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Chemistry: Reviewing Oxidation-Reduction Reactions


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  • Type: Video Tutorial
  • Length: 10:16
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 110 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Electrochemistry (12 lessons, $19.80)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

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I'd like to remind you about two kinds of reactions we've seen before. In the beaker here, I have a solution that started out as silver nitrate. And then we put a piece of copper wire in it. And what you're seeing is the formation of silver crystals - and a lot of silver has already dropped down to the bottom of the beaker. And the blue color, that you're seeing, is copper two plus, or cupric ion, so that the chemical reaction that I'm showing is right here. It's very exothermic, very spontaneous reaction, in this case. And it's continuing to go here.
The reaction down below is an example of a combustion reaction. This happens to be a little bit of methane and oxygen and we'll just supply a spark. So this is a very exothermic reaction and a very spontaneous reaction. These are very similar in one level but they're both examples of redox reactions - at least in a formal sense - in that we can identify a reductant and an oxidant.
So let's get a little review here on what redox reactions are. In the above reaction, copper is being oxidized to copper 2 plus. Remember, oxidation is removing electrons whereas reduction is adding electrons. The word "reduction" comes from the sense that we're reducing charge, but we're actually adding electrons to it. And in the same sense, formally at least, we could describe oxidation states for the components down here. We see that carbon is going through an oxidation state change, in this case, of minus 4 to plus 4 and in the oxygen's case, it's going from zero to minus 2, in both the CO[2 ]and the water form. So we can talk about 8 electrons being transferred, overall, coming from the oxygen. I forgot, I should just say, in this case, methane, we'd be describing as the reductant, and the oxygen, here of course, would be described as the oxidant - so, very, very similar reactions formally. Wouldn't it be wonderful if we could somehow just directly, on a microscopic level, connect a wire to something somehow, to get these electrons to transfer through the wire rather than directly in solution? If we could do that, we could use that wire with the electrons transferring through it, and we could maybe harness the driving force of this reaction and get some useful energy out of it. Likewise, it would be wonderful if we could do that for the bottom reaction.
Well, it turns out that one of these reactions is much more readily available for that type of a procedure than the other one is. In this case, the electrons are transferred with a minimum amount of rearrangement of other atoms. And so the kinetic barrier for this type of a process is very low, and it happened without doing anything. It's just happening before our eyes. The kinetic barrier for this process is extremely high, and so as a result, this does not happen on any kind of a reasonable time scale. Remember, we had to apply a spark. And by the time we did that, there were literally 10 or 20 or 30 different reactions happening - and maybe hundreds - going on at the same time - just complete chaos at that high temperature, of radicals being sent all over the place, components, pieces of fragments of molecules. So the overall reaction delivered plenty of energy, but not in a very controlled manner. And again, in order to get anything going, we had to get over a very high barrier, whereas up here, the barrier was very low. So it turns out that not all reactions that, at least in a formal sense, undergo a redox process - in other words, a change of oxidation states - are readily available for harnessing electricity from, but some are - some that involve very minor changes. And so what we're going to focus on in this unit is how, in fact, one goes about taking what we'll refer to as half-reactions, or even half-cells, eventually, as in a battery cell. And, in this case, the copper going to copper 2 plus would be a half-cell. And the silver plus going to silver metal would be a half-cell. And how we can actually physically separate those two ideas and get useful electricity - electrical energy - out of that kind of system.
