- Oxidation and reduction
- Oxidation state trends in periodic table
- Practice determining oxidation states
- Unusual oxygen oxidation states
- Balancing redox equations
- Oxidizing and reducing agents
- 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
- Redox titrations
- Oxidation–reduction (redox) reactions
Trends in common oxidation states for main group elements. Created by Sal Khan.
Let's see if we can come up with some general rules of thumb or some general trends for oxidation states by looking at the periodic table. So first, let's just focus on the alkali metals. And I'll box them off. We'll think about hydrogen in a second. Well, I'm going to box-- I'm going to separate hydrogen because it's kind of a special case. But if we look at the alkali metals, the Group 1 elements right over here, we've already talked about the fact they're not too electronegative. They have that one valence electron. They wouldn't mind giving away that electron. And so for them, that oxidation state might not even be a hypothetical charge. These are very good candidates for actually forming ionic bonds. And so it's very typical that when these are in a molecule, when these form bonds, that these are the things that are being oxidized. They give away an electron. So they get to-- a typical oxidation state for them would be positive 1. If we go one group over right over here to the alkaline earth metals, two valence electrons, still not too electronegative. So they're likely to fully give or partially give away two electrons. So if you're forced to assign an ionic-- if you were to say, well, none of this partial business, just give it all away or take it, you would say, well, these would typically have an oxidation state of positive 2. In a hypothetical ionic bonding situation, they would be more likely to give the two electrons because they are not too electronegative, and it would take them a lot to complete their valence shell to get all the way to 8. Now, let's go to the other side of the periodic table to Group 7, the halogens. The halogens right over here, they're quite electronegative, sitting on the right-hand side of the periodic table. They're one electron away from being satisfied from a valence electron point of view. So these are typically reduced. They typically have an oxidation state of negative 1. And I keep saying typically, because these are not going to always be the case. There are other things that could happen. But this is a typical rule of thumb that they're likely to want to gain an electron. If we move over one group to the left, Group 6-- and that's where the famous oxygen sits-- we already said that oxidizing something is doing to something what oxygen would have done, that oxidation is taking electrons away from it. So these groups are typically oxidized. And oxygen is a very good oxidizing agent. Or another way of thinking about it is oxygen normally takes away electrons. These like to take away electrons, typically two electrons. And so their oxidation state is typically negative 2-- once again, just a rule of thumb-- or that their charge is reduced by two electrons. So these are typically reduced. These are typically oxidized. Now, we could keep going. If we were to go right over here to the Group 5 elements, typical oxidation state is negative 3. And so you see a general trend here. And that general trend-- and once again, it's not even a hard and fast rule of thumb, even for the extremes, but as you get closer and closer to the middle of the periodic table, you have more variation in what these typical oxidation states could be. Now, I mentioned that I put hydrogen aside. Because if you really think about it, hydrogen, yes, hydrogen only has one electron. And so you could say, well, maybe it wants to give away that electron to get to zero electrons. That could be a reasonable configuration for hydrogen. But you can also view hydrogen kind of like a halogen. So you could kind of view it kind of like an alkali metal. But in theory, it could have been put here on the periodic table as well. You could have put hydrogen here, because hydrogen, in order to complete its first shell, it just needs one electron. So in theory, hydrogen could have been put there. So hydrogen actually could typically could have a positive or a negative 1 oxidation state. And just to see an example of that, let's think about a situation where hydrogen is the oxidizing agent. And an example of that would be lithium hydride right over here. Now, in lithium hydride, you have a situation where hydrogen is more electronegative. A lithium is not too electronegative. It would happily give away an electron. And so in this situation, hydrogen is the one that's oxidizing the lithium. Lithium is reducing the hydrogen. Hydrogen is the one that is hogging the electron. So the oxidation state on the lithium here is a positive 1. And the oxidation state on the hydrogen here is a negative. So just, once again, I really want to make sure we get the notation. Lithium has been oxidized by the hydrogen. Hydrogen has been reduced by the lithium. Now, let's give an example where hydrogen plays the other role. Let's imagine hydroxide. So the hydroxide anion-- so you have a hydrogen and an oxygen. And so essentially, you could think of a water molecule that loses a hydrogen proton but keeps that hydrogen's electron. And this has a negative charge. This has a negative 1 charge. But what's going on right over here? And actually, let me just draw that, because it's fun to think about it. So this is a situation where oxygen typically has-- 1, 2, 3, 4, 5, 6 electrons. And when it's water, you have 2 hydrogens like that. And then you share. And then you have covalent bond right over there sharing that pair, covalent bond sharing that right over there. To get to hydroxide, the oxygen essentially nabs both of these electrons to become-- so you get-- that pair, that pair. Now you have-- let me do this in a new color. Now, you have this pair as well. And then you have that other covalent bond to the other hydrogen. And now this hydrogen is now just a hydrogen proton. This one now has a negative charge. So this is hydroxide. And so the whole thing has a negative charge. And oxygen, as we have already talked about, is more electronegative than the hydrogen. So it's hogging the electrons. So when you look at it right over here, you would say, well, look, hydrogen, if we had to, if we were forced to-- remember, oxidation states is just an intellectual tool which we'll find useful. If you had to pretend this wasn't a covalent bond, but an ionic bond, you'd say, OK, then maybe this hydrogen would fully lose an electron, so it would get an oxidation state of plus 1. It would be oxidized by the oxygen. And that the oxygen actually has fully gained one electron. And you could say, well, if we're forced to, we could say-- if we're forced to think about this is an ionic bond, we'll say it fully gains two electrons. So we'll have an oxidation state of negative 2. And once again, the notation, when you do the superscript notation for oxidation states and ionic charge, you write the sign after the number. And this is just the convention. And now, with these two examples, the whole point of it is to show that hydrogen could have a negative 1 or a positive 1 oxidation state. But there's also something interesting going on here. Notice, the oxidation states of the molecules here, they add up to the whole-- or the oxidation state of each of the atoms in a molecule, they add up to the entire charge of the molecule. So if you add a positive 1 plus negative 1, you get 0. And that makes sense because the entire molecule lithium hydride is neutral. It has no charge. Similarly, hydrogen, plus 1 oxidation state; oxygen, negative 2 oxidation number or oxidation state-- you add those two together, you have a negative 1 total charge for the hydroxide anion, which is exactly the charge that we have right over there.