Simple redox reactions | Chemistry | Khan Academy

You have probably heard about this word: oxidation, oxidizing, or antioxidants, and stuff like that. But what exactly does it mean, and what can you do knowing about it? Well, let's find out.

Oxidation has the word oxygen in it; you can see that, right? So we can guess that it's probably got something to do with the oxygen, right? And that's how we defined it earlier—not currently, but that's how it started. So let's start with that. Oxidation can be thought of as adding oxygen atoms.

For example, take a look at these reactions. What do you notice? We see that, like for example in this first reaction, oxygen is getting added to magnesium, right? Therefore, we say magnesium is being oxidized, or magnesium is undergoing oxidation. Remember, oxygen is diatomic in nature; that's why there's an O2 over here. Okay, but the same thing is happening over here: look, oxygen is being added to ion; you can see that over here. Therefore, we say ion is being oxidized; our ion is undergoing oxidation. Same thing over here: oxygen, look, gets attached to carbon, so carbon is being oxidized; carbon is undergoing oxidation.

But a question I always had was, well, why do we give a name to these kinds of reactions? What's so special about oxygen? Well, I don't think there's anything special about oxygen, but I think these reactions are important. For example, if you look at ion oxide, which we call it ion (III) oxide, because the charge on the ion over here will be +3, we'll get to that. But this is the molecule that you find in rust. This is rust. I mean, rust also involves water, so the actual reaction also involves water molecules over here. But since rusting involves oxygen, that's an important reaction for us, because we want to prevent rusting.

And when you're burning stuff, well, guess what? Oxygen is involved. Combustion involves oxygen; that's also another important reaction for us. And so, since these reactions are important and they involve oxygen, that's probably the reason why we gave a name to it, and we call it oxidation.

But anyways, this was the earlier definition of oxidation. Later on, we thought, okay, let's just think about what's really going on with the electrons over here. So, for example, over here, we realize that since magnesium is a metal, it loses electrons. In fact, over here, it's going to lose a couple of electrons, and magnesium will now become magnesium 2+ because it has lost two electrons. And oxygen, well, it gains two electrons, and therefore oxygen becomes 2-. And that's why they stick to each other.

Okay, this is how that reaction is formed. So what we notice is that when magnesium is undergoing oxidation, it's losing electrons. Well, let's see if the same thing is happening over here. Well, yeah, ion is also a metal, and it's losing electrons over here. In fact, ion loses three electrons to form ion 3+, and that's why this is called ion (III) oxide, because it has a charge of +3. It loses three electrons over here, and oxygen gains those electrons. I mean, each atom of oxygen gains two electrons, making it oxygen 2-.

And so because oxygen has a 2- charge and ion has a 3+ charge, now you can notice three atoms of oxygen have a total of -6, and two atoms of ion have a total of +6. +6 - 6 very nicely becomes neutral. But again, what's important is that ion, which is undergoing oxidation, is losing electrons.

And what about this one? Well, this is actually a slightly more interesting case, so we'll get back to it towards the end of the video. But this means in most cases when you add oxygen to something, that something loses electrons. So we said, "Hey, hey, hey! Oxidation is involving loss of electrons." So let's call that as the definition of oxidation, and that is our new definition. Because it's generalizing the definition of oxidation, we now say forget about oxygen atoms: anything that loses electrons we will call it oxidation.

So magnesium is being oxidized because it's losing electrons; ion is being oxidized because it is losing electrons. That is the current definition of oxidation. Okay, now if something loses electrons, something else should gain electrons, and in this case, the oxygen is the one that gains electrons. So we thought, okay, let's give a name to gaining electrons as well. Gaining electrons is what we call reduction.

And now immediately looking at this, it feels weird, right? You're gaining electrons; why do we call it reduction? Okay, this sounds weird, and there might be some history behind it. We'll not worry too much about the naming and where it came from, but here's how I think about it. When you gain electrons, since electrons are negative, you are gaining negative charge.

Okay, and that means look at your charge: your charge was zero to begin with over here; oxygen has zero charge to begin with. But now it has a negative charge; that is, in some sense, reduction, right? It has reduced. So gaining negative reduces your overall charge because you gain more negative charge. And so that kind of sort of makes sense to me in calling that reduction.

And so oxygen over here is getting reduced. Okay, but of course, what if we forget this? How do we remember that oxidation is losing electrons and reduction is gaining electrons? There are a couple of mnemonics you can use. For example, one mnemonic that I like is: Leo the Lion goes "g!" Here, losing electrons is oxidation, gaining electrons is reduction. Another mnemonic you can use is: oil rig; oxidation is losing, reduction is gaining.

My friends keep debating about which one is better, but I personally like to think in terms of charge. When I gain electrons, I gain negative charge, and therefore it's a reduction. And that's how I like to remember that gaining electrons is reduction; therefore, losing electrons must be oxidation.

But now we have multiple ways to remember—whichever one works for you. Anyways, now think about this: if you have a reaction in which something is getting oxidized, meaning something is losing electrons, then something else must be gaining electrons. So something else must be reduced. And therefore, such reactions in which something gets oxidized and something else gets reduced, we give a name to it: we call them redox reactions, because both reduction and oxidation must be happening.

And now here's my question: do redox reactions need to involve oxygen? The answer is no, because our definition has become more general. Sorry, oxygen, but we don't need oxygen anymore. Any reaction in which something is losing electrons and something else is gaining electrons becomes a redox reaction.

