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# Redox reaction with iron

Video transcript
So here, we have an interesting reaction. We have some iron with some hydrochloric acid, if we are assuming it's in an aqueous solution. And as a product, we get iron II chloride. And you'll see in a second why it's called iron II chloride. And you get some molecular hydrogen. And what I encourage you to do right now is pause this video, figure out the oxidation states for each of the components of these molecules before the reaction and after the reaction, and then try to write your own half reactions and determine who has been oxidized and who has been reduced. So I'm assuming you've given a go at it. So let's think about it. So right here, this iron, it's not bonding to anything. So there's no reason that its going to gain electrons or lose electrons in our hypothetical oxidation state world. So it has an oxidation number of 0. Now, you have hydrogen bonded to chlorine. Chlorine is very electronegative, hydrogen not so much. So chlorine, in our hypothetical world, is going to take the electrons. So it's going to have an oxidation number of negative 1. And the hydrogen is going to, in our hypothetical world, lose the electron. And so you have a positive 1 oxidation number. We could write 1 plus here if we wanted or 1 minus there if we want on the chlorine. Now let's go on to the other side of the reaction. So our iron now is bonded to two chlorines. Now let's just think about what chlorine typically does. We've already seen chlorine. It's Group 7 element right over here. It would love nothing more than to gain an electron so that it has a complete eight electrons in its valence electron shell. And so it typically has an oxidation number when it's bonded with other things of negative 1. Since I'm not writing it as a superscript, I could just write it as negative 1. So let's go with the flow there. Let's assume that each of these chlorines are taking an electron. So it has an oxidation number or oxidation state of negative 1. Well, if each of these have an oxidation state of negative 1, and this whole thing, this iron II chloride is a neutral molecule, then the iron is going to have to-- you'll see 2 times negative 1 is negative 2. The iron, in order to net out, is going to have to have a positive 2 oxidation number. And that is why this is indeed called iron II chloride. The iron II tells us that the oxidation number on the iron is positive 2, iron II chloride. And what about the hydrogens here? Well, hydrogen bonded to hydrogen. There's no reason why one hydrogen would hog it from the other hydrogen or let their electrons be hogged by another hydrogen. And so, each of these hydrogens would have a oxidation state of 0. Now, let's think about the half reactions, decide who's getting oxidized and who's getting reduced. So let's start with the iron. So the iron on the left hand side, 0 oxidation number. And then on the right hand side, iron has a positive 2 oxidation number. So how do you go from 0 to positive? Well, you've got to lose electrons. Every time you lose an electron, your oxidation number, your hypothetical charge-- or if you're really losing electrons, your real charge-- is going to go up. It's going to become more positive. So this character right over here, he is losing electrons. So we could say this right over-- he has lost electrons. So let's throw those electrons. Here he hypothetically had them. Here he lost them. Where are those electrons? They are right over here. So what do we call it when you are losing electrons? Well, we call that oxidation. What is being done to you is what oxygen typically does to things. It takes electrons away. Even though there's no oxygen here, this oxygen typically takes electrons away. He's having his electrons taken away from him. So he is being oxidized. So the iron is being oxidized. Now, let's think about the hydrogens. On the left-hand side, you have two hydrogens with a plus 1 oxidation number. Now, on the right-hand side, you have still two hydrogens. Otherwise, it wouldn't be a balanced equation. You have two hydrogens. But what's their oxidation number now? It is now is 0. So how do you go from having a positive oxidation number to a neutral one? So you must have gained electrons. That's the only way to bring your charge, your hypothetical charge or your real charge down, is to gain electrons. So this guy right over here, each of the hydrogens must have gained an electron. There's two of them, so collectively they must have gained two electrons. So what do we call it when you're gaining electrons in these reactions? Well, reduction is gaining-- RIG. Hydrogen is being reduced here. Let me write that down. So hydrogen is being reduced. Now what about the chlorine? Chlorine has a negative 1 oxidation state on the left-hand side. And there's two of them. And there's two chlorines on the right-hand side. But they still each have a negative 1 oxidation state. So the chlorine was neither oxidized nor reduced in this reaction. But this is definitely an oxidation-reduction reaction, or a redox reaction, because you had something getting oxidized and something being reduced. Iron was oxidized by the hydrogen, and the hydrogen was reduced by the iron.