Now I want to show you how to fill in electrons
on a molecular orbital energy diagram, and
we’re going to use O2 as an example.
The first thing we want to do is we want to
put electrons on each side here for the Oxygen’s
that are not bonded together, and then what
we’ll do is we’ll combine those electrons
into the center to show what happens when
the Oxygen’s get bonded together.
All right so let’s look at Oxygen first,
we know that it’s got 8 electrons because
it’s got an atomic number of 8, which means
it has 8 protons.
So we’ve got 8 electrons, we’ll put 2
in here, and we’ll put 2 in here, just according
to the principle, and then obeying Hund’s
Rule, put 1 in here, 1 in here, 1 in here, that’s 7 so far.
So now we can put 1 in here, and that makes
up our 8.
Let me do a better job of that.
There we go.
So now we’ve got 8 electrons put in on that
side, let’s do the same on this side, 1,
2, 3, 4, 5, 6, 7 and then 8.
It doesn’t matter which orbital that the
other electron goes into.
All right so now we’ve got 8 on this side,
and 8 on this side, that means we’ve got
a total of 16 electrons, that we want to deal
with.
So what I’m going to do is I’m going to
fill that in.
It’ll go 1, 2, 3, 4 and we’re filling
this in with the same rule, we’re starting
from the bottom and moving to the top.
We’ve got energy on this side over here,
so that’s 4, 5, 6, 7, 8, 9 and 10, now here
we use Hund’s Rule, and we put 1 electron
in here, 1 electron in here, that’s 12, 13 and then 14.
We’ve still got 2 more to put in.
That’s 15 and 16 up there.
All right so now what you can see is we’ve
got our molecular orbital diagram filled in
and as you’d expect you can see that the
1s’s don’t combine.
You see, you’ve got the 1s here, and the
1s here, and we got 2 in the anti-bonding
and 2 in the bonding.
And what that just tells us is that the 1s
orbitals don’t bond.
And that’s what we expect and the reason
is because that’s the inner electron’s
of Oxygen, the inner electrons aren’t expected
to bond.
All the bonding we expect to see will be in
the valence electrons, and you can see those
are the ones up here.
All right, so let’s take a look at how we
calculate bond order.
And the equation I showed you earlier was
bonding electrons minus anti-bonding divided by 2.
So what this is going to tell us is how many
bonds we expect to see between Oxygen, I know
I’ve only put 1 here but we’ll see in
a second that that’s not really just 1 bond between the 2 Oxygen’s.
So here’s what we do, we have to calculate
how many bonding electrons there are, now
the way we do that is any orbital that doesn’t
have a star on it is considered to be a bonding
orbital, anything that has a star on it is
considered to be anti-bonding.
So these don’t have stars, this one, so
that’s 2 electrons there, there’s 2 electrons
there, that’s a total of 4, we got 2 electrons
here, that’s 6, 2 here, that’s 8, and 2 here, that’s 10.
So we have a total of 10 electrons that are
in bonding orbitals, and the anti-bonding,
we’ve got 2, and 2 is 4, and 2 is 6.
Ten minus six divided by two equals two.
All right so that’s what we expect to see
in Oxygen, is a double bond, or 2 bonds between
the Oxygen’s, and that’s exactly what
we would expect to see, if we take the lewis
structure of Oxygen, we’ve got 6, 12 valence
electrons in total, now not total electrons,
but valence electrons, these are just electrons
in the outer shell.
If we look at what that would do, we end up
getting 2 in the middle from each of them,
for a double bond, and that’s how the Lewis
Structure would look, and the other way of
showing that is just as a double bond here,
and that’s the only way that we can arrange
it, so that each of those Oxygen’s end up
having 8 electrons, they have to share a double
bond, and that would be what’s predicted
by the bond order as well.
