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Chemistry: Oxidation Numbers

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  • Type: Video Tutorial
  • Length: 13:55
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 150 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Final Exam Test Prep and Review (49 lessons, $64.35)
Chemistry: Looking In-Depth at Redox Reactions (5 lessons, $7.92)

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 http://www.thinkwell.com/student/product/chemistry. 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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Thinkwell
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Want to see the real power of a chemical reaction? We're going to give our little friend here the ride of his life. You just saw an example of a reaction that we call a combustion reaction. It also could be classified as a redox reaction in that electrons transferred from one compound to another. If we look at the formula, the chemical equation describing what just happened, it is not apparent where electrons are coming from and going to. Everything is neutral. There are no charges giving us any hints as to where the electrons are.
What we're missing to be able to analyze a reaction like this is a tool called oxidation states or oxidation numbers. Oxidation states or numbers attempt to take account of where electrons are and assign them to individual atoms. Now we know right away that this is going to get us into trouble if we take this too literally, because we learned already that in molecules electrons are in molecular orbitals, not just in atomic orbitals so electrons can't conveniently be assigned to one atom or another. Nonetheless, as long as we're staying honest with each other about what reality is, it still is a valuable tool to be able to assign electrons to individual atoms as long as we don't take that completely literally.
Let's look at an example here. Consider the molecule hydrogen chloride. We know that this bond is a very polarized bond. We know that the electronegativity of chlorine is much greater than hydrogen. The electrons involved in this bond are pulled very much toward the chlorine. We've talked about this molecule as having a dipole moment where, again, there is a partial separation of charge. The chlorine is more negative. Hydrogen is more positive. The notion of an oxidation number is to say, okay, let's take those two electrons and formally, we'll assign those completely to chlorine. The electrons involved in bonds go entirely to one atom or another atom. We remove completely this notion of sharing electrons. Again, we know that's wrong. We know that that is purely a formalism, but it's a form of bookkeeping, just making sure that we keep track of where all electrons are. We would assign an oxidation number of +1 to the hydrogen, and an oxidation number of -1 to the chloride.
Right away, let me show you where we can get in trouble with this. This is not too far from reality. We know that hydrogen chloride, when it's dissolved in water is hydrochloric acid, and it easily breaks up into these ions. This is not very far from the truth. However, consider the molecule methane. In the case of methane, the difference in electronegativity between carbon and hydrogen is very slight. Yet, by this formalism, we'll go ahead and describe carbon as having a 4- charge or, in other words, having a 4- oxidation number, and hydrogen as having a 1+ oxidation number. If we take this literally, we'd assume that methane was very acidic and we know for a fact that this is not true. Again, just keep in the back of your mind as we go through the rest of this unit, that oxidation number is only a formalism, but it gives us some insight, in particular, for instance, by knowing oxidation numbers for the hydrogen and chlorine, it gave us insight about where the electron-rich portion and where the electron-deficient portion of the molecule was.
Finally, while we're here, let's consider something where we're talking about two atoms that have the same electronegativity. What would be the oxidation number in dichlorine? Well, in this case, there is just no way that we can describe those electrons as being more in one side than the other. So we agree by convention to call the oxidation number in this molecule "0". One of our first rules that we'll encounter is that elements, in their elemental forms, will have an oxidation number of 0.
Let's go through these rules. Oxidation numbers must add up to the total charge. That's a different way, maybe a more general way, of saying what I just said before. If we have a neutral element, then the oxidation number of that element must be 0. But if we have an ion, then we have to make sure that the overall sum of oxidation numbers is, in fact, equal to that ion. Or, if it's a neutral molecule with different atoms, the sum of the oxidation numbers again must equal whatever the overall charge is - in the case of a neutral, 0.
Alkaline metals, we know by experimental evidence, almost always takes the 1+ oxidation state in that almost always it forms one-to-one salt, as we've talked about with halides, for instance. As our next rule, we will always assign the alkaline metals an oxidation state of +1, the only exception being, if it is the element itself. So sodium metal will have an oxidation state of 0, but otherwise, sodium in the form of salt is going to carry an oxidation state or oxidation number of 1. By the way, regarding those two terms - oxidation number is a little more modern nomenclature, oxidation state was the older nomenclature. Alkaline earth, we'll go ahead and assign an oxidation state or oxidation number of 2+. Hydrogen will normally assign an oxidation state of +1, but there's a very important exception. Remember that hydrogen is capable of accepting electrons to reach a filled shell configuration as well as losing electrons. We have to pay a little bit of attention here. Hydrogen normally will assign an oxidation state of []+1, but the important exception being, when it is connected to a more electropositive element such as alkaline metals, for instance, or alkaline earth metals which are very electropositive and tend to release electrons better than hydrogen does. In those cases we would assign the hydrogen and oxidation state, an oxidation number of -1 rather than +1. Of course, if it's elemental hydrogen it would have an oxygen state of 0.
