Chemistry: Color and Transition Metals

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I mentioned to you earlier that I have a real soft spot in my heart for transition metals. I do that in my research. I did that in my graduate work. But actually it was back in high school when I first fell in love with these types of compounds. And it was the rich colors that you could get out of these materials that fascinated me so much.
Think about this. One transition metal, let's consider chromium for a moment, one transition metal, depending on what's mixed in with it, gives rise to a whole spectrum of different colors, just beautiful colors. And that's very uncharacteristic of the main group elements. So what we're going to talk about now is where this color comes from, and how we can start to make sense on a molecular level of the origin of the absorptions that give rise to these colors.
Let's start out with a curious observation. And that is I've got a solution here of hexafluora cobalt[ ]three. So this is the anion in this solution. And you'll see that it's this beautiful deep green color. And if I consider breaking this down into the metal and the ligands, this is again a cobalt in a three plus oxidation state. Now this also is cobalt in a three plus oxidation state. This compound is cobalt hexamine, or hexamine cobalt[ ]three. And again it's the exact same metal. It's the same oxidation state. And in fact, in the metal it's exactly the same electron configuration. Yet they are clearly very different as far as our observation goes. Their colors are vastly different. And in addition to the color, their magnetic properties are different. This guy turns out to be paramagnetic. In other words, when you bring it into a magnetic field there's an attraction to that magnetic field. This is diamagnetic. You bring it into a magnetic field and there's a very weak repulsion. And we'll have a lot more to say about magnetism a little later. But what these observations tell us, I'll just tell you this right now, is that when it's paramagnetic, that tells us that there are unpaired electrons in the molecule. And if it's diamagnetic, there are no unpaired electrons in the molecule. So something is very different. Given that we have an equal number of electrons, why are they pairing in one case and not in the other? Why are their colors so different?
So let's start with what is the electron configuration for cobalt three plus and see where that leads us. So why cobalt three plus, again I remove the ligands, I end up with just the metal in the correct oxidation state to account for the overall charge for the ions. That's cobalt three plus. If it were cobalt neutral, we would count a total of nine electrons, two in the 4S orbital and then seven in the 3D orbital. Now we want to remove three electrons from that. Now something curious happens. Let me remind you of why the 4S orbital is filled first. And we can look, in fact I'll highlight it for you, look at what happens to the 4S orbitals as we go across the Periodic Table, and you'll notice that they drop in energy here as a result of the increasing nuclear charge. But there's a lot of shielding going on in the 3D orbitals, much more so than in the 4S. So the 3D doesn't drop as fast, and actually crosses over the 4S right about the time we start to fill the transition metals. So the 4S orbital actually fills before the 3D, because it's a little lower in energy. But that is not true anymore once we start to take out electrons.
When we take out electrons, we generate a cation. The cation has less shielding, so this crossover doesn't occur anymore. So we take electrons out of the 4S before taking them out of the 3D. Thus for this cobalt example, although this is the neutral configuration with the 4S below the 3D, and so we put in electrons there first. That's the total nine electrons. When we remove three electrons, we end up now with this configuration, with the 3D below the energy of the 4S now. So that's where the electrons stay. And so we have a net of six electrons in the 3D orbitals for either of these complexes.
Well what does that have to do with color? Well just hold on to that idea, that we have six electrons in the d orbitals. Now the observation is that hexamine is diamagnetic, meaning no unpaired electrons, and the hexafluoride is paramagnetic. So that means that there are unpaired electrons. Well if all of the d orbitals were exactly the same energy still, like we would expect for a neutral atom, well then we could explain why it's paramagnetic. Because of Hund's rule, all of these guys are going to be unpaired. Remember repulsion? They're going to try to avoid pairing up because of electron-electron repulsion. And our prediction would be that the compound should be paramagnetic.
On the other hand, maybe we could explain why the compound was diamagnetic if we said, okay, we know about hybridization. We know about valence bond theory. We have to take away a couple of the d orbitals to mix with the s orbital and the p orbitals in order to get six hybrid orbitals. So if that sounds a little unfamiliar to you, you may want to review hybridization theory. The same exact idea, take enough valence orbitals, atomic orbitals to make hybridized orbitals. And if this is an octahedral complex, we need six of them. That means two of the d orbitals aren't there anymore, because we've made hybrids of them. That leaves us three other d orbitals. And maybe the electrons are in those orbitals, and that should be diamagnetic. So what the heck is it that makes cobalt decide whether it wants to do this, or do this, because apparently both are happening, but in different compounds? And how do we try to get to a point where we can predict something about this?
Well the color we know must come from the fact that light is absorbed. And that's going to be our big clue. If light is absorbed, we are going to see the complimentary color. And I'll remind you about that. If red light is absorbed, our brain says, "Hey, you know, red light must be green." Likewise if yellow light is absorbed, we perceive that as blue and so on. So we see a color that we can measure what light is actually absorbed by the material.
We can convert that to the energy of a photon. And that tells us about the difference between two different energy levels, an electron in a low energy level jumping to an empty place, a higher energy level, a higher state. And we know what that energy difference would be, because we know the energy of the photon. So okay, there is our big clue. We'll look at the absorption spectra of these two things. And what it will tell us is that this guy, the hexafluora cobalt has got a smaller separation in energy levels than this guy does.
Well okay, so what does that mean and how come are they different colors? Well we know why they're different colors. Or at least we could explain that by this difference in energy. But again, why the different magnetic properties, and what causes this separation to be different? Just to clarify here, in case I didn't make this clear, this separation is going to be the lowest energy separation in orbitals we can find in order to get absorption in the visible region. And that turns out to be the d orbitals. So this is corresponding to an electron going from a lower energy d orbital to a higher energy d orbital. And that's one of the reasons, by the way, one of the big reasons, transition metals have all of this color whereas main group doesn't, is that these d orbitals apparently are very close together in energy. So instead of absorptions in the ultraviolet, where we can't see them so materials look white, these transition metals are absorbing in the visible region, because these energy levels become very close together. Again these are both d orbitals. So they are fairly close in energy. And that's again what's going to give us this color.
But why does it change, and where does that separation actually come from? Because we had this kind of primitive picture of all of the d orbitals at the same level, so something is going on. Something is causing the separation in the d orbitals to change as we change the ligands. What's happening? Stay tuned and we'll find out.
Transition Elements
Bonding in Coordination Compounds
Color and Transition Metals Page [2 of 2]