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Chemistry: Electron Configurations beyond Neon


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  • Type: Video Tutorial
  • Length: 9:39
  • Media: Video/mp4
  • Use: Watch Online & Download
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  • Download: MP4 (iPod compatible)
  • Size: 103 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Electron Configurations and Periodicity (11 lessons, $17.82)
Chemistry: Electron Spin & Pauli Exclusion (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 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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Now we're in high gear as far as understanding the electron configurations of atoms. We've been using the Pauli Exclusion Principle, which says that no two electrons can have the same four quantum numbers; Hund's Rule, which says that an electron, given a choice of being in a separate orbital or pairing with an electron in a different orbital of equal energy, will choose to be in its own orbital; and the basic notion that an electron wants to be in the lowest energy state possible. Now, the process that we've been describing is actually referred to as the Aufbau Process or Aufbau Principle, describing the building-up--which literally that means--building-up of the electron configurations for atoms.
So let's continue on a little bit here and do our third series. We'll start with sodium. Sodium has an electron core of neon but one additional electron now. So that electron, we would guess, is going to go into the 3s orbital. For magnesium we have one space left in the 3s orbital so we can put it in as long as we change the spin state, if you remember. So we would describe its electron configuration as a neon core plus two electrons in the 3s orbital, or in other words, 3s2. For aluminum we have the neon core but also two electrons in the 3s and one electron in the 3p. For silicon we have neon core, 3s2, 3p2. Going to phosphorus, same story. We have 3s2 and now we have three electrons in the 3p. For sulfur we have a 3s core, four electrons now in the 3p, and notice we're starting to pair. Once again, notice that given the choice--if we go back to silicon for a moment--electrons chose to be in separate orbitals rather than pair up. That's Hund's Rule again.
Chlorine, we end up with seven electrons in our valence shell now. And finally argon, all eight electrons--all spaces that are available in the n=3 level, at least for the s and the p orbitals, are filled. Now notice again a very strong similarity between argon and neon, in that in argon the valence shell has its s and p orbitals filled, and so it we might expect to behave chemically similar to neon in that it's going to be difficult to remove an electron because we've reached a maximum and effective nuclear charge as the charge of the nucleus continues to climb as we go across, but we've also reached a point where there are no more vacancies for an electron at that level unless we go to the 3d level.
Well, what about that 3d level? Let's take one step beyond argon and we'll go to potassium. Potassium, we would expect by our logic that that next electron would want to go into the 3d orbital because, after all, that should be the next available place. But in fact experimentally we find that that's not true, that the last electron in potassium is in fact in a 4s orbital rather than a 3d orbital. Now that doesn't make sense to us, so let's see what's going on here.
In order to explain that, we have to go back to this notion of shielding. And remember, shielding is what was responsible for making the 2s energy below the energy of the 2p when we were talking about our first introduction to shielding. Likewise, when we're at the n=3 level, shielding is going to differentiate between the 3s orbital, the 3p orbital, and the 3d orbital. We have more shielding in the 3d than we do the 3p or the 3s. So just like we saw at the n=2 level, although they're all the same energy for hydrogen, they start to split with the 3s below the energy of 3p, and that below 3d.
Now look what happens. The n=4 level also has that same differentiation between s, p, d, and now f orbitals, and the 4s orbital actually drops below the energy of the 3d right at the point where we start to put in electrons. Right at potassium, as it turns out. So potassium unexpectedly has an electron in the 4s orbital rather than the 3d. And in fact we can more or less rely on that order, that the first two electrons as we continue to go on will go into the 4s orbital, then go into the 3d orbital. And what we're going to start to see is that these two orbitals in fact are very close to each other, and are so close that sometimes predicting where the electrons are can be a little tricky. So we'll try to generalize what's going on through this region next.
Now, just kind of recapping what I was just saying, for potassium then we have now our n=3 core, at least the s and the p orbitals completely filled. So let's put in our argon core, but now we want to put in one more electron. And what I was just saying is that the 4s energy has dropped below the energy of 3d. Simply put, putting an electron in a 4s orbital is a lower energy than putting it in a 3d orbital, once again because the shielding is not as complete for the 4s orbital. It can penetrate deeper towards the nucleus. So an electron here feels a little bit higher effective nuclear charge. So we're putting an electron in for potassium.
For calcium, there is still one space left in the 4s orbital, and that also will get an electron. Then not until scandium do we finally have a situation where there is no more room in that 4s orbital and we're going to now start to put electrons for the first time into the d orbitals, and one electron now will go into the 3d orbital.
So at this stage let's turn to a Periodic Table and start to make sense of this. Here's where we just were. We filled up this series. We filled in 3s and then we filled up the 3p and we got to argon. Then we said, okay, that next electron ends up going into the 4s orbital rather than the 3d, as is the case for calcium. And in fact we're going to start to know here that chemically potassium looks a lot like sodium, lithium, and rubidium. We'll say more on this later. But by having an electron in the s orbital rather than the d orbital, potassium will start to take on the chemical properties of these other elements which had one electron in an s orbital. Likewise, calcium is going to mimic some of the behavior of magnesium for the same reason.
Now here's where we started to fill the D orbitals. And with the help of our graphic here, you'll note scandium is going to be 4s2, 3d1. Titanium is 4s2, 3d2. Vanadium is 4s2, 3d3. Now, chromium is a little bit odd. Chromium elects to take one of its electrons out of the 4s orbital and put it also in a d orbital, and by doing so it gives a half-filled shell. Chromium's electron configuration would be, again, an argon core, 4s1, 3d5. So you'll note that every one of the d orbitals has one electron in it, and there seems to be an extra little bit of stability having to do with half-filling that set of d orbitals.
Manganese behaves back to normal. And not until we get to copper do we find another one of these odd exceptions. Now what's really going on here again is that the energy of the 4s and the 3d orbitals is almost on top of each other. They're almost identical, and so small differences can determine whether an electron goes to a d orbital or an s orbital. And in copper, once again, it elects to remove an electron from an s orbital so that it may completely fill the d shell, in this case. Zinc is back to normal. We've got 4s2, 3d10. Now we have no more places to put an electron.
And let's turn to this graphic now. We're at the point where we have filled our 3d orbitals and the next logical place to put an electron would be the 4p orbitals. So now what would you expect to be true about an electron that has one electron in a 4p orbital? You'd expect it to behave the same as elements that have one electron in a 3p or a 2p orbital. And indeed that's the case with gallium. It behaves very much like aluminum or boron in that they all have in common one electron in their valence p orbital. We could continue to fill up this row then, taking us all the way to krypton, and everything else behaves as we've seen previously.
So we've seen another case where shielding has dictated where the electrons go. Our next step then is to look at the Periodic Table in whole and start to see the relationship within families.
Electron Configurations and Periodicity
Electron Spin and Pauli Exclusion Principle
Electron Configurations Beyond Neon Page [1 of 2]

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