Oxygen we'll normally find in the oxidation state of 2-, once again with a couple of important exceptions. If oxygen is connected to itself in the form of neutral O[2], we know that that oxidation number is going to be 0. If it is connected to a more electronegative element than itself, and there really is only one we have to worry about, then oxygen actually could have a different oxidation state. It's very rare when that would happen. In fact, there are probably only one or two compounds known that oxygen would have a positive oxidation state like that. That would be the rare exception. Normally we're going to see it as a -2 or sometimes a -1 in the form of a peroxide. We'll see that's when we have an oxygen/oxygen single bond or, most commonly, the 2- or the element itself, 0.
Finally, kind of the catch-all last rule is that when we have other elements that don't neatly fit into this scheme, we're going to go ahead and allow the element that is more electronegative to receive as many electrons as it wants to reach a noble gas configuration. For instance, if it was nitrogen bound to something that is less electronegative, we would assign it a formal oxidation state of -3, or a formal oxidation number of -3, to allow it to reach the noble gas configuration. If you look at the graphic to the side here we illustrate that point.
What remains is for us to do a few examples. Let's start out with the nitrate ion, NO[3], and walk through how we're going to do this. In this case, going through our rules, the first rule that we come to is that we assign oxygen an oxidation state of -2. There are a total of three oxygens in this case. That would give us a total charge, just taking into account the oxygens of -6. The nitrate anion overall has got a charge of -1. Of that -6 charge, we still need to account for a +5 in order to end up with an overall charge of -1, thus the nitrogen will be assigned, in this case, an oxidation state of +5.
Let's consider another common nitrogen-containing ion, the ammonium ion, NH[4]. In the case of ammonium, once again going through our rules, we come to the fact that hydrogen normally will be in a +1 oxidation state. So we have hydrogen in a +1 oxidation state but there are 4 of them, so that gives us a total of 4+, just considering hydrogen. We know it has an overall charge of +1. That would require the nitrogen, then, to have a charge of -3 in order to give us that total. So we're seeing the whole range of oxidation states for nitrogen, ranging from -3 all the way to +5. Once again, like in the case of methane, this does not promise that nitrogen has really got a +5 charge, but what it does tell us is that the nitrogen is probably electron deficient in this molecule. Formally, it is missing a lot of its electrons.
Let's look at an example now where we have two of the same elements. We'll do dichromate, Cr[2]0[7]. This has an overall charge of 2-. Once again, in the case of oxygen, we know that that's -2 x 7 in this case, a -14 charge. We have an overall 2- charge so what we're missing here, apparently is a +12. The +12 now corresponds to both chromiums. Right? If both chromiums account for +12 charge units, then we have the correct charge overall for the anion. In this case we'd take 12 divided by 2 chromiums and we'd end up with a +6 charge or oxidation number for the chromium.
Let's look for a moment at oxidation states in organic compounds. We'll consider now a series of one carbon-containing organic compounds. We'll start with carbon dioxide. In the case of carbon dioxide, we have a total of -2 x 2 atoms. That's going to give us, then, carbons. Since this is neutral, we're going to have an overall +4 oxidation state in the case of carbon dioxide. If we go now to an example of an aldehyde, and I'm going to draw the way an organic chemist would draw it, formaldehyde, CH[2]O, and then we'll just redraw it, simply keeping track of atoms in this case. Once again, just taking account of the oxidation states of the more common elements, oxygen is going to be -2. Each hydrogen is going to be +1. In this case we end up with a balancing of the charge, and so that gives us an oxidation state for this aldehyde, formaldehyde in this case, of 0. Even though it is not in its elemental form, it still can have an oxidation state of 0.
Finally, let's look at an example of carbon with a very negative oxidation state. We'll look at the oxidation state, in this case, of methane. As we've described before, hydrogen must have, overall, +1 charge times 4, which would give us an oxidation state of -4 for the carbon. We see that in these neutral organic compounds, we are going from a +4 oxidation state all the way down to a -4 oxidation state. It doesn't mean that carbon is missing four electrons or have picked up four extra electrons. But what it does tell us is that we're going from electron-deficient to electron rich in this process. That will start to give us insight into how molecules behave in chemical reactions.
Oxidation-Reduction Reactions
Looking in-Depth at Redox Reactions
Oxidation Numbers Page [2 of 3]

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