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Chemistry: Properties of Transition Metals


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  • Type: Video Tutorial
  • Length: 13:28
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 144 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Transition Metals (9 lessons, $14.85)
Chemistry: Examining Transition Metals (2 lessons, $3.96)

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.

About this Author

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In the heart of the periodic table lies a set of elements called transition metals. Transition metals are very close to me. My graduate work had a lot to do with transition metals and my group, at the University of Virginia, now is doing research using transition metals to do organic reactions with.
Now, what characterizes transition metals - we can talk about these in several different contexts. But from an electronic standpoint, the reason they show up in this part of the periodic table is, of course, they're electron configurations. And what characterizes all transition metals is that they all involve partially filled d-orbitals in their valence shell. Now sometimes, the very last column of the transition metals - zinc, cadmium and mercury - these are also sometimes lumped into the transition metals. But technically, they actually are not transition metals in that these guys have filled d shells. But all of the rest of these guys showing up in orange, you'll recall, have partially filled d-shells. And that gives some of their chemical properties, which we'll talk a lot more about in a moment.
Some basic characterization - some basic properties - of the transition metals, these are all metals, in that they have valence electrons that are delocalized. These are typically harder elements, in their elemental form - harder metals than the metals associated with the main group elements. They're higher melting. Tungsten, for instance, is what's used in light bulbs - very, very high melting - 3000° - very, very high melting metal. These materials, when they make compounds, are very often colored. And that's one of the attractive features of working with these things, from my standpoint, just beautiful colors of compounds. And we'll talk a lot more about where that color comes from and why it's so particular to transition metals, what types of electronic transitions gives rise to these colors that you can have in transition metals that you don't have, necessarily, with the main group elements. So we'll have a lot more to say about that.
Likewise, the magnetic properties are really interesting for these compounds. And again, these have to do with these partially filled d-orbitals. So, once again, we'll talk about magnetic properties, as well as the optical properties, or the colors that come from these things.
So let's talk a little bit more about the periodicity of the transition metals and how that compares to the main group. We know about the trends in the periodic table, as far as size, as far as ionization energy and so forth. What about within the transition metals? Well, one thing that we see that's different is that if you look at the size of the elements, remember the general trend, in the periodic table, is as we go from top to bottom in the periodic table, the size increases. Well, not so, as it turns out, for transition metals, especially if we compare the second and the third rows here, these guys have actually identical sizes, almost - almost absolutely identical, within error, at least. And this is a phenomenon referred to as lanthanide contraction. Basically, what's going on here is that is remember that this is actually kind of a distorted picture of a periodic table in that for convenience - for logistics purposes really - we've removed the rare earths from this section right here. And so, in fact, we take a big jump between barium and luticium as far as the charge of the nucleus. And so this buildup of positive charge in the nucleus is what's responsible here for actually causing the third row to be pulled in to the point that it's essentially the size of the second row in the periodic table. Now that has another very immediate consequence. Think about what that means as far as how easy it is to remove an electron from these things. The ionization energy is going to also mirror that idea in that the ionization energy for the third row actually is higher than the ionization energy of the second row. And one of the consequences, although this is a broad overgeneralization, is that the third row is not chemically as reactive as the second row, because it's more difficult to get those electrons away from the nucleus. Again, that's a very broad oversimplification, but certainly a consequence of this lanthanide contraction.
So those are the general chemical properties that we'll be looking at. The important things here are going to be the spaces in the d orbitals. And we'll see the profound consequences of that shortly, as far as the unique chemistry that we get out of the transition metals.
So let's go ahead and what we'll do is look at the first period of the transition metals and look at a little more detail at their electron configurations, and basically just come back and reemphasize some of the points we made just a moment ago. If we look at that first period and we look at the electron configuration, I remind you that the first thing that happens is that we put a pair of electrons in the 4s orbital. That happens actually prior to the 3d orbitals filling. And so we have this peculiar situation of the 4s and 3d orbitals being very close together in energy. And so in their neutral forms, we actually have 2 s electrons in the 4s orbital and then 1 d electron, 2 d electrons, so on. And notice that there is an exception here. We take out an electron of the s orbital to give us a half-filled shell, in the case of chromium. And then we continue to put electrons in. And again, we see an exception here at copper. These aren't so important to remember, at least in my opinion, because this is only pertaining to the atoms in their gaseous elemental form. And as soon as you put them in compounds, the ordering changes a bit, as we've talked about earlier.
