinds of atoms that can
have double and triple bonds.
So be careful where
you're putting
your double and triple bonds.
If any valence electrons
remain, those are loan pairs.
And then lastly, you want to
figure out the formal charge
on all of the atoms in
your structure to make sure
that this is a valid
structure, and we're going
to talk about formal charge.
So first, let's
just try an example.
And your sheet has two
examples on the same page.
We have HCN and we
also have CN minus.
So we're going to do HCN first,
so don't fill your entire page
because you're going
to have to write
things for CN minus as well.
But before we start, we need to
figure out which atom is likely
going to be in the middle.
And so why don't
you tell me what you
think on the clicker question.
OK, 10 more seconds.
Yup.
So here again we want to have
one that has a lower ionization
energy, and you also want
to consider other things
like hydrogen can't
be in the middle.
OK.
And it was written that
way, but sometimes they're
written in a way that is
not as straightforward.
OK.
So I'll put that up.
All right so we'll
go through the rules
and we'll try to work this out.
So first I can write--
I'm going to start--
I guess I'll write over here.
So number one, we're just
going to write HCN with C
in the middle.
So that's the first
thing we're going to do.
Next we're going to consider
the valence electrons.
And you can just help me
out by yelling things out.
How many valence electrons
does hydrogen have?
AUDIENCE: One.
CATHERINE DRENNAN:
What about carbon?
AUDIENCE: Four.
CATHERINE DRENNAN: Nitrogen?
AUDIENCE: Five.
CATHERINE DRENNAN: And you can
always check me on my math.
How much does that equal?
AUDIENCE: 10.
CATHERINE DRENNAN: Excellent.
There's nothing like
adding simple numbers
in front of 350 people to really
put the stress in one's day.
OK, so to have a complete
full valence shell, what
do I need for hydrogen?
AUDIENCE: Two.
CATHERINE DRENNAN:
What about carbon?
AUDIENCE: Four.
CATHERINE DRENNAN:
Eight to be complete.
Nitrogen?
AUDIENCE: Eight.
CATHERINE DRENNAN: Eight.
And I think we add this
up, we should get 18.
How's that?
All right.
So for four, now we're going
to subtract these numbers
from each other to tell us how
many bonding electrons we have.
So we have 18 minus 10.
And we should have eight
to bonding electrons.
For five I'm going to
now assign two per bond.
So I'm going to put one here.
Another here.
Another here.
Another here.
So I've assigned two per bond.
And now I see if
I have any left.
Do I have some left?
I've used four, I had eight.
So yes, I have four more.
And if you have more
bonding electrons,
then you are supposed to
assign those bonding electrons.
And think about
whether it's allowed
to have double or triple bonds.
Can hydrogen be involved
in the double bond?
No.
Carbon nitrogen?
AUDIENCE: Yes.
CATHERINE DRENNAN:
Yes, or a triple bond.
So I have four
more, so I'm going
to put one, two, three, four.
So I'm going to
have a triple bond
between my carbon and my
nitrogen. And now I'm good.
So now I want to see do I
have any extra electrons.
So for this, I had 10, I used
eight, so I have two left.
So I'm going to assign
those two as a lone pair.
Should I put them on hydrogen?
What about carbon?
No.
Carbon already has its eight,
because it has four bonds.
So then I'm going to
put it on nitrogen.
And then the only thing
left is formal charge, which
we haven't talked about yet.
So before we do
formal charge, I just
want to do the same
thing with CN minus.
And I will say if
you want to draw
triple bonds, that's fine too.
You don't have to indicate
the dots as bonds.
It's perfectly fine to
write out a triple bond.
So I could have
written this as well.
OK, so let's look at CN minus.
So how many valence electrons
does carbon have again?
Four.
Nitrogen?
Five.
Am I done?
AUDIENCE: No.
CATHERINE DRENNAN: No.
I need to add one
because there's
a charge on this
molecule of minus one.
So now so I have 10 again.
Three I'm going to
figure out how many
electrons I need to
complete my valence shell.
How many from carbon?
AUDIENCE: Eight.
CATHERINE DRENNAN: Eight.
Nitrogen?
Eight so now I have 16.
I will subtract.
Now 10 from 16 and get
six bonding electrons.
And I'm going to assign
first just two of them.
So I'll assign one, two here.
And then we-- this is to assign.
Six said, do you
have any left over?
I had six, and I only used
two, so the answer is yes.
I have four more.
So I can put those in
one, two, three, four.
Again, we're going to
have a triple bond.
And then we ask are
there any electrons left?
So we had 10, we
used six of them.
