More on Oxidation States

More practice calculating oxidation states

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.

Minus 2.

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

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.

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.