So we have two different molecules here.
This is hydrogen peroxide.
We call it peroxide, because it has this oxygen-oxygen bond.
And here we have oxygen difluoride,
where oxygen is bonded to two different fluorines.
And what I want you to do is pause this video, use
this periodic table of elements I have here,
and this is more than just a typical information
of periodic table of elements.
It also gives you the electronegativities
of these different elements.
And these electronegativities are
based on the Pauling scale named after famous biologist
and chemist Linus Pauling.
And so using the information here
and what you know already about oxidation states,
think about the oxidation states or the oxidation numbers
for each of the constituent elements in these molecules.
So pause the video now.
So I'm assuming you have given a shot at it.
And you might have immediately realized
that something very interesting is going on.
We've said in the past that because it's
two valence electrons away from a full valence shell,
because it is so electronegative,
oxygen typically takes electrons from other things, typically
two electrons, which typically gives it
an oxidation state of negative-- an oxidation number
or oxidation state of negative 2.
This is so electronegative, and it so typically oxidizes
other things that we've called the whole phenomenon
"oxidation."
But what's interesting here is that oxygen
isn't purely bonded to things less electronegative
than itself.
And the hydrogen peroxide, yes, it is bonded to the hydrogen.
But it's also bonded to another oxygen.
And obviously, these two are going
to be equally electronegative.
So what would be the oxidation states
or the oxidation numbers here?
Well, hydrogen, once again, we portend-- hydrogen,
because it's less electronegative,
it would have a partially positive charge,
because the electrons would spend
more time around this oxygen.
But when we're talking about oxidation states,
we don't like this partial charge business.
We want to pretend like these covalent bonds
are ionic bonds, hypothetical ionic bonds.
And if they were hypothetically ionic bonds, what would happen?
Well, if you had to give these electrons to somebody,
you would give them to the oxygen,
the electrons in this period, give them to the oxygen,
giving it an oxidation state of negative 1.
With the hydrogen having these electrons taken away,
it's going to have an oxidation state of positive 1.
And the same thing's going to be true for that oxygen
and that hydrogen right over there.
So this is fascinating, because this
is an example where oxygen has an oxidation state not
of negative 2, but an oxidation state of negative 1.
So this is already kind of interesting.
Now it gets even more interesting
when we go to oxygen difluoride.
Why is this more interesting?
Because fluorine is the one thing on this entire table
that is more electronegative than oxygen.
This is a covalent bond, but in our hypothetical ionic bond,
if we had to give these electrons to one
of these atoms, you would give it to the fluorine.
So the fluorine, each of them would
have an oxidation state of negative 1.
And the oxygen here-- now, you could imagine,
this is nuts for oxygen.
The oxidation state for oxygen, it's giving up these electrons.
It would be a positive 2.
And we talk about oxidation states
when we write this little superscript here.
We write the sign after the number.
And that's just the convention.
But it has an oxidation state of positive 2.
Oxygen, the thing that likes to oxidize other things,
it itself has been oxidized by fluorine.
So this is a pretty dramatic example
of how something might stray from what's typical oxidation
state or it's typical oxidation number.
And in general, oxygen will have an oxidation state or oxidation
number in most molecules of negative 2.
But unless it's bonded with another oxygen
or it's bonded to fluorine, which is a much more
electronegative-- or actually, not much more,
but it's the only atom that is more electronegative than--
or the only element is more electronegative than oxygen.