And so we're going to
have now four more here.
So now I can assign--
this only has-- this
is not complete for carbon.
So I can put a lone
pair on carbon,
and I can put a lone
pair on nitrogen.
And now they have
their complete octet.
And we get to assign
formal charge.
So I could write it this way,
or I could have written it
with a triple bond here.
And don't forget to put
the charge on the end.
All right.
So now let's consider
formal charge,
because we're never done
with our Lewis structures
until we've considered
formal charge.
So formal charge is a
measure of the extent
to which the atom has
really lost or gained
an electron in the process of
forming this covalent bond.
So as we'll talk about, it's
not the same as oxidation number
where you have like
sodium plus one,
but again, there's
some differences
in how many are
brought to the table
and what it ends
up with in the end.
So there's an equation
which you'll have to learn,
but if you do enough
of these problems,
it'll be stuck in your
brain and you can't purge it
even if you want to.
And so formal charge, FC,
is equal to the number
of valence electrons,
symbol V, here,
minus the number of
lone pair electrons
minus half of the number
of bonding electrons.
So at least this
equation makes sense.
If you forget what they mean,
you can probably think about it
and it'll come back to you.
So in doing these
formal charges,
you want the formal
charges to add up
to the charge on the molecule.
So if we had HCN, that's
a neutral molecule
so the sum of all of the
formal charges must be zero
or you did something wrong.
If it's CN minus, the
sum of the formal charges
has to add up to minus one
or you did something wrong.
So this is a good way
of checking your math.
So always remember
that the sum needs
to add up to the total
charge on the molecule.
If you remember that, that's a
really good check to make sure
you didn't make some kind
of weird math mistake
and add things wrong and have
an appropriate number of loan
pairs or something going on.
OK.
So let's calculate
formal charge now
on our CN minus
molecule up here.
So the formal charge
now on carbon here.
So how many valence
electrons do carbon have?
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN: Four.
How many lone
pairs does it have?
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN: It has
one lone pair, two lone pair
electrons.
This is the total number of
lone pair electrons here,
and then half the number
of bonding electrons.
So it is expanding electrons
and half of that is three.
And so that should add up
to a charge of minus one.
You can also, instead of
thinking half the number
of bonding electrons,
you can also
just think about number of
bonds if you want to to do this.
All right so to see if
you have the hang of it,
let's do a clicker question.
OK, 10 more seconds.
Most people got it right.
I can't actually read the
number, but that was very good.
It always seems wrong to
put zero as the answer,
but that is, in fact,
the answer here.
So if we look at this again,
we have five valence electrons
on nitrogen. We have
two lone pair electrons,
and we have six
bonding electrons.
Half of six is three,
and so that's zero.
Or you could have said,
well, if this is minus one,
and the charge is
minus one, then
that had to have been zero,
otherwise Professor Drennan
did something wrong, and
that's just not possible.
So the answer there
would be zero.
And you can see that the total
formal charge is minus one,
and the total charge of the
molecules also minus one.
So again formal charge does
not equal oxidation number,
it's something special.
It tells you kind of in
this arrangement of atoms
in the molecule,
did this atom kind
of come up with a little
bit more at the end
or a little bit less, depending
on where it started from
and where it is.
And where you want
in these structures
for the formal charge to be low.
So we can use a formal charge to
decide between Lewis structures
so that the structures
with the lowest
absolute values of
this formal charge
are more stable structures.
So if you have really
high formal charges,
that means that molecule
isn't really very stable
because you want low charges.
You want lower energy.
You want things to be in a
more neutral and happy state.
So we want to figure
out which ones
are going to have low charges.
So let's look at another
example, Thiocynate ion,
and it has a carbon,
sulfur, and nitrogen in it,
and it has a charge
of minus one.
So I might tell you the
ionization energies for carbon,
sulfur, and nitrogen,
and then ask you
which atom do you
think is going to be
in the center of the molecule?
So what do you think?
What's in the center
of this molecule based
on those numbers?
Just yell it out.
AUDIENCE: [INAUDIBLE]
CATHERINE DRENNAN: Sulfur.
So sulfur has the lowest
ionization energy,
and I told you that's usually
the thing in the center.
But you can start with that.
It's always good
to start with that,
but then you want to
check the structure
and make sure that a structure
with sulfur in the middle
has the lowest formal charge.
So let's take a look at that.
So we can draw structure A
with sulfur in the middle,
and then calculate the
formal charges on that.
