Let's figure out the oxidation
states for some more

constituent atoms
and molecules.

So let's say I had
magnesium oxide.

MgO.

I'll do oxygen in a
different color.

So what are going to be their
oxidation states?

And you might know this already,
but let's look at the

periodic table, because
it never hurts to get

familiar with it.

So we have magnesium.

Magnesium has two valance
electrons.

It's a Group 2 element.

It would love to lose
those two electrons.

Oxygen, we already know,
is one of the most

electronegative atoms. It's so
electronegative that oxidized

has essentially been
named after them.

And we know that oxygen loves
to gain two electrons.

So this is kind of a marriage
made in heaven.

This guy wants to lose two
electrons and this guy wants

to gain two electrons.

So what's going to happen?

The magnesium is going to
lose two electrons.

It was neutral.

So it's going to have a plus
2 charge, hypothetically.

And then, the oxygen is going
to have a minus 2 charge,

because it gained the
two electrons.

So in this molecule of magnesium
oxide, the oxidation

state of magnesium is plus 2.

And the oxidation state of
the oxygen is minus 2.

Now let's do a slightly
harder one.

Let's say we had magnesium
hydroxide.

So hydroxide is OH2.

OH2 right there, where there's
two hydroxide groups in this.

So, my temptation would
still be, look.

Magnesium likes to lose its
electrons, its two electrons,

which would make it's charge
positive-- it's hypothetical

oxidation state positive.

So my temptation is to say,
hey, magnesium here

would be plus 2.

So let me write that there.

And remember, in order for
everything to work out, if

it's a neutral compound, all of
the oxidation states in it

have to add up to 1.

So let's see if that's
going to work out.

Now, oxygen.

My impulse is that its
oxidation state

tends to be minus 2.

So let me write that down.

And hydrogen, when it's bonded
with an oxygen-- remember.

In this case, the hydrogen is
bonded with the oxygen first.

And then that's bonded
to the magnesium.

So the hydrogen is bonded
to an oxygen.

Hydrogen, if it was bonded to a
magnesium, you might want to

say, hey, maybe it'll take the
electrons and it'll have a

negative oxidation state.

But when hydrogen is bonded
with oxygen, it

gives up the electrons.

It only has one electron
to give up.

So it has a plus 1
oxidation state.

So let's see.

At first, you might say,
hey, I'm adding up

the oxidation states.

Plus 2 minus 2 is 0 plus 1.

I get a plus 1 oxidation
state here.

That doesn't make sense, Sal.

This is a neutral compound.

And what you to remember is oh,
no, but you have two of

these hydroxides right here.

So what you do is you figure out
the sum of the oxidation

states of the hydroxide.

So that's minus 2 plus 1.

So for the entire hydroxide
molecule, you

have a minus 1 sum.

And then you have two of them.

Right?

You have two hydroxide
molecules here.

So the contribution to the
entire compound's oxidation

state will be minus 1
for each hydroxide.

But then you have two of them.

So it's minus 2 and then plus
2 from the magnesium.

And it all adds up to 0.

So that worked out.

Now, I want to do a little
bit of an aside.

I want to go back to doing
some problems again.

But I want to do a little bit
of an aside on some of my

terminology.

Because I've kind of used
oxidation state, and oxidized,

and reduced interchangeably,
to a certain degree.

But, we've done so many
problems with water

autoionizing into-- actually,
let me do 2 moles of water.

And it's in equilibrium with 1
mole of H30 plus OH minus.

And obviously, everything is
in an aqueous solution.

Now, let's look at the water.

What are the oxidation states
in this water right here?

Well, we've done this already
in the previous video.

Oxidation state of oxygen is
minus 2, because it's hogging

the two electrons from
the two hydrogens.

Each hydrogen is giving
up an electron.

So it has an oxidation
state of plus 1.

And we see this molecule.

Everything adds up.

Because you have two hydrogens
with a plus 1.

So that's plus 2.

Plus 2 minus 2 for that one
oxygen, and you get to 0.

And it's a neutral compound.

Now here, what are the
oxidation states?

So one of these hydrogens left
one of these water molecules

and joined the other of the
water molecules without taking

its electron with it.

So it left the electron
over here.

So that oxygen still has a
minus 2 oxidation state.

And this hydrogen still
has a plus 1.

And that's why you do
minus 2 plus 1.

You get minus 1.

And this time, it works out,
because that's the actual

charge on this hydroxide ion.

Now, here, what are the
oxidation states?

Each of the hydrogens have
a plus 1 oxidation state.

And then this oxygen
has a minus 2.

And so if you look at the
charge for the entire

molecule, plus 1 on three
hydrogens, so that's plus 3.

I just added them up.

Minus 2.

So plus 3 minus 2, I should have
a plus 1 charge on this

entire molecule, which
is the case.

Now, my question to you is has
any of the oxidation states

changed for any of the atoms?

All of the hydrogens here--
and we could call

this 2 moles of water.

Or maybe I just have two
molecules of water.

But I have four hydrogens
here.

