What I want to do in this video is review what we
learned from our chemistry classes about oxidation and
the opposite of oxidation, reduction.
And then see how what we learned in our chemistry class
relates to the way that a biologist or biochemist might
use these words.
And hopefully we'll see that they're the same thing.
So just as a bit of review, if you watched the chemistry
playlist. Oxidation, you can view it-- and actually there's
a famous mnemonic for it.
It's: OIL RIG Where the oil tells us that oxidation is
losing-- I put it in quotes because you're not necessarily
losing the electrons; I'll show you what I mean-- is
losing electrons.
This is what you should have learned in
your chemistry class.
And then you also learned that reduction is gaining.
And I'll put that in quotes as well.
Is gaining electrons.
And I put that in quotes because you're not necessarily
gaining electrons.
You're more hogging it.
And the reason why it's called reduction, is because if you
are gaining electrons your notional charge, if you really
were gaining them, is being reduced.
And the reason why this is called oxidizing is because
you tend to lose electrons to oxygen.
Although it doesn't have to be oxygen.
It could be any molecule that will hog
electrons away from you.
And I think a nice example would be fair to kind of make
this a little bit more concrete.
Let's say I took some molecular hydrogen, it's in a
gaseous state, and I were to combust that with some
molecular oxygen.
This is what happened on the Hindenburg.
They filled a balloon full of hydrogen and you get a little
bit of spark, expose it to oxygen, and you're going to
have a big explosion.
But in the process, for every mole of molecular oxygen, if
you have two moles of molecular hydrogen-- I'm just
making sure the equation is balanced-- you're going to
produce two moles of H2O plus a ton of heat.
This thing is really going to blow.
What I want to do, I mean we could talk about the
Hindenburg but really, the whole reason why I even wrote
this is, I want to show you what is getting oxidized and
what is getting reduced.
So in this situation right here on the hydrogen, the
molecular hydrogen just looks like this.
You have a hydrogen-hydrogen bond.
They're each sharing an electron with the other one so
that they both can pretend their 1s orbital
is completely filled.
So they're not losing electrons to each other.
They're not hogging electrons one from the other.
So we say that they have a neutral oxidative state.
They haven't gained or lost electrons.
They're just sharing them.
And the same thing is true for the molecular oxygen.
And here you actually have a double bond
with the two oxygens.
But they're both oxygens, so there's no reason why one
would gain or lose electrons from the other.
But when you go on this side of the equation, something
interesting happens.
You have, for every oxygen is connected to two hydrogens.
And the way to think about is that oxygen is hogging each of
these hydrogen's electrons.
So hydrogen has this one electron on its valence shell.
The deal with most covalent bonding is, hey, I give you an
electron, you give me an electron and we both have a
complete pair.
But we know, or hopefully we can review, that oxygen is
much more electronegative than hydrogen.
This is a little bit of glucose that's left over from
our cellular restoration video.
You can ignore it for now but I'm going to connect all this
in a future video.
But if we look at our periodic table, if you remember from
the chemistry playlist, electronegativity increases as
we go to the top right of the periodic table.
These are the most electronegative elements over
here, these are the least electronegative.
And all electronegative means is, likes to hog electrons.
So even though oxygen and hydrogen are in a covalent
bond in water-- they're sharing electrons-- oxygen is
more electronegative, much more electronegative than
hydrogen, so it's going to hog the electrons.
And actually if you take some elements on this side and you
bond them with some guys over here, these guys are so much
more electronegative than these left-hand elements that
they'll actually completely steal the electron, not just
hog it for most of the time.
But when you talk electronegativity, it just
means, likes the electrons.
So when you look at this bond between hydrogen and oxygen,
we saw from the periodic table, oxygen is a lot more
electronegative, so the electrons spend a lot more
time on oxygen.
We learned about hydrogen bonding.
We learned that it creates a partial negative charge on
that side of the water molecule and creates partial
positive charges on this side.
And electrons still show up around the hydrogens
every now and then.
When you talk about oxidation and reduction you say, look
there's no partial charge.
If one guy is kind of hogging the electron more, for the
sake of oxidation states, we're going to assume that he
took the electron.
So for an oxidation state, we'll assume that the oxygen
in water takes the electron and we'll give him an
oxidation state of one minus.
Or the convention is, you write the charge after the
number for oxidation states.
So you don't confuse it with actual charges.
So this has a one minus because, from an oxidation
state point of view, it's taking the electron.
It's gaining the elctron.
That's why I put it in quotes.
Because you're not really gaining it.
You're just gaining it most of the time.
You're hogging electrons.
And likewise, this hydrogen-- let me be careful, this
isn't-- he got one electron from this hydrogen and you got
another electron from this hydrogen.
So instead of saying one minus, it should be two minus.
It should be two minus, because he's hogging one
electron from here and one electron from there.
And in general, when oxygen is bonding with other non-oxygen
atoms or non-oxygen elements, it tends to have a two minus
or a negative two oxidation state.
So if this guy's two minus, because
he's gained two electrons.
Let me write that in quotes.
Gained two electrons.
We know that he really didn't gain them, that he's just
hogging them.
These guys lost an electron each.
So this guy's oxidation state is going to be one plus.