And if we do that, we see
for the nitrogen here,
nitrogen has five
valence electrons,
it has four lone pair electrons,
and it has half of four bonding
electrons, so it would have
a formal charge a minus one
in this particular structure
that has sulfur in the middle.
We can look at carbon.
Carbon has four valence
electrons, Four lone pair
electrons, and half of
four bonding electrons,
so it has a charge of minus two.
Then we can look at sulfur.
Sulfur has six valence
electrons, zero loan pairs,
and half of eight
bonding electrons,
so it has a formal
charge of plus two.
So overall, this does
equal the minus one.
So it's a valid structure.
But is that the
lowest energy one?
We also could put carbon
in the middle or nitrogen
in the middle.
So let's look at what
this does for us.
So with the formal
charge on nitrogen now,
we have five minus four lone
pair electrons, minus half
of four bonding
electrons, minus one.
So that's the same.
Now we can look at carbon.
We have now just no
lone pair electrons.
It has four minus
zero minus four,
half of eight bonding
electrons, or zero.
And for sulfur, six minus
four lone pair electrons,
and half of four bonding
electrons, or zero.
Next, structure C, five minus
zero lone pairs minus half
of eight, plus one.
So that one's different.
Carbon, we have four
valence electrons
minus four lone pairs minus
half of four, bonding minus two.
And for the sulfur, six
minus four lone pairs
electrons, half
of three or zero.
So now with the clicker,
tell me which is most stable.
All right, 10 seconds.
I think this can be 98%.
That's what I'm thinking.
I'm feeling good.
Well, close.
Yeah.
What?
No, no.
Sorry.
It should be B. Yeah,
it should be B. Yeah.
Sorry I actually-- Yay.
There we go.
Yeah, B is correct.
So if we just look
at it over here,
it has the lowest number
of formal charges,
so the answer is B. OK.
Let's start with this simply.
Who wants to tell me why one,
how they could look at this
and realize one was
not the correct answer?
I think this is on.
Give it a try.
AUDIENCE: One's not correct
because if you look back
at your atomic radius chart,
this is pretty much doing
the exact opposite of that
CATHERINE DRENNAN: Yeah,
so helium definitely
not the biggest atom there is.
OK so six got a lot of
attention, and so did two.
And ionization energy, electron
affinity, and electronegativity
are definitely
connected to each other,
but there is a clue that
electron affinity would not
be the correct.
Does someone want to say
what you might have noticed?
AUDIENCE: For electron
affinity, it increases and then
stops at the noble gases
because noble gases do not
want electrons.
So in this particular
chart, all the noble gases
are like the highest--
are the highest ones
in the relative
area, which would
mean that electron affinity
would be incorrect.
CATHERINE DRENNAN: Yep.
That's right so the noble
gases were the clue.
So [INAUDIBLE].
Yep.
And so if electron affinity also
is not high at the noble gases,
they're also not
electronegative.
Noble gases just don't
want extra light.
They don't want
to lose electrons,
they don't want
to gain electrons,
they just want to be left alone.
So this trend is for
ionization energy.
And because noble gases
want to be left alone,
they don't want to lose
any of their electrons.
Great.
So this is good to
be thinking about,
because we're finishing up now
for the handout from last time.
And we're going to be talking
about electronegativity again.
We never move very far away
from a lot of these topics.
They just keep coming back.
So we just keep reviewing them.
All right, I don't
need to microphones,
although, I don't know, I
kind of like having this one.
Anyway.
So let's take out the
handout from last time,
and let's finish it up.
We were talking
about formal charge,
and we had looked at
examples where we calculated
formal charge,
and then we looked
at which structure would
be lowest in energy,
and that was the structure
where you had the least
separation between
the charges on them,
so the smallest
absolute numbers.
If you have formal charges
of zero, that's fantastic.
That's what the molecule wants.
Minus one, plus
one, if you must,
but when you start having
plus two minus two, that's
a lot of charge separation,
so that's less favorable.
So we're going to have
more examples of that
as we go along.
But now, what-- if you had
calculated formal charge
and they're all the
same, how do you
know which structure is correct?
So what if you have-- and
this is-- just some people who
are having trouble.
This is the top of page
four from the handout.
You have two valid
Lewis structures
that have the same
formal charge,
how do you know where it goes?
And the answer is that
the negative formal charge
should go on the most
electronegative atom.
And so that's why we are sort of
talking about electronegativity
again.
And so electronegativity--
remember electronegativity
is high when the electron
affinity is high,
meaning that the atom
wants to get an electron,
has a high affinity
for electrons,
and also a high ionization
energy, which means it
doesn't want to give
up its electrons.
