-
-
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.
-