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Chemistry: Electron Shielding


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About this Lesson

  • Type: Video Tutorial
  • Length: 9:05
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 97 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.

About this Author

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In trying to describe Helium, the strategy that we took was to put one of each of the two electrons of Helium in the 1s orbital, and this didn't violate the poly-exclusion principle because we could take advantage of the fact that those two electrons could have different spins. Remember, the poly-exclusion principle says in English that two electrons can't be at the same place at the same time, essentially, meaning we need a unique set of quantum numbers to describe each of those two electrons. So at the end of the day we were able to get both of the electrons in the Helium atom into the same orbital.
Now, what happens with Lithium when we now go to three electrons in the system instead of two? We have no more choices at the 1s level because there is no other way to describe an electron at that lowest energy state, so we're forced to go to the n=2 level; in other words, to put an electron in either the 2s or the 2p orbital. Which should we use? They're the same energy. At least we found they were the same energy in the Hydrogen atom. Well, it's going to turn out that they're not the same energy in the Lithium atom; that the 2s, we'll find, is a little lower in energy than the 2p. And our goal for this tutorial is to understand why that would be.
Okay, before we talk about that, I want to give you an analogy that will help us picture what's going on. Suppose now that you're a moth. Now, I know I've asked you to think of some strange things, but go with me here. You are a moth flying at night. It's dark. Suddenly you feel the intense traction of a light bulb. Someone has turned on their porch light and you feel sucked in towards this light bulb. As much as you fight it, trying to fly away from it, you're pulled in, into an orbit around the light bulb. Okay? Now, suppose that you're not alone. You're joined by other moths that get sucked into this light as well, drawn into this source. As more and more moths come in, at first they don't really make any difference as far as how much light you see. You might be feeling the light but another moth may be over here at the same distance, and it doesn't really interfere with your ability to see that light source. But as more and more moths come in, and maybe you lose your initial front-row seat and you kind of get pushed out a bit--there's a whole bunch of moths surrounding the light and you're out here--you're not going to see as much of the light as if those other moths weren't there. We use the term "shielding" to describe how electrons around a nucleus, if they're inside other electrons, partially block or partially shield the nuclear charge, much in the way that these moths that I'm describing are partially blocking the light from a moth out here. So hold onto that idea and let's then remind ourselves a little bit about energy, and then we'll put this all together.
Remember that we said that if a ball is on top of a hill, we refer to that as having a positive potential energy. If we decide to define zero energy at ground level, then how do we describe a ball that's fallen into a hole that's below the surface of the ground? We have to use the term "negative energy" to describe the fact that it is lower energy than zero energy. So remember, in my discussion here, as I'm talking about more negative energy or lower energy, I'm talking about more negative again, lower than zero. And one of the consequences of having our nuclear charge increase as we go from Hydrogen to Helium to Lithium is that because there's a more intense attraction, this well, if you will, becomes deeper and deeper. There becomes a stronger attraction. Electrons drop to lower and lower energies, meaning below zero, more and more negative below zero.
Okay, now, hold onto that idea for a moment. Back to this notion of the 2s orbital versus the 2p orbital. Once again, the question we're trying to answer is, "Between these two orbitals, which is a better deal for this third electron, the last electron in Lithium?"
Okay, now, let me remind you of something else that we've talked about, and that's this notion that we describe what referred to as a "radial distribution function," which answers for us, "At what radius are you most likely to find the electron?" In other words, at what distance from the nucleus are you most likely to find the electron? So this is the radial distribution function for an s orbital and this is the radial distribution function for a p orbital. And we'll notice that even though in the p orbital, essentially the most probable place to find the electron is just about the same as for the 2s orbital, the 2s orbital differs in a very important way. It has a small probability of being very, very close to the nucleus, whereas the 2p orbital doesn't.
Now, if there's no other electrons in the system, putting an electron here or putting an electron here--both electrons will have the same exact energy. On average, they see the nucleus about the same amount. But when we put electrons in the 1s orbital--now, remember, Lithium has got two electrons in the 1s orbital and then one electron that we're going to put in the n=2 level, either the 2p or the 2s, and we're trying to decide which one. Okay, so what I'm going to do is take a crayon just to remind us about shielding here, and say, all right, this is, let's say, the nucleus, and let's say this represents the electron density for those two electrons in the 1s orbital. They are in here someplace.
Okay, now, the 2s orbital and the 2p orbital spend most of their time on the outside of that 1s orbital. So again, it's being shielded. These electrons--an electron in either of these orbitals--are going to be shielded by the electrons in the 1s orbital. So they're not going to see the same nuclear charge that these electrons do. These electrons in the 1s orbital would see essentially a +3 charge, but electrons out here are going to see much lower effective charge. We use that term, by the way, "effective nuclear charge," to describe the charge after taking into the account the fact that there are other electrons partially blocking the nuclear charge. So once again, the effective nuclear charge out here is lower than the nuclear charge would be in close to the nucleus. But again, that screening is not going to be the same for the 2s and the 2p orbital because the 2s orbital, although it spends most of its time out here, it has a finite probability of spending time inside the 1s orbital. And that means for a small amount of its time, it's actually inside the screening electrons. Inside the screening electrons, for part of the time it feels a higher effective nuclear charge. The higher the charge, the lower the energy, meaning more negative the energy. And so the consequence, the bottom line, is that the 2s orbital drops below the energy of the 2p orbital because it can get inside the shielding 1s electrons.
So our picture then becomes this. Here is the energy of the Hydrogen atom, with one electron in there. And again, you'll note the 2s and the 2p orbitals are the same energy. If we put this electron up here either here or here, it's going to be the exact same energy. But as we go from Hydrogen to Helium to Lithium, notice what happens. The 2s and 2p orbitals split in energy, with the 2s orbital dropping in energy below the level of 2p. Don't forget, zero energy is up here. So I'm talking about the 2s orbital being more negative than the 2p orbitals are. And down here someplace, again, are the 1s orbitals.
Once again, the crucial point here is that because the 1s orbital is filled, those electrons block this electron from seeing the nucleus, and it would block this electron if it were up here from seeing the nucleus. But the difference between an s orbital and a p orbital is that the s orbital spends a little bit of its time inside the 1s orbitals and so it drops to a little lower energy. At the end of the day, the 2s orbital is below the energy of the 2p orbital, and so that is where we're going to put the third electron.
Electron Configurations and Periodicity
Electron Spin and Pauli Exclusion Principle
Electron Shielding Page [1 of 2]

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