So that's something-- it
likes to have electrons,
and so you want to put
a negative value, which
indicates there's more
electrons on something
that's electronegative.
So negative on
negative over here.
And so that's what
you're looking for.
That's how you're going
to make a decision.
So let's look at an example.
So here is a molecule,
and we're going
to look at two possible
Lewis structures of this
with similar formal
charges, and decide which
has the correct structure.
So first let me give you a
couple hints that can be useful
in problem sets and in exams.
When you see CH3,
that's a methyl group,
and that's going to
be terminal so you're
going to have these three
hydrogens associated
with that carbon, and
that's going to be
at an end of the molecule.
So it could look like
one of these two things.
So we have the carbon,
we have it attached
to three hydrogens, a carbon
attached to three hydrogens,
and then attached to something
else, this nitrogen here.
And this structure where
it's just kind of written out
in a line or a chain
of atoms, what we
call chain molecules sometimes.
Often the atoms are
actually written
in the order in which
they're attached,
so that's definitely true here.
CH3, three hydrogens
attached to the carbon,
so they're attached to the atom
that came before in the chain.
The nitrogen is also going
to be attached to the carbon.
Even though it
follows the hydrogen,
you're not going
to have hydrogen
in the middle of a bond.
It's not going to be
bonded to two things.
Hydrogen is always terminal, so
even though nitrogen is here,
it's got to be
attached to the carbon.
So we have three
hydrogens and then
a bond between the
carbon and the nitrogen.
Now we have a hydrogen
after the nitrogen,
and by this rule it
should go on the nitrogen.
But you might want to double
check that that's true.
And then we have
an oxygen. Again,
the oxygen is going to have
a bond with the nitrogen.
You're not going to have a bond
with hydrogen in the middle.
Hydrogen's always terminal.
So the only real
choice we have here
is we can put the
hydrogen on nitrogen
or we can put the hydrogen
on the oxygen here.
And so we can use this rule
about electronegativity and
formal charges to figure out
which of these structures
is right.
So in this particular case,
all of the formal charges
are zero on all of the
atoms, except there's
one minus one charge.
And in this structure,
the minus one charge
would be on the oxygen,
and in this structure
the minus one charge
would be on the nitrogen.
And if everything
else is zero then you
have the sum of
your formal charges,
minus one, equal to the
charge on the molecule,
which is minus one.
So both of these are
valid structures.
Both have low values of
formal charge, which is right.
So it's going to be
the structure that
has the negative charge on the
most electronegative atom is
the right structure.
And so here you need to
remember some of your rules
about electronegativity.
And in terms of
electronegativity,
we see that oxygen has
greater electronegativity
than the nitrogen.
And so that's where
we would want to put
our negative charge--
on the negative charge
goes on the atom that's
the most electronegative.
And that would generate
the lower energy structure.
So if you're given a table
of electronegativities, which
you often are,
you can look it up
and validate that that's
going to be the correct place.
And in fact, experimentally
we know that that's
the right structure.
So that works.
So if you have two structures,
identical formal charges,
valid structures,
then the last step
is to think about where
that negative charge should
go, and pick the atom that
is the most electronegative.
All right we have
one more thing we
need to talk about
in Lewis structures
before we start
violating various rules
that we've learned,
and that is that we
need to talk about
resonance structures.
And so to explain to you what
a resonance structure is,
it's really helpful to
start with an example,
so that's what
we're going to do.
And we're going
to consider ozone,
which is three atoms of oxygen.
And we have the ozone layer,
which protects us from UV
damage and is very valuable,
and we should not destroy
it with chemicals being
released into the environment.
And because you don't have
complete say over that,
always wear sunscreen.
OK.
So let's build up
these Lewis structures,
and then consider what's meant
by a resonance structure.
So here we have part
of the Periodic Table
that you're going
to need, and we
need to think first about
the valence electrons.
So oxygen has six
valence electrons
and there are three
oxygens, so that's 18.
To get a full octet for each
of the oxygens, three oxygens,
an octet is eight valence
electrons, so that would be 24.
To figure out the number
of bonding electrons,
we're going to be subtracting
our octet electrons
from our valence electrons.
So 24 minus 18 is six.
And then our next step is
to assign those bonding
electrons two at a
time, two per bond.
So let's take a look at that.
We can put one bond here
between these two oxygens,
one bond here between
these two oxygens,
and then ask do
we have any more?
And we do.
Because we have six bonding
electrons, we used four,
so we have two more
bonding electrons.
So we need to make
a double bond.
But now we have the question of
where to make that double bond.
Am I going to put it
between these two oxygens
or am I going to put it
between those two oxygens?