So let's look at a few. Here are a few reactions; our goal now is to analyze each reaction and think about which one is undergoing oxidation and which one is undergoing reduction.

So let me start with this one: so we have ion that reacts with sulfur to give you ion sulfide. So how do I know which is undergoing oxidation and which is undergoing reduction? Well, I know ion is a metal, so it loses electrons. So I just keep track of charge over here. Ion is neutral, a single atom of ion, zero charge; sulfur is also neutral, zero charge.

Sulfide usually gets a charge of -2, so ion must have gotten a charge of +2, meaning ion must have lost two electrons. Sulfur must have gained two electrons, giving you ion sulfide, or we say ion (II) sulfide. But what's important for us is that look, ion is the one that has lost electrons, which means ion got oxidized. Sulfur is the one that gained electrons, and therefore sulfur is reduced. Again, sulfur's charge has reduced; sulfur is the one that underwent reduction.

Now, it would be a great time for you to pause the video and try to analyze the remaining three reactions.

All right, let's look at this one. Sodium has a charge of zero; chlorine also has a charge of zero. Over here, I know sodium is a metal, so it's going to lose an electron. In fact, it loses one electron, giving a +1 charge, and chlorine gains an electron, giving it a -1 charge. So look, it's the sodium that lost electrons; sodium got oxidized, and chlorine that gained electrons, so chlorine got reduced.

All right, here's the next one. This looks slightly more complicated, but we'll work it out. Silver, being a metal over here, loses an electron. So it loses one electron; it actually has a +1 charge. Sulfur gains two electrons, meaning it has a -2 charge. In fact, that's the reason why we have two AGs, giving you +2 and -2; that balances it out.

Okay, anyways, that's the charge over here. Aluminium—look, it's not bonded to anything, so it has a zero charge over here. But what happens over here? Hey, now aluminium gets bonded to sulfur, and aluminium is a metal; it's the one that loses electrons, and it usually loses three electrons, so it gets +3. Sulfur gets -2; you can just look at the molecular formula and also guess what the charge they must have had, how many electrons they must have lost and gained.

And finally, now look: silver is not bonded anymore, so it has a zero charge. So if you look at silver, it went from +1 to 0. Ooh, ooh! Silver got reduced because its number has reduced; it has gained electrons. Ah, so silver underwent reduction. What about aluminium? Aluminium—look, it underwent oxidation because it must have lost electrons. So it has gone from 0 to +3, so aluminium underwent oxidation, and silver underwent reduction. What about sulfur? Hey, sulfur's charge has not reduced at all; it has not changed at all. Therefore, sulfur neither underwent oxidation nor reduction.

So this was an interesting one. Okay, that brings us to the last one. Again, if you haven't tried it, feel free to try it now, because this is an interesting one.

All right, here it goes. So sodium chloride, we already saw, +1, -1 charge. Silver fluoride: silver +1 and fluorine -1. Sodium fluoride: sodium is again +1, fluorine -1. Silver chloride: silver is +1, chlorine is -1. What do you notice? Which elements undergo oxidation? Which elements undergo reduction? None of them, because none of their charge has changed, which means that means none of them have lost or gained electrons. So this is not a redox reaction, because nothing underwent oxidation; nothing underwent reduction.

When I first came across such reactions, I was surprised, because in my mind, any chemical reaction involves transfer of electrons. So in my mind, it was like every reaction must be redox reactions. But right in front of your eyes, you can see that not every reaction is a redox reaction.

Okay, finally this brings us back to our carbon dioxide example. Why was this interesting? You know? Well, because in all these cases, something was losing electrons and something else was gaining electrons, right? But over here, carbon does not lose electrons to oxygen. Instead, carbon and oxygen share the electrons, forming a covalent bond.

So how does it make sense according to this new definition? According to the current definition, how does it make sense to talk about oxidation and reduction over here? Well, it is slightly more complex and interesting, but here's the gist of it: the electrons are not shared equally; oxygen has a tendency to pull those electrons more towards itself compared to carbon.

We say oxygen is more electronegative. Electronegativity is the tendency of pulling the shared pair of electrons towards yourself. Okay? So, oxygen, since it's more electronegative compared to carbon, it pulls the shared pair of electrons more towards itself, kind of hogs on those electrons. And as a result, oxygen also ends up getting a negative charge— but not a complete negative charge, because it’s not a complete transport. We say it gets a partial negative charge.

I know it sounds complicated, but at the end of the day, it's still kind of sort of getting a partial negative charge. So in that sense, it still makes sense to say oxygen is reduced, and therefore carbon gets a partial positive charge, meaning oxidation: carbon is oxidized.

Okay? So even here, we can think in terms of oxidation and reduction.

Anyways, our last question could be: what can we do now that we understand what oxidation and reduction really is? Well, here's an example: now that we know that rusting is basically ion losing electrons to oxygen, one way to prevent rusting is to coat it with a metal that loses electrons more readily compared to ion. And one such example is zinc.

And that's why if you take an ion nail and you don't want it to rust, we usually coat it with zinc. It kind of gives you double protection. First of all, it kind of forms a barrier so it doesn't even allow oxygen to reach ion. It’s the zinc that gets oxidized. But even if there's a breach somewhere over here, an ion is exposed; even at that site of exposure, oxygen is much more likely to react with zinc because zinc much more readily gives out its electrons compared to ion.

We say zinc is more reactive compared to ion. So using this knowledge, we can protect our precious ion from getting rusted. That's awesome, right?