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Chemistry: Electron Affinity


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About this Lesson

  • Type: Video Tutorial
  • Length: 12:46
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 137 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: Periodicity (4 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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Recent Reviews

Great! Thank you! The best!
~ afernandes

I've checked other lessons on the same subject in the Internet.

But that is the best (despite the HORRIBLE musics and sounds using as a helpers in the intermissions. Usefull, but not fun at all).

Brazil - Rio de Janeiro

Great! Thank you! The best!
~ afernandes

I've checked other lessons on the same subject in the Internet.

But that is the best (despite the HORRIBLE musics and sounds using as a helpers in the intermissions. Usefull, but not fun at all).

Brazil - Rio de Janeiro

We talked about the notion that it takes energy to remove an electron from an atom because that electron is attracted to the nucleus. What about an atom that had a space available, if you would, for an additional electron--an orbital that's low in energy, meaning physically that if we were to put an electron in that orbital, that electron could benefit from the charge of the nucleus? So again we could talk about a potential well. In this case, it would be a potential well that had an empty space available to accommodate an electron. And we could ask, "What happens in terms of energy if an electron is dropped into that well?" This electron is not part of the atom. This is an electron that we're introducing to the atom now, in this case.
Well, for that process--remember, we're going from the energy of that electron being zero to a negative value. That means that energy, in order to conserve energy, energy has to be released, and that's going to be released in the form of a photon. So the energy difference between an electron dropping from here to here is going to be made up in the energy of the photon that comes out as a result. By knowing the frequency of that photon, we can know the energy difference, and that term is referred to as "electron affinity." It describes the general appetite of an atom for an additional electron.
Now, this is a very easy term to get confused with the opposite of ionization energy. In one sense, it's the opposite in that we're giving an electron rather than taking away. But don't make the mistake of thinking that it is indeed the exact opposite of ionization energy. The exact opposite of ionization energy would be putting an electron into a positive ion to make the neutral atom. That would be the exact reverse. I'll get our graphic and show you that.
In electron affinity, we are offering an electron to a neutral atom, generating as a result an anion. So let's again define our terms here. We're going to describe electron affinity as the energy released when we give a neutral atom an additional electron. So in this case we make an anion and we get energy released, usually in the form, again, of a photon. So in cartooning them, we put the electron into an empty space in an atom and a photon is released. The energy of that photon is what we would describe as the electron affinity.
Now, again, let's try to make some predictions from what we know now about orbitals and electrons and electron configurations. As we go from hydrogen to lithium to sodium--in other words, as we go down in the periodic table--we know that in all of these cases there is indeed a vacancy in the valence orbitals. In other words, hydrogen is 1s1, where we have one space left in that s orbital. Lithium is 2s1, and you have one space left in that s orbital. Likewise with sodium, there's one space in the 3s orbital. What is the difference between putting an electron in there--in fact, let me back up a second and say we might predict that because there is that one empty spot, the atom indeed may want to take another electron. And if an electron did go in, again, energy would be released, that being the electron affinity.
How do those electron affinities change as we go down the periodic table? Well, the big difference is that if we were to put an electron in this orbital, it's going to have the benefit of being closer to the nucleus than if we put an electron in this orbital. Effective nuclear charges are going to be more or less the same as we're going down in the periodic table, and so the big difference is that electron being able to get a lot closer to the nucleus on average. So our prediction would be that electron affinity would be highest at the top of the periodic table and decrease steadily as we go down in the periodic table.
Also, we might expect that as we go across the periodic table, where effective nuclear charge is increasing as we go across the periodic table, that we'd expect, since the vacancy in the valence orbitals is going to be in an orbital where it feels a higher effective nuclear charge, we'd expect the electron affinity to increase as we go across the periodic table. We'd expect the most electron affinity from elements in the right part of the periodic table and the lowest electron affinity for elements in the left part of the periodic table.
