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Studying for a test? Prepare with these 6 lessons on Redox reactions and electrochemistry.
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Now that we know a little bit about oxidation and reduction, what I want to do is really just do an exercise to just make sure that we can at least give our best shot at figuring out the oxidation states for the constituent atoms that make up a compound. So, for example, here I have magnesium oxide, which is used in cement. It has other applications. And this is magnesium hydroxide, which is actually used in antacids. It's used in deodorant. And what I want you to think about, and I encourage you to pause this video right now, is given these two molecules, these two compounds, and what we know about the periodic table, try to come up with the oxidation states for the different elements in each of these compounds. So I'm assuming that you've given a go at it. Now let's try to work through this or think through this together. So first of all, magnesium. Magnesium right over here. We see it's group two. It's an alkaline earth metal. It has two valence electrons. It's not that electronegative. We've already seen that something in this group right over here with two valence electrons, it's likely to give them away. So if it were to form ionic bonds, or if it were to be ionized, it's likely to lose two electrons. If you lose two electrons, you would have a plus 2 charge. So magnesium would typically have a plus 2 oxidation state. On the other side of the periodic table, oxygen, group seven. It has six valence electrons. It's very electronegative, so electronegative that oxidation is named for it. It likes to take electrons from other elements. And oxygen in particular likes to take two electrons. So it's not unusual to see, actually anything in this group, but especially oxygen, taking two electrons from something else. If you take two electrons, and you started off neutrally, or you started in a neutral state, it's not unusual to see oxygen at a negative 2 oxidation state. So given that, it seems like this could work out. Magnesium could have a positive 2 oxidation state. And actually when you write it as a superscript here, the convention is to write the positive after the 2. And oxygen would have or could have a negative 2 oxidation state. And this makes sense relative to the overall charge of the molecule. Positive 2 plus negative 2 is going to be 0. And that makes sense. This thing overall is a neutral molecule. And not only in this case is the oxidation state a hypothetical ionic charge, if these were to be ionic bonds, this actually is an ionic compound. Oxygen actually does take two electrons. And magnesium actually does give away two electrons. So in this case the oxidation state is actually describing what is happening ionically. Now let's think about this one right over here, magnesium hydroxide. Well, just like before, magnesium typically has an oxidation state, likes to give away its electrons. So it could have an oxidation state of positive 2, which would imply that the entire hydroxide anion-- And let's just say hydroxide for now. Well I'll say hydroxide anion. I kind of gave it away a little bit-- that this hydroxide, or this part of the molecule, the right-hand part of what I've written here, for this whole thing to be neutral, it should have a negative 2 oxidation state. Now how does that make sense? Well we have two hydroxides here. Notice this subscript right over here. So if each of those hydroxides has a negative 1 charge, or a negative 1, I guess you could say, total oxidation state, then when you take two of them together, they would net out against the magnesium. And that does seem to make sense. If oxygen has a negative 2 oxidation state, hydrogen has a positive 1 oxidation state. Each hydroxide part of this molecule is going to have a net oxidation state of negative 1. But then you have two of them. So the net oxidation for this part of the molecule or the compound is going to be negative 2 nets out with the positive 2 from magnesium. So once again, it makes sense.