Learn about Senior Chemistry, Redox 3, in this comprehensive video by bannanaiscool.
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Now you may say, “That oxidation number stuff I understand I get it, but is there another way to balance this reaction successfully without using those numbers and stuff?” Okay, there is. Now that doesn’t mean that this method which sometimes I do prefer, it doesn’t mean it takes the place of understanding oxidation numbers because you may be just asked in a question, what substance here in this reaction it’s always in the reactant side what’s undergoing oxidation? By using oxidation numbers you can easily determine something that’s undergoing oxidation or reduction. But this is called the half reaction method where you’d break this equation down into two half reactions and then you build your own half reactions. And watch how this is done. The rules are as we go through I will try to explain. What we’re going to do is we’re going to take two principal chemicals in here and break them down into half reactions. Well look at this MnO4 negative here has to somehow turn in to Mn2 positive. So we write MnO4 negative turns into Mn2 positive. Hydrogen and oxygen are not necessarily the principals in a reaction. It’s the other elements that look kind of different from reaction to reaction that we actually split up. Fe has to turn from 2+ to 3+. With that we’ve got two half reaction now that we’re going to add together to get this net reaction. But it doesn’t look like—where is this going to go? Make your own and watch how you do that. When you need oxygen you add water. When you need hydrogen you add H+ and then in order to balance the charges you add electrons. So follow those three rules and everything works out great. Here is how it goes. I have a half reaction here that doesn’t have oxygen on this side. When you need oxygen add water. Oh yeah, wait a minute there are four of those here. I’ll put 4H20 in front. Now you’ve got the oxygen balanced. Now I've got hydrogen on this side but not over here. When you need hydrogen add H positive. Now you’ve got to balance the charges. One negative here—but there isn’t one negative here, there are eight H’s. You’ve got to be careful. You’ve got to go through this and make sure that you catch yourself like I just did in terms of the number of atoms you have all time. That’s eight positive and one negative is seven positive on the reactant side, but only two positives here on the product side. So what do you do? You actually have to add negatives to which side? To this side over here and if you have five negatives, six negatives, eight positives makes two positive total on this side and two positive total on this side. So just be careful and do that and you'll get it. Two positives here, three positives here you don’t have to add hydrogen or oxygen what do you do? Well we add one electron here. And hey, we’re ready to add them together. But look at just like we discovered with the oxidation number balancing of this. You’ve got one electron here and five here. And so you take this whole half reaction and you multiply it by five, so you get five electrons here canceling five there and electrons have to cancel out for the net reaction completely and then of course you get the eight—well if we add these two reactions together we’re going to get the eight in front here, the five here, the five here and the four here that we did in our last example. So this is a reduction half reaction, oxidation half reaction. The two are added together to give you this net redox reaction. So let’s practice another one. But now we have to balance in a base.