Right?

And all of them had an
oxidation state of 1.

On the right-hand side,
I have four hydrogens.

All of them have an oxidation
state of 1.

So although their oxidation
state is 1, in this reaction--

and you can pick either
direction of the reaction--

hydrogen has not
been oxidized.

Its oxidation state
did not change.

Maybe it was oxidized in a
previous reaction where the

water was formed, but in this
reaction, it was not oxidized.

Likewise, the oxygens-- we have
two oxygen molecules, or

atoms, here.

Each have a minus 2
oxidation state.

Here, we have two oxygen
molecules.

Each have a minus 2
oxidation state.

Due to this reaction, at least,
no electrons changed

hands in our oxidation
state world.

So this is not an oxidation
or a reduction reaction.

And I'm going to cover that in
detail in the next video.

And I just want to be clear that
nothing here was oxidized

or reduced, because
their oxidation

states stayed the same.

Because sometimes I'll
say, hey, look.

Magnesium has an oxidation
state of plus 2.

And oxygen has an oxidation
state of minus 2.

Magnesium was oxidized.

Two electrons were taken
away from it.

And oxygen was reduced.

Two electrons were
given to it.

And I'll say that implying some
reaction that produced

it, but that's not
always the case.

You could have a reaction
where that

necessarily didn't happen.

But the oxidation state
for magnesium is

definitely plus 2.

And the oxidation state for the
oxygen, or the oxidation

number, is minus 2.

But I think you know what I'm
talking about when I say it

was oxidized.

At some point, it went from
a neutral magnesium to a

positively charged magnesium
by losing two electrons.

So it got oxidized.

Now, let's do some harder
problems. So hydrogen

peroxide-- I've said multiple
times already that oxygen

tends to have a minus
2 oxidation state.

This is minus 1.

I think you see the pattern.

These guys are plus 1.

Hydrogen is plus or minus 1.

These guys are plus 2.

I think you see the pattern.

It's whether you want to
lose or gain electrons.

You might say, well see,
water normally

has a minus 2 oxidation.

So you might be tempted
to do-- OK.

Hydrogen has plus 1, because
it's bonding with water.

And oxygen has a minus 2.

But when you do that, you
immediately have a conundrum.

This is a neutral molecule--
let's see.

Two hydrogens plus 2.

Two oxygens at minus 2.

Minus 4.

So this would end up with
a minus 4 total

net oxidation state.

And that's not the
case because this

doesn't have any charge.

So there's a conundrum here.

And the conundrum is because,
if you actually look at the

structure of hydrogen peroxide,
the oxygens are

actually bonded to each other.

That's where the peroxide
comes from.

And then each of those are
bonded to a hydrogen.

So in this case, especially in a
first-year chemistry course,

the peroxide molecules,
especially hydrogen peroxide,

tends to be that one
special case.

There are others, but this is
the one special case where

oxygen does not have a minus
2 oxidation state.

Let's look at this and try to
figure out what oxygen's

oxidation state would be
in hydrogen peroxide.

So in this case, the
hydrogen-oxygen bond, oxygen

is going to hog the electron
and hydrogen is

going to lose it.

So it's going to have
a plus 1 there.

Same thing on the side.

Oxygen, at least on this bond,
is going to have a plus 1.

It's going to gain
an electron.

What about from this other bond
with the other oxygen?

Well, there's no reason why
one oxygen should hog the

electron from the
other oxygen.

So it's not going to have
any net impact on

its oxidation state.

So in this case, this oxygen's
oxidation state is plus 1.

This oxygen's oxidation
state is also plus 1.

So each of the hydrogens have an
oxidation number of plus 1.

You said the oxygens have an
oxidation number of minus 1.

And so you have a net of 0.

2 times plus 1, plus 2
times minus 1, is 0.

So that's just a special case.

That's a good one to
be familiar with.

Let's do another one.

Iron 3 carbonate.

And now, for the first time--
I remember when we first

encountered iron 3 carbonate.

You probably thought, hey, why
is it called iron 3 carbonate

when there are only two
iron molecules, or

two iron atoms, here?

And you're about to learn why.

Let's look at the oxidation
numbers.

So oxygen.

Oxygen's oxidation number
tends to be minus 2.

Now, if carbon is bonding with
oxygen-- let's look at the

periodic table.

We have carbon bonding
with oxygen.

Carbon can go either way.

Carbon, sometimes it likes
to give away electrons.

Sometimes it likes to
gain electrons.

When carbon is bonding with
oxygen, this right here is the

electron hog.

If we had to say who's
taking the electrons,

it's going to be oxygen.

Right?

So carbon is going to be giving
away its electrons.

But how many electrons
can carbon give away?

Well, let's see.

It has 1, 2, 3, 4 valence
electrons.

So the most it can really
do is give away

four valence electrons.

So let's go back to
the carbonate.

So the carbon could at
most give away its

four valence electrons.

So what will be the net
oxidation number for the

carbonate molecule?

For the CO3?

