- Oxidation–reduction (redox) reactions
- Worked example: Using oxidation numbers to identify oxidation and reduction
- Balancing redox equations
- Worked example: Balancing a simple redox equation
- Worked example: Balancing a redox equation in acidic solution
- Worked example: Balancing a redox equation in basic solution
- Oxidation–reduction (redox) reactions
By assigning oxidation numbers to the atoms of each element in a redox equation, we can determine which element is oxidized and which element is reduced during the reaction. In this video, we'll use this method to identify the oxidized and reduced elements in the reaction that occurs between I⁻ and MnO₄⁻ in basic solution. Created by Sal Khan.
- [Instructor] What we have here is a reaction that involves iodine, manganese, oxygen, and hydrogen. And what we wanna do in this video is think about which of the elements are being oxidized in this reaction and which of the elements are being reduced in this reaction. And pause this video and see if you can figure that out before we work through it together. All right, now let's work through it together. And the way that I will tackle it, and you might have tackled it or I suggest you tackle it, is to figure out the oxidation numbers for each of the elements as we go into the reaction, as they are entering the action and as they are exiting the reaction, or I guess you could say on either side of the reaction. So first, let's look at this iodine right over here. Well, each iodine has a negative one charge. And so it's quote hypothetical charge, which isn't so hypothetical in this case, which would be its oxidation number is negative one. Now let's move over to this permanganate ion right over here. Now this one's a little bit more involved to figure out the oxidation numbers. But what we generally remember is that oxygen is quite electronegative. It is likely to hog two electrons and when we think about hypothetical charge with oxidation numbers, oxygen is going to have eight negative two oxidation number because it likes to hog those two extra electrons. And so if each of these four oxygens has a hypothetical charge of negative two, that would be negative eight total and we see that this entire ion has a negative one charge. So that means that the manganese has to have a hypothetical charge, an oxidation number of plus seven. So I just wanna review that one again because this is a little bit involved. We said oxygen, we're gonna go with the negative two 'cause it likes to hog two electrons. We have four of them. So if you add all that together, you're at negative eight and the whole ion has a negative one. So what plus a negative eight is going to be negative one? Well, positive seven. And so that's manganese's oxidation number as we enter into the reaction on this side of the reaction. And then let's look at the water. Well, water, both the hydrogen and oxygen, these are ones you'll see a lot. This oxygen is going to have a negative two oxidation number and each of those hydrogen atoms are going to have a plus one oxidation number because in that water molecule. We know that the oxygen hogs the electrons, these are covalent bonds. But if we had to assign kind of a hypothetical charge where we said, all right, well, let's just say the oxygen takes those two electrons and each of those hydrogens will lose an electron and have a plus one oxidation number. Now let's look at the right-hand side of this reaction. What's going on with these iodines here? Well, in this iodine molecule, they aren't gaining or losing electrons, so your oxidation number is zero. Then let's move on to the next compound. Each of these oxygens have an oxidation number of negative two. And so what would be manganese's oxidation number? Well, the compound is neutral. Two oxygens at negative two is gonna be negative four. So in order to be neutral, the manganese must be at plus four, an oxidation number of plus four. And then last but not least, if we look at these hydroxide anions, each of the oxygen is going to have a negative two oxidation number. And then the hydrogen is going to have a plus one and we can confirm that that makes sense. Negative two plus one is going to be negative one for each of these ions. So now, let's just think about who's been oxidized and who's been reduced. And remember, oxidation is losing electrons. Oil rig, reduction is gaining electrons, or reduction is a reduction in the oxidation number. So first, let's look at the iodine. We go from an oxidation number of negative one to zero. So to go from an oxidation number of negative one to zero, you need to lose electrons. So it has been oxidized. Oxidized. Let me write that down. The iodine has been oxidized. Now let's look at the manganese. We go from a plus seven to a plus four. Our oxidation number has gone down. It has been reduced. Now let's look at the oxygen. Well, everywhere, the oxygen has an oxidation number of negative two, so nothing there. And then same thing for the hydrogens. Plus one on both sides, so nothing there. So the iodine has been oxidized and the manganese has been reduced.