Oxidation and reduction review from biological point-of-view

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.