Resonance and Dot structures - Chemical Bonding Video Lecture | Inorganic Chemistry

Now that we know how
to draw dot structures,
let's apply our rules
to the nitrate anion.
And we're going to see that we
can draw a few different dot
structures for this anion.
And we're going to
call those resonance
structures of each other.
But first, we need to
calculate the total number
of valence electrons.
And so nitrogen is in Group
5 in the period table,
therefore, five
valence electrons.
Oxygen is in Group 6,
therefore, six valence electrons
for each oxygen.
I have three of them.
So 6 times 3 is 18 valence
electrons, plus the 5
from the nitrogen gives me 23.
And I have a negative charge.
This is an anion here.
So we have to add
one electron to that.
So 23 plus 1 gives us a
total of 24 valence electrons
that we need to represent
in our dot structure.
So we know that nitrogen is
going to go in the center,
because oxygen is
more electronegative.
So nitrogen goes in the center.
Nitrogen is bonded
to three oxygens.
So I can go ahead and put
them in there like that.
And let's see.
How many valence electrons
have we represented so far?
2, 4, and 6.
Therefore, 24 minus 6 gives us
18 valence electrons left over.
We're going to put those
leftover valence electrons
on our terminal atoms,
which are our oxygens.
And oxygen's going to
follow the octet role.
Currently, each oxygen has two
valence electrons around it,
the ones in magenta.
So if each oxygen
has two, each oxygen
needs six more to
complete the octet.
And so I go ahead and put
six more valence electrons
on each one of my oxygens.
Now each oxygen is surrounded
by eight electrons.
So the oxygens are happy.
We added a total of six valence
electrons to three oxygens.
So 6 times 3 is 18.
So we've used up
all of the electrons
that we need to represent.
And so this dot
structure, so far,
it has all of our
valence electrons here.
Oxygen has an octet.
So oxygen is happy.
But nitrogen does
not have an octet.
If you look at the
electrons in magenta,
there are only six electrons
around the nitrogen.
And so the nitrogen
wants to get to an octet.
And there are a couple
of different ways
that we could give
nitrogen an octet.
For example, we could take
a lone pair of electrons
from this top oxygen here
and move them into here
to share those electrons
between that top oxygen
and that nitrogen.
So let's go ahead and draw
that resulting dot structure.
So we would have
our nitrogen now
with a double bond
to our top oxygen.
Our top oxygen had three
lone pairs of electrons.
But now it has only two,
because electrons in green
moved in to form a double bond.
This nitrogen is
bonded to an oxygen
on the bottom left and an
oxygen on the bottom right here.
So this is a valid
dot structure.
We followed our steps.
And we'll go ahead and
put this in brackets
and put a negative charge
outside of our brackets
like that.
So that's one possible
dot structure.
But we didn't have to take
a lone pair of electrons
from the top oxygen.
We could've taken a lone pair
of electrons from the oxygen
on the bottom left here.
So if those electrons
in blue moved in here,
we could have drawn
another dot structure which
would have been equally valid.
We could have shown this oxygen
on the bottom left now bonded
to this nitrogen, and it used
to have three lone pairs.
Now it has only two.
And now this top oxygen
is still a single bond
with three lone pairs around it.
And this bottom
right oxygen is still
a single bond with three
lone pairs around it.
So this is a valid
dot structure as well.
So let's go ahead
and put our brackets
with a negative charge.
And then, of course,
we could have
taken a lone pair of
electrons from the oxygen
on the bottom right.
So I could have moved these
in here to form a double bond.
And so now, we would have
our nitrogen double bonded
to an oxygen on
the bottom right.
The oxygen on the
bottom right now
has only two lone
pairs of electrons.
The oxygen at the top, single
bond with three lone pairs.
And then the same situation for
this oxygen on the bottom left.
And so this is, once again,
another possible dot structure.
And so these are
considered to be
resonance structures
of each other.
And the way to
represent that would
be this double-headed
resonance arrow here.
And I think when students
first see resonance structures,
the name implies
that, in this case,
the ion is resonating back
and forth between these three
different possible, equally
valid dot structures.
And that's not quite
what's going on here.
Each of these dot
structures is an attempt
to represent the
structure of the ion.
But they're really not the
best way of doing that.
You need to think about
combining these three dot
structures in a resonance
hybrid of each other.
And so let's go ahead and draw
just a simple representation
of a way of thinking
about a resonance hybrid.
So if I combined all
three of my dot structures
here into one picture,
I had a double bond
to one oxygen in each of my
three resonance structures
here.
And so the top oxygen had a
double bond in one of them,
the bottom left in the middle
one, and then the bottom right
in the third one.
So, in reality, if we take a
hybrid of all those things,
we could think about the
electrons being delocalized
or spread out among all
three of our oxygens.
And so instead of giving
our top nitrogen-oxygen,
instead of making
that a double bond,
we can just show some electrons
being delocalized in that area,
so stronger than a single
bond, but not as strong
as a double bond.
And we could do the same
thing between this nitrogen
and this oxygen.
So the electrons are
delocalized a little bit here.
It's not a double bond.
It's not a single bond.
And the same idea for this
nitrogen-oxygen in here.
And one way we know that the
ion looks more like this hybrid
is because of bond length.
When the ion is measured in
terms of the bond length,
all the nitrogen and oxygen
bonds are the same length.
And of course, if
we thought about one
of these resonance
structures as being
the true picture
of the ion-- let's
say this one, for
example-- that wouldn't
be the case for this ion,
because this double bond here,
we know that would
be shorter than one
of these single
nitrogen-oxygen bonds.
And so it's actually
more of a hybrid
with the electrons
delocalized throughout.
And that's the idea of
resonance structures here.