Now before we go on, what we need to do is just remind ourselves a little bit about how we assign oxidation states - how we balance redox reactions - and then we'll go ahead and start to look at what an electro-chemical cell looks like. So, let's just quickly remind ourselves about oxidation state rules, because if we can't figure out where electrons are coming from and going to, well then, we're up a creek. Basically, the idea here is those most electro-negative elements are going to get their way. They're going to get, in a formal sense, all the electrons that they can to fill their valence shell. So alkali metals, in this case, are going to be the ones taken advantage of. They're almost always going to be in an oxidation state of plus 1, unless they're all by themselves in their elemental form. Alkaline earths normally will have an oxidation state of plus 2. Hydrogen normally will have an oxidation state of plus 1. With the exception of metal hydrides, it would be more electro-negative than some of the alkaline metals. Oxygen almost always will have an oxidation state of minus 2. The exception is if it's connected to itself or if it's connected to a more electro-negative element, such as fluorine. And then finally, if none of these rules apply, then if there's a bond, the electrons in that bond, in a formal sense, will go to the element that needs them the most, that are the most electro-negative.
So let's look at balancing a redox reaction again. This is a typical example of something you might see on an exam. And there are a couple of different strategies that we've discussed. We will summarize the half-cell method. I will summarize the half-cell method here, and that's simply to look at a redox reaction, identify the redox partners - in this case bromine going from neutral to a 1 minus and sulfur going from, in this case, a plus 4 to a plus 6. Those are the changes in redox chemistry here. And so, I'm going to break those ideas down as two half reactions - the bromine taking two electrons to go to two bromides. In this case, again, think about what we're doing here. The bromine is acting as an oxidant. This is a reduction step, though, in that the electrons are going to the bromine. So that's the reduction step. And down here, the SO[2] is going to SO[4]. It's increasing its oxidation state. It's giving up electrons in the process. So this would be the reduction step.
These are not balanced, so our next step is to balance oxygens, for instance. In this case, this one is balanced. But this guy needs his oxygen balanced. That would be our next step, to balance our half-reactions. And so, that's going to require - I'm in acid in this particular problem, which you may have seen. And so, I need to adjust for oxygens and hydrogens in such a way that it's consistent with being in acid. So I'm not going to use hydroxides here. I'm going to just use H plus and water on each side of the equation as I need. So in this case, I'm going to need to have 2 waters on this side to have the oxygens completely balanced, the 4 oxygens over here. That, in turn, is going to require 4 hydrogens that need to show up on this side in order to have all atoms balance. And then, I'll just check myself, make sure that all my charges balance on both sides. Everything looks good. So once that all looks good, we'll go ahead and just add the two half-reactions together, but making sure that the electrons gain as the electrons receive. We want to make sure - this is probably the most crucial step - that when we accept 2 electrons up here that those same number of electrons come from the bottom reaction. So those two numbers have to be same. In this case, they are. If they weren't, we'd multiply each through by a multiplier to make sure that they were. Then we'll add them together and end up with a balanced redox reaction. So, if that seemed a little rushed to you or you're a little foggy on that, I suggest you go back to an earlier tutorial where we go through, in a little more detail, strategies for balancing redox reactions.
Okay, so we're all set then. Let's go back to our idea of copper and, in fact, let's go ahead and look at that equation again. Let's go back to the idea of copper combining with silver to give us copper 2 plus and silver metal. And we're ready to go. What we're going to do, to kind of keep things as simple as we can, is let's take everything in their standard states. Okay, that means just so I don't have to worry about huge changes in concentration, I'm going to go ahead and start with copper and copper 2 plus in their standard state. So one molar concentration of copper 2 plus, copper metal. It is in its standard state. It's a metal. And then, we're going to have one molar concentration of the silver. And silver metal, again, will be silver metal. We'll go ahead and make up the half reactions and what we'll call half-cells. This is the piece of copper and we're going to go ahead and put it in one molar copper nitrate. And then, this is a piece of silver and we'll put it in one molar solution of silver nitrate, and we're ready to go. We'll just connect up those two electrodes together and hopefully, we'll get some work out of it. And you know what? It doesn't work. So we'll let you think about that for a couple minutes. In our next tutorial, we'll talk about the missing piece that will make this work by way of introducing electro-chemical cells.
Principles of Electrochemistry
Oxidation-Reduction Reactions Page [1 of 2]

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