So this is a plus 4 oxidation,
because it only has

four to give away.

If it's bonding with oxygen,
it's going to give them away.

Oxygen is more of a hog.

Each oxygen has a minus 2.

So let's think about it.

I have plus 4 minus,
3 times minus 2.

Right?

I have 3 oxygen molecules.

So I have 4 minus 6 is
equal to minus 2.

So we can kind of view it as the
oxidation number for the

entire carbonate molecule
is minus 2.

Now, if this entire carbonate
molecule is minus 2, its

contribution to the oxidation
state for this whole kind of--

the carbonate part
of the molecule.

We have 3 carbonate molecules.

Each of them is contributing
minus 2.

So I have a minus
6 contribution.

If this is minus 6 and this is a
neutral molecule, then our 2

irons are also going
to have to have a

plus 6 oxidation state.

Because it all has
to add up to 0.

If both irons combined have
a plus 6 contribution to

oxidation state, then each
of the irons must

have a plus 3 oxidation.

Or that, in our hypothetical
world, if this happens, at

least three electrons are going
to favor the carbonate

from each of the irons.

So why is it called
iron 3 carbonate?

I think you may have figured
this out by now.

Because this is iron in its
third oxidation state.

Iron-- a lot of the metals,
especially a lot of the

transition metals-- can have
multiple oxidation states.

When you have iron 3 carbonate,
you're literally

saying, this is the third
oxidation state.

Or iron's oxidation number
in this molecule

will be positive 3.

Now, let's do another one.

This is interesting.

Acetic acid.

And I think is the first time
that I've actually shown you

the formula for acetic acid.

I won't go into the whole
organic chemistry of it.

But let's try to figure out
what the different charges

are, or the different
oxidation states.

Sometimes you'll just see
it written like this.

You'd say, OK.

Oxygens, each of those are
going to have minus 2.

Hydrogens are each going
to have plus 1.

So how are we doing so far?

So these oxygens are going
to contribute minus 4.

And then the hydrogens--
here you have plus 3.

And then here you have plus 1.

You add these up and
you get to 0.

And you're like, oh.

So the carbons must have
no oxidation state.

They must have an oxidation
number of 0.

Because we're already at 0, if
we just consider the hydrogens

and the oxygens.

So let's look at that and see
if that's actually the case.

So when carbon is bonding with
hydrogen, who's going to hog

the electrons?

When carbon is bonding
with hydrogen.

Electronegativity-- as
you go to the right.

Carbon is more electronegative.

It likes to keep the electrons,
or hog them, more

than hydrogen.

So hydrogen is going to lose the
electrons in our oxidation

state world.

It's actually a covalent bond,
but of course, we know that

when we're dealing with
oxidation states, we pretend

that it's ionic.

So in this case, your
hydrogens are

going to lose electrons.

So they're each going to have an
oxidation state of plus 1.

That's consistent with
what we know so far.

And actually, that's
another thing.

When I did this exercise, right
here, I immediately

assumed hydrogen has an
oxidation state of plus 1.

I did that because, oh,
everything else in the

molecule is carbon and oxygen,
which are more electronegative

than the hydrogen.

So the hydrogen is going to
go into its plus 1 state.

If, over here, I had a bunch of
alkali and alkaline earth

metals, I wouldn't be so sure.

I'd say, oh, maybe hydrogen
would take

electrons from them.

But anyway.

So these all gave an electron
to this carbon.

So just from these hydrogens,
that carbon would have a minus

3 oxidation state, right?

These lost electrons.

This guy gained three electrons,
so his charge

goes down by 3.

The carbon-carbon bond.

Well, there's no reason one
carbon should take electrons

from another carbon.

All carbons are created equal.

So there should be
no transfer here.

So this carbon's oxidation
status is 3.

Now what about on this side?

So we know that this hydrogen
is going to have a plus 1

oxidation state.

It's going to give its electron
to this oxygen.

This oxygen, like most oxygens,
are going to take up

two electrons.

One from this carbon, and
one from this hydrogen.

So it's going to have a minus
2 oxidation state.

This oxygen is also going
to take two electrons.

In this case, both of
them are going to be

from this orange carbon.

So it's going to have a minus
2 oxidation state.

So what's the oxidation
state of this carbon?

It lost two electrons to this
guy up here, and it lost one

electron to this oxygen
down here.

Remember, this guy got one
electron from the carbon and

one from the hydrogen.

So it lost one electron
here, two there.

It lost three electrons.

So in that reality, it would
have a plus 3 charge.

So it turns out that the average
oxidation state for

the carbon in acetic
acid is 0.

Because if you average minus
3 and plus 3, you get to 0.

And that's why I said, oh,
maybe these are a 0.

But if you actually write out
their oxidation numbers, this

green C has a minus
3 oxidation state.

And this orange C, this
orange carbon, has a

plus 3 oxidation state.

If you got this one, and I
don't think it's overly

complex, you will be an
oxidation state jock.

So I think you're all set now.

In the next video, we're going
to start exploring oxidation

reduction reactions.