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.