But I do want to point out, about electron configurations, is that the oxidation states that are common for these elements directly reflects the electron configuration, as would be expected. So the oxidation state for scandium is 3. We've got 3 valence electrons. For titanium, it can be 2, 3 or 4. And notice, we've got 4 electrons in the valence shell. For vanadium we have 5 electrons total in the valence shell and we have oxidation states ranging from 2 all the way up through 5. And this is one of the interesting properties of transition metals is that, unlike main group elements, they enjoy a wide range of different oxidation states. And as we'll see in a little while, they can easily change from one oxidation state to another. So transition metals very often undergo 1 electron transfer, which is not nearly as common for main group elements. And we'll see, again, this plays a really important role when we look at enzymes in your body and we look at specific characteristics of transition metals and how they differ from main group elements. So the one other thing I point out on this chart is look what happens to density as we go across the periodic table. We increase effective nuclear charge. We know about that as a general periodic trend. So not only that, we're increasing the mass of the nucleus and so the density is reflecting that. The density is steadily climbing as we're going across that period.
Okay, so now that we understand some of the basic properties of the transition metals in their elemental forms, let's talk for just a moment about where they come from. Transition elements are relatively electropositive, meaning it's very easy for them to be oxidized by oxygen, for instance. And so you don't find the transition metals in their elemental forms in nature. These guys are going o be in high oxidation states. They'll typically be found as oxides, like rust, for instance, sulfides, halides in the form of different minerals, carbonates.
And so the challenge to the chemist or to the engineer is to get the transition metal in its elemental form, and you need to be able to reduce the element, to go from these high oxidation states back to the element. And so, if you think about the development of chemistry and materials and so on, it turns out that copper, as we've discussed earlier, is easier to reduce than iron. Now copper and iron both have fairly high abundances compared to some of the other transition metals. Copper is a lot lower abundance than iron - about 1000 times less abundant than iron - but copper is a lot easier to get to, because you can reduce it a lot easier. You don't need as high of a potential.
And so, historically, the first thing to develop was copper, and then we'll talk about iron in a moment. Minerals that contain copper or iron, just in general, before I talk more about copper here. The first step in the processing is to just get the minerals in their pure form. And so you go and you dig a huge hole in the ground. There's a hole. You're digging out all this stuff. You get your minerals separated from all the mud and what's called gangue. And then the trick is just converting that mineral into the elemental form. Now there are two industries that have built up over this - hydrometallurgy and pyrometallurgy. This has to processing from water. This has to do with processing at high temperatures.
Now, we really won't say a lot about this, but let me just give you a flavor for what kinds of things you're going to do. If you're interested in making copper, you've got to be able to reduce the copper from its mineral form. So various different sulfides of copper, which are common minerals containing copper, these guys actually can be reduced to copper metal with oxygen. Now that seems a little bit odd that you're using an oxidant to do this, but in fact the oxidant is serving the role to oxidize the sulfide here. And the copper is reduced in that process. So this is a way to just get us down to the elemental form of copper. Copper can be used in alloys, such as bronze. That's an example of an early alloy that was used extensively.
But things got a lot better when people figured out a way to reduce iron, because there's a lot more iron around and iron, in general, is going to be a lot stronger than copper. So the iron age was a couple of thousand of years later, so it took folks a little while to figure out how to do it, but now that we know how to do it, iron is certainly the metal of choice, because it's abundance is so high.
You can get to elemental iron by reducing it with carbon monoxide. The carbon monoxide comes from taking coal and oxidizing the coal under starving conditions, meaning you're not going to give it an unlimited supply of oxygen, in which case you would have gotten carbon dioxide. But instead, you're starving it for oxygen, so you only partially oxidize the carbon here up to carbon monoxide and then that acts as your reducing agent then, when you throw that in with what's effectively rust - an iron oxide. And that gives you iron in its metallic form - in its elemental form - along with carbon dioxide. So the most common use for iron is steel, certainly. We use it structurally all over the place. That's actually a mixture - an alloy - of iron and other transition metals. So stainless steel, for instance, has actually has got about 20 percent chromium mixed in with the iron, and a little bit of nickel. And depending on compositions of the transition metals that you mix in with your iron, you get all kinds of different properties of your metals, as far as strength and malleability and so forth. So we certainly are familiar with iron in our world.
Another property we associate with iron is its magnetism, more correctly, it's ferrimagnetism - and we'll have a lot more to say about that in a moment, but just to remind you about ferrimagnets, you can see here the ability of a magnet to pick up something else that's magnetized. And so you can have lots of fun playing with these things. But anyway, we'll talk about what it is that gives rise to this unusual property of magnetism here. And we'll understand that only by looking at the molecular or atomic level for these compounds. So that's the next step. We're going to now leave the elemental forms and focus on transition metals as they show up in compounds.
Transition Elements
Transition Metals
Properties of Transition Metals Page [3 of 3]

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