Unusual oxygen oxidation states

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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.