And so I could say put it
there, but I could also,
in structure two, put the
double bond over here.
All right.
We'll come back to that
question in a minute.
Let's first figure out if we
have any remaining valence
electrons, and we do.
So we had 18 and we've only
used six, so we have 12 left.
So we're going to put
those in as lone pairs.
And so I can put
them in over here.
One set here, one, two,
three, four, five, six.
And I can also do
that over here.
One, two, three,
four, five, six.
And now we have two
structures, so we
need to think about
formal charges
to see if that can help
differentiate structure
one from structure two.
OK.
And let's look at that.
Be sure everyone's ready.
And that is a clicker question.
So I'll put that back up.
All right.
Let's just take 10
more seconds, and we'll
talk about what the right answer
is, and a little bit of a trick
for doing these,
perhaps, a bit faster.
OK.
So that was pretty good.
So let's look at why
that's the right answer.
And we'll take a look
at that over here.
So let's do the calculations.
So you have to remember the
equation for formal charge
for sure and once you
do enough problems
it should stick in your
head pretty easily.
So if we look over
here-- so this
is the formal
charge on oxygen A.
There were six
valence electrons,
there are four lone pair
electrons, minus half
of the bonding electrons.
There are eight bonding
electrons, so that's two.
So that's a formal
charge of zero.
For oxygen B over here, we again
have six valence electrons,
and we have two
lone pair electrons.
We have six bonding
electrons, so half of six
is three, which is plus one.
And on oxygen C over
here, six minus six
lone parallel electrons, half
of one bond, so half of two
is one, minus one, overall
the charge here is neutral.
And that's good because
it's a neutral molecule.
So to do this
structure faster, you
have to realize that
oxygen A over here
is the same as C over there.
So you can just
copy down what you
had for C is now A. B is exactly
the same in the two structures,
so it's the same as
what you calculated.
And now this C was
the same as this A,
so we can put that
same value that we
calculated for A into C.
So they're both the same.
Same formal charges, which
structure is correct?
And the answer is both of them.
And in fact, you
need both of them,
so there's data--
chemists love data--
and the data is that
the bonds are actually
equivalent in the molecule.
So it isn't that there's
one double bond and one
single bond, there's just
one kind of bond in ozone,
and it's between a
single and a double bond.
And so this is how one would
draw that kind of thing.
You would have structure one
here and structure two here.
You would put them
both in brackets,
and you would put this
special double headed
arrow between them
and that indicates
that both of these structures
are needed to describe
the properties of the molecule.
You do not have a stationary
double bond and a single bond.
You have kind of a
mixture between these two,
and that's what a
resonance structure is.
So let's take a little bit
more of a look at this.
So this is just what I had
before, experimental evidence
is that the bonds
are equivalent.
There isn't a
double and a single,
there's something in
between, a one and a half.
And this is called
a resonance hybrid,
and it's of length of
these two structures.
So this structure is
blended with that structure.
So some of you are aware
of hybrids from biology,
and now with cars.
So a mule would be an example.
And if you're thinking
about what a mule is,
you don't walk out into
your barnyard one day
and see a donkey one day
and you go out the next day
and you see a horse.
A mule is a hybrid between
a donkey and a horse.
So if you were a chemist,
you would do this.
You would put the
donkey in brackets,
and the horse in brackets, and
put your double headed arrow.
And if you see this,
you'd say, oh yes, a mule.
So that is what this is.
Both structures are
needed to describe ozone.
One structure isn't enough.
You need both of them.
They're in resonance
with each other.
And so what's true about
the electrons in ozone
is that they're delocalized
across all of these bonds,
so there isn't like
a double single.
All of those electrons
are delocalized.
They're shared over
this set of atoms.
And you can have two resonance
structures, you can have three,
you can have four.
It depends on the
molecule in question.
And in all those cases, those
electrons would be delocalized.
So just to sum that up,
resonance structures,
two or more, same arrangement
of atoms-- and that's important.
It's the electrons
that are different.
And this isn't in your
handout, out but just think
about this for a second--
because it was in your hand
out a minute ago.
Are these resonance structures?
No.
They're not
resonance structures.
The atoms are in
different positions.
So one of these structures is
right one of these structures
is wrong.
With resonance structures,
they're both correct
and both needed to
define the structure.
So pay attention.
This is a common mistake.
Pay attention.
Ask yourself, are the atoms
in a different position?
That's not resonance.
You're just looking-- atoms
are the same, formal charge
are the same,
you're just looking
at whether you have different
arrangements of electrons.
