We now know what an oxidation state or an oxidation number
is and what it means when things
are oxidized or reduced.
So let's see how that actually happens in reactions.
So what we're going to study in this video is oxidation
slash reduction reactions.
All that is is a reaction where somebody's being
oxidized, which means that electrons are being taken away
from them, and someone's being reduced, which means that
they're being handed electrons, or they're taking
electrons away from someone else.
Sometimes this has been termed-- because you have the
R-E-D in reduction and you have the ox in oxidation, and
they switch them around and they call this a redox
reaction, which sounds like a very fancy chemistry term, but
it just means a reaction where something is getting oxidized.
And if something is getting oxidized, something else is
getting reduced.
So let's study a couple of them.
So right here, I actually have combustion.
This is actually methane.
So methane is a fuel.
You can make a motor powered on methane right there.
It's a hydrocarbon.
Most fuels that we use are hydrocarbons, which just mean
carbons bonded in a bunch of different ways to a bunch of
different hydrogens.
If you add enough heat for the reaction to happen-- so its
activation energy you have to put it in there with some
oxygen around-- it's going to combust and it's going to
produce carbon dioxide and water-- and I didn't put it
here-- and even more heat than you put into it.
So it's actually an exothermic reaction.
It produces more heat than you put into it.
I'll do a lot more of that in future videos on endothermic
and exothermic reactions.
But anyway, we care about the oxidation and the reduction.
So let's see if anyone's getting oxidized, or anyone is
getting reduced here.
So let's look at their oxidation
numbers, or their states.
Here, carbon is bonded to 4 hydrogens.
Who's giving, who's taking?
So let's go to our Periodic Table.
Carbon is here.
Hydrogen is here.
Carbon is more electronegative.
It's adjacent to the Three Musketeers of
electronegativity.
These guys are the most electronegative.
We always ignore the noble gases because they pretty much
don't react at all.
They're always happy with their 8 valence electrons.
These guys love to get electrons.
They're small molecules.
Their outermost shell is close to the nucleus.
They're so close to becoming noble that they just love
hogging electrons.
Carbon's almost there.
Carbon's much further to the right on the periodic table
than the hydrogen.
So if you have carbon bonding with hydrogen, carbon is going
to be the electron hog.
So in the situation, let's go down here.
So if carbon's the electron hog and hydrogen is having its
electrons taken away from it-- remember, this is all kind of
hypothetical.
It's more partial, is the reality.
But if you had to pick or choose, hydrogen's going to
lose an electron each.
So it's going to have an oxidation state of plus 1 for
each hydrogen.
You have 4 hydrogens each giving up an electron.
So the carbon must be taking 4 electrons.
So its oxidation number is minus 4.
It's taken 4 electrons so its charge will go down by 4.
So that's why it's negative.
Fair enough.
What's the oxidation state of this oxygen right there?
Well, it's just bonded to itself.
There's no reason to believe that 1 oxygen should be able
to take any electrons from another oxygen.
So it has a 0 oxidation state.
It's not hogging more than its fair share of electrons that
it was essentially born with.
Now.
After the combustion occurs, what are
the oxidation numbers?
Well, I have 2 oxygens bonded with carbon.
Now we know oxidation was, to some
degree, named after oxygen.
It is one of the most electronegative.
Almost anything bonded with oxygen is going to be giving
up its electrons.
We also know that oxygen likes to take 2 electrons.
If it takes 1, it gets here.
If 2, it gets here.
It has a configuration like it has 8 valence electrons.
So its typical oxidation number is minus 2.
So carbon, in this in the situation, each oxygen is
going to have a minus 2 oxidation number.
The carbon dioxide molecule is neutral.
So everything has to add up to 0.
The oxygens, you have 2 of them, each of them with a
minus 2 oxidation number.
So that's minus 4.
So carbon must be plus 4, which means that it has given
up 4 electrons, right?
Really it only has 4 electrons to give up.
It has 1, 2, 3, 4 valence electrons in its second shell,
which is its reactive shell.
So its oxidation number is plus 4.
Now let's look at the water.
We've done about a bunch.
When the hydrogens each give up an electron, they have an
oxidation number of 1.
Each oxygen takes 2 electrons.
There's only 1 of them.
So it's minus 2.
So what's going on here?
What's getting oxidized, what's getting reduced?
You have the carbon.
It went from an oxidation number of minus 4 an oxidation
number of plus 4.
So let me just draw what's happening to the carbon.
Carbon goes from minus 4, which means it's hogging 4
electrons, to a situation where it's having 4 electrons
being hogged from it.
It's kind of giving away 4 electrons.
So someplace in this process, this guy must have given away
8 electrons.
This is a difference of 8 electrons.
So plus 8 electrons.
So what's happened to carbon?
Carbon, has it been oxidized or reduced?
Its charge has gone up.
So it hasn't been reduced.
Or the other way is that you can say electrons have been
taken away from it.
So it has been oxidized.
Now, let's look at the oxygen itself.
Oxygen, over here you have 4 oxygen molecules.
So I'll just write 4 oxygen molecules here.
And they all have an oxidation state of 0.
From an oxidation-number point of view, they're neutral.
On the side, we have 2, and then we
have 4 oxygen molecules.
What's their oxidation state?
They're all minus 2.
So each of these oxygens must have gained an electron.
So actually, let me let me re-write this reaction.
Let me erase a little bit of it.
Even better, let me just move this over.
There you go.
Let me fill in that, so aesthetically it's pleasing.
There you go.
All right.
So we have 4 oxygens with a 0 oxidation state turning into 4
oxygens with a minus 2 oxidation state.
