- [Instructor] Let's look at a few redox reactions
and assign oxidizing and reducing agents.
So, this is a famous reaction, the rusting of iron,
so here I have some iron metal,
and we add some oxygen gas to it,
and we form iron three oxide as our product.
So, let's assign some oxidation states,
so let's ignore the coefficients, right,
the coefficients are for balancing chemical reactions.
So, let's just think about those iron atoms
as atoms being in their elemental state,
and so therefore we know the oxidation state
of those atoms is equal to zero.
I move over here to oxygen gas and ignore my coefficient,
and think about the oxygen atoms in the O2 molecule.
Right, those atoms are in their elemental state,
and so therefore they have an oxidation state,
or an oxidation number, equal to zero.
I move over here to my product, ignore my coefficient,
I know that for compounds like this,
oxygen usually has an oxidation state equal to negative two,
so I put a minus two here,
and I like to circle these just to distinguish them
from all the other numbers on the screen.
And I have three oxygens,
each one has oxidation state of negative two,
so negative two times three gives me a total
of negative six for all my oxygens.
I know for a neutral compound,
the sum of all my oxidation states
has to add up to equal to zero,
and so therefore I know the sum
of all my irons has to be plus six,
since plus six and minus six would give me zero here.
Since I have two irons, each one must be plus three,
and so now I've assigned all my oxidation states
for this redox reaction.
Let's think about what's happening to the atoms here,
so if I think about an iron atom,
going from an oxidation state of zero
to an oxidation state of plus three,
that's an increase in the oxidation state.
All right, and we know by definition
an increase in the oxidation state means oxidation.
So, we can go ahead and say that iron is oxidized here,
and we know if iron is oxidized,
it is losing electrons, right?
So, loss of electrons is oxidation.
In this video, we're not so concerned
with how electrons iron is losing,
we'll worry about that in balancing redox reactions.
So, for right now we just know that iron is oxidized
and it is losing electrons in this reaction.
Oxygen, let's look at the oxidation states.
It's going from an oxidation state of zero
to an oxidation state of negative two,
that's a decrease in the oxidation state.
A decrease, or a reduction in the oxidation state is,
of course, the process of reduction.
So, oxygen is being reduced here.
And so, we know that when something is being reduced,
it's gaining electrons, Leo the lion goes ger,
so gain of electrons is reduction, like that.
Now, once again, for this video,
we're not concerned with the exact number of electrons,
but we know that the electrons that iron loses
are the same electrons that oxygen gains.
And so, that number is actually the exact same.
So, iron, by losing electrons,
is allowing oxygen to be reduced.
And so, therefore you could say that iron
is the agent for the reduction of oxygen,
and so we're going to say that iron
is the reducing agent here.
So, iron is the reducing agent, all right.
Thinking about it the other way, all right,
so if we think about oxygen,
oxygen is gaining electrons from iron,
so oxygen is allowing iron to be oxidized.
So, oxygen is the agent for the oxidation of iron,
and so therefore oxygen is the oxidizing agent.
So I'll go ahead and write oxidizing agent here.
And so, this could be confusing to students because
if you think about it, right,
it makes sense that iron's being oxidized here,
but then you're like, you're changing, you say oh,
it's the reducing agent.
It's not really changing anything,
it's just the definition of your reducing agent.
Your reducing agent itself is being oxidized,
and it's allowing something else to be reduced.
And so, if you think about it here for oxygen, right,
so those oxygen atoms have a decrease
in their oxidation state, they are being reduced,
but by the fact that they're being reduced,
they allow something else to be oxidized,
so they are the agent for the oxidation of something else.
So, you could say that O2 is your oxidizing agent.
So, let's go ahead and look at this little table here
where I've summarized this for you,
so we have reducing agents and oxidizing agents, right.
So, we'll stick with the same colors
that we've been using here, so a reducing agent
is the agent for the reduction of something else,
but in itself is being oxidized and so therefore
it's losing an electron or many electrons, right?
So, common reducing agents would be things
that lose electrons easily, so things like metals.
If I think about oxidizing agents, right,
so oxidizing agents over here,
so the agent for the oxidation of something else,
therefore in itself must be being reduced.
And if something is being reduced,
it's gaining an electron or gaining many electrons, right?
So, something that would gain electrons,
you could think about very reactive non-metals,
so things like oxygen that we talked about
in the previous example.
So, let's do one more, and let's see if we can figure out
the oxidizing agent and reducing agent for this reaction.
So, once again we'll ignore our coefficients.
Think about our rules for oxidation states,
we know that for a monatomic ion,
the oxidation state is equal to the charge on that ion,
and so if the charge is two plus,
we know the oxidation state is two plus, or plus two.
Right, I think about chlorine,
think about the atoms in the chlorine molecules,
those atoms are in their elemental state,
and so we know an atom in its elemental state
has an oxidation state equal to zero.
I move over here to iron three plus, ignore my coefficient,
a monatomic ion, has an oxidation state
equal to the charge on that ion,
so the oxidation state is plus three, here.
And I think about the chloride N ion, I ignore the two,
I think about the negative charge, and negative one charge,
meaning the oxidation state is negative one.
So, when we look at what happened in this redox reaction,
right, we know that we went from an oxidation state
of two plus to an oxidation state of three plus,
that's an increase in the oxidation state, all right?
So, we know an increase in the oxidation state
means that iron two plus was oxidized in this reaction,
so iron two plus, let me go ahead and use these colors here,
so iron two plus was oxidized.
All right, let's look at chlorine now,
so a chlorine atom went from an oxidation state
of zero to an oxidation state of negative one,
that's a decrease in the oxidation state,
a decrease in the oxidation state is reduction.
So, we can say that chlorine was reduced.
All right, so, let's now think about
oxidizing and reducing agents.
So, iron two plus was oxidized, right,
and the fact that it was oxidized, right,
it's losing electron, right?
So, it is itself being oxidized,
but it's causing the reduction of something else,
so it is our reducing agent.
So, iron two plus is being oxidized,
but it is our reducing agent.
And we think about chlorine, right,
chlorine is being reduced,
the chlorine atoms are being reduced,
by in that in so doing, right,
they cause the oxidation of the iron two plus.
And so therefore, we can say that chlorine
is our oxidizing agent.
And so, that's how to think about redox reactions.
If your oxidation states change,
you can consider it to be a redox reaction,
and then once you look at those oxidation states,
you can figure out what is oxidized and what is reduced,
and you can also figure out
which one is your oxidizing agent,
and which one is your reducing agent.