And this guy's oxidation state is going to be one plus.
So you could say, by combusting the hydrogen with
the oxygen, that the hydrogens-- before they had a
zero oxygen state, each of these hydrogens had a zero
oxygen state-- now they have a one plus oxidation state
because they lost their electrons when they bonded
with the oxygen.
So we say that these hydrogens have been oxidized.
So, due to this reaction, hydrogen has been oxidized.
Why has it been oxidized?
Because before, it was able to share its
electrons very nicely.
But then it bonds with oxygen, which will hog its electrons.
So the hydrogen is losing its electrons to the oxygen, so
it's been oxidized.
Similarly, the oxygen, due to this combustion reaction, has
been reduced.
Why has it been reduced?
Here it was just sharing electrons.
It wasn't losing or gaining it.
But here when it's bonded with an element with much lower
electronegativity, all of a sudden it can start hogging
the electrons, it gains electrons.
So this hypothetical charge is reduced by two.
And if I wanted to actually account for all of the
electrons, because we're talking about losing electrons
and gaining electrons, we can write two half reactions.
This should all be a little bit of review from your
chemistry class.
But it never hurts to see this again.
I'm going to throw this in the biology playlist so that you
biology people can hopefully refresh your
memory with this stuff.
We can write two half reactions.
We could say that we started off with two moles of
molecular hydrogen.
And they have no oxidation states, or they're neutral.
So I could write a zero there if I want.
And then I end up with-- on the other side-- I end up with
two moles of H2.
But each of the hydrogens now, have a plus
one oxidation state.
Or another way to think about it is, each of these-- there's
four hydrogens here.
This is molecular hydrogen has two hydrogens and we have two
moles of this.
So there are four hydrogens here.
Each of the four hydrogens lost an electron.
So I can write this.
So, plus four electrons.
That's the half reaction for hydrogen.
It lost four electrons.
So this is another way of saying that hydrogen is
oxidized because it lost electrons.
OIL: oxidation is losing.
And then the other half reaction, if I were to write
the oxygen.
So I'm starting with a mole of molecular oxygen and I'm
adding to that four electrons.
I can't make electrons out of nowhere.
I'm getting the electrons from the hydrogen, I'm adding to
the oxygen.
And so the half reaction on this side, I end up with two
moles-- I could write it like this-- two moles of oxygen.
And each of them have an oxidation state of two minus.
So these are the half reactions.
And all this is showing is that the hydrogen, over the
course of this combustion reaction, lost electrons.
And that the oxygen gained the electrons that the hydrogen
lost. So this tells us that oxygen is reduced.
Now this is all fair and good and this is all a bit of
review of what you learned in chemistry class.
But now I'm going to make things even more confusing.
Because I'm going to introduce you to how a biologist
thinks about it.
So-- and it's not always the case.
Sometimes the biologist will use the definition you learned
in your chemistry class.
But a biologist-- or many times in many biology
textbooks-- they'll say-- and this used to confuse me to no
end, really-- that oxidation is losing hydrogen atoms. And
reduction is gaining hydrogen atoms.
And at first when I got exposed to this, I was like, I
learned it in chemistry class and they talk about electrons.
Hydrogen atoms, you know it's a proton and an electron, how
does it relate?
And the reason why these two definitions-- this is really
the whole point of this video-- the reason why this
definition is consistent with this one is because in the
biological world hydrogen is what tends
to get swapped around.
And it tends to bond with carbon, oxygen,
phosphorous, nitrogen.
And if we look at the periodic table, and we see where
hydrogen is, and we see where carbon, nitrogen, oxygen and
phosphorous and really all this other stuff is, you see
that all of the stuff that in biological systems, hydrogen
tends to bond with, the things it tends to bond with are
much, much more electronegative.
So if a carbon is bonding with a hydrogen, the carbon is
hogging that electron.
And then if that hydrogen gets transferred to an oxygen,
along with the electron, the carbon will lose the hydrogen
atom, but it really lost the electron that
it was hogging before.
And now the oxygen can hog that electron.
So these are really consistent definitions.
And the whole reason why I showed you this example is
because the biological definition doesn't apply here.
I mean, you could say, well, oxygen is definitely gaining
hydrogens in this reaction.
So we can definitely say that oxygen is being reduced still,
according to the biological definition.
But you can't really say that hydrogen is
losing hydrogens here.
In this situation, hydrogen is just losing electrons.
It's not losing itself.
I guess you could say it's losing itself because it's
being taken over.
But the biological definition just comes
from the same notion.
That when hydrogen bonds with most things in biological
compounds, it tends to give the electrons.
So if a carbon loses a hydrogen and gives it to an
oxygen, the carbon will lose that hydrogen's electron that
it was able to hog.
And now the oxygen is hogging it.
So the carbon would be oxidized and the oxygen would
be reduced.
Hope that doesn't confuse you.
In the next video I'll show you a couple more examples.
And the whole reason why I'm doing this is to apply this to
cellular respiration.
So that you don't get confused when people talk and say that,
oh the NAD is being reduced when it picks up the hydrogen.
Or it's being oxidized when it loses the hydrogen, and so
forth and so on.
I wanted you to see that these are the same definitions that
you learned in your chemistry class.