So each of these 4 oxygens took 2 electrons.
There's 4 of them, so we must have gained 8 electrons.
So what happened to oxygen?
It's oxidation state went down, or its hypothetical
charge went down, or its charge was reduced.
So this is reduced oxygen.
What was it reduced by?
It was reduced by the carbon.
What was the carbon oxidized by?
It was oxidized by the oxygen, which oxygen tends to do.
It tends to oxidize things.
What is the oxidizing agent?
Well it's the thing that did the oxidizing.
So the oxygen is the oxidizing agent.
What is the reducing agent?
Well it's the thing that did the reducing?
It's the reducing agent, is the carbon.
So you see in this redox reaction, carbon was oxidized.
It gave away 8 electrons from this state to that state,
hypothetically, and oxygen was reduced.
Its charge was reduced by a total of 8 electrons, 2 for
each oxygen.
So its oxidation number went down.
Let's do a couple more.
Each of the things that I wrote-- just to kind of do a
side-note-- this is called a half reaction, because I'm
writing just what happened to the carbon.
Here I'm writing just what happened to the oxygen.
I did ignore something.
I did ignore the hydrogen.
I encourage you to do something like
this with the hydrogen.
But you'll see that hydrogen was
neither oxidized or reduced.
On both sides of the equation, all the hydrogens had a plus 1
oxidation state.
Let's do another one.
Here I have another combustion situation.
This is what happened to the Hindenburg.
They filled a balloon full of hydrogen because they wanted
it to float, because hydrogen is a very light gas.
Unfortunately, there must have been a spark in the presence
of oxygen, and it combusted.
And actually, this is what is used for a
rocket fuel as well.
If you have liquid hydrogen-- actually, I don't think they
have liquid hydrogen.
Well, I don't know enough about rocket science.
I'll have to do another video that in the future.
But let's look at the oxidation states.
What's the oxidation state of hydrogen here, in its
elemental form?
It's 0.
2 hydrogens bonded to each other, no reason why they
should hog or be hogged by another hydrogen.
2 oxygen molecules or atoms bonded to each other?
Once again, 0 oxidation state.
Once they combust and form water, what's
their oxidation states?
We've done this multiple times.
The oxygen has a minus 2 oxidation state.
Each hydrogen has a plus 1.
So let's write the half reactions.
We have 2 hydrogens that are just happy as they are in
neutral state.
They end up being 2 H2's, or we could have written 4
hydrogens, either way.
They turn into 2 hydrogen molecules, because there's
actually 4 hydrogen atoms, with a plus 1 oxidation state.
So they must have each given away 1 electron.
Now, how many total hydrogens are there?
There are 4.
So we must have given away 4 electrons.
So this is a half reaction for hydrogen.
Now.
Let's do the same thing for our oxygen.
We have some oxygen here.
On the left-hand side, it has a neutral oxidation state.
Then we end up with 2 oxygens on the right-hand side.
I'll write it like this, 2O, each with the minus 2
oxidation states.
So each of these oxygens must have gotten 2 electrons.
So it got 2 electrons each, or it gained 4 electrons.
So what was oxidized?
Oxidized means that electrons were taken away from you.
The hydrogen was oxidized, oxidized by oxygen.
What was reduced?
The oxygen was reduced by the hydrogen.
If you actually add these two reactions up, if you make this
the left-hand side of your equation and you make this the
right-hand side of your equation, you can say, OK,
let's remove the electrons from both sides.
Then you'll end up with your original reaction.
Just to make sure our terminology is right, what's
our oxidizing agent?
It's the thing doing the oxidizing.
It's the oxygen.
What's our reducing agent?
It's the hydrogen.
Let's do one more.
So here I have iron plus some hydrochloric acid.
Let's say this is in an aqueous solution.
You end up with iron (II) chloride plus some hydrogen.
So let's do some oxidation numbers.
I'll do it fast this time.
Iron is just by itself.
It's got a 0 oxidation state.
Hydrogen with chloride, chloride's a halogen.
These guys love to take electrons.
They love to take 1 electron.
They typically have a minus 1 oxidation state.
So the chlorine is going to have minus 1.
The hydrogen is going to be plus 1.
You add them together, you get to a neutral compound.
Fair enough.
Now you go on this side.
What is chlorine's oxidation state?
Well, once again, it always likes to take 1 electron.
So this is minus 1.
But I have 2 chlorines here.
This is a neutral compound.
So the iron oxidation state must be plus 2.
What about the hydrogen?
Now it's just in its elemental form so
it's going to be neutral.
So it has a 0 oxidation state.
So what was oxidized?
Our iron went from neutral to having 2 electrons taken away
from it, which may give it a positive charge.
So plus 2 electrons got taken away.
So this was oxidized.
What about the hydrogen?
The hydrogen went from 2 hydrogens with a plus 1 state,
and essentially, they went to 2 hydrogen
with a neutral state.
So 2 electrons must have been added to the hydrogens.
Their charge was reduced.
So they were reduced.
What were they reduced by?
They were reduced by the iron.
What was the iron oxidized by?
It was oxidized by the hydrogen.
What about the chlorine?
The chlorine has a minus 1 oxidation number here.
It has a minus 1 oxidation number here.
It was a neutral relative to the redox reaction.
Anyway, I think you get the point now.
You can do this with a bunch of reactions.
But this will give you a little bit more insight of
actually what's going on, and who's
gaining or losing electrons.
In the future, it will also help us think about a lot of
the structures of molecules.
Anyway, see you in the next video.