Well, in general, this is in fact what we observe, but there are a number of exceptions to this. So let's take a look at the data and see if we can at least get comfortable with the notion that these generally are the observed trends, and where there are exceptions to this, let's see if we can start to rationalize why that would be.
Again, for convenience sake, I'm going to not consider the transition metals at this stage and just look at our main group elements. What you're seeing here is that, again if we look at the alkaline metals, the general trend is observed that if we go from hydrogen, which is not truly an alkaline metal but we'll include it here, and it has a 1s electron in its valence shell--as we go from hydrogen to lithium to sodium to potassium to rubidium, we notice that the electron affinity is indeed decreasing. Good, we feel good about that and we understand why that would happen. So far, so good.
As we go across the periodic table, indeed it is the case that in general, over here in the right part of the periodic table, the halogens and the chalcogens, these guys have the highest electron affinity. They're the most desiring of an additional electron, if you will. We get the most energy out if we give them an electron. But you'll notice that it is not a smooth trend. In fact, something goes terribly wrong not too far across the periodic table. By the time we get just into the second period, the alkaline earth, notice that the electron affinity drops to zero or is even below zero. Now, experimentally, that just simply means that we can't make these elements take the electron. That must mean that there is overall no energy benefit for these atoms to pick up an additional electron.
Why would that be? Let's take the simplest case of beryllium and again go to an electron configuration. In the case of beryllium, we have four electrons. So again, let's not lose sight of what electron affinity is. It's the energy that we get released by giving this neutral atom an additional electron. Well, if we gave it an additional electron, where is it going to put it? We know from the Pauli Exclusion Principle it can't be down here, so it's going to have to go into the 2p orbital. Well, we know the 2p orbital is higher energy than the 2s, and in fact, these four electrons are going to do a really good job of shielding the entire nuclear charge. Now, beryllium has got a 4+ nuclear charge. These four electrons are all inside this 2p orbital. They effectively completely shield an electron that we would be putting here. As a result, there's no net energy benefit to putting an electron in this orbital, and so beryllium says, "No, thank you. I'm not interested in the electron," and we have a net energy that's less than zero, meaning again that it's not a stable process to give beryllium the electron.
Moving into the third row and the fourth row, now we start to see the general trend picking up again, but wham, we run into trouble again when we get to Nitrogen. Okay, once again, let's go to nitrogen's electron configuration and make sense out of this. One, two, three, four, five, six, seven--that's nitrogen. We do have spaces in the 2p orbital, room for rent in other words. We can bring in another electron here. But look at the problem. If we bring in another electron, we're forcing it to pair. We saw this before with ionization energy. Electrons don't want to be in the same general volume as other electrons if they can avoid it. Apparently, that price to pay is enough that nitrogen again is not interested in taking another electron. It has a place to put it but it costs too much to try to make that electron pair with another electron. So its electron affinity is not something we can measure because it refuses to take the electron.
Oxygen, fluorine--the trend again picks up. You'll notice, in fact, with the halogens, this is where we're going to find the highest electron affinities. The general personality, if you will, of these elements is to take additional electron density. We've already discussed why that is: if they get one more electron, they complete their valence shell. This is the maximum effective nuclear charge for a given period over here. And now, apparently, in the case of fluorine, even if we've got to pair the electron, the effective nuclear charge is so high now that it says, "Fine, give me that electron. We'll make them pair." And so these guys have high electron affinities, and again, also in this family. So this part of the periodic table is where we're going to find the elements that want to have extra electrons the most.
Now, one last thing to point out. You'll notice that there's a little bit of a problem as we go down the periodic table, on this side of the periodic table. In general, it is true that as we drop down the periodic table, electron affinity is going to decrease. Remember our explanation for that was that we're putting electrons in orbitals that are further and further away from the nucleus, so it makes sense that we don't get quite as much energy back if we're putting electrons in a place that's not quite as good as we go to further and further places from the nucleus. But notice that there's a reversal of that trend at the very top of the periodic table. Fluorine actually does not have as high of an electron affinity as chlorine. We get more energy out of chlorine by giving it an electron than we do fluorine. And in fact, you'll see the same thing is true with oxygen and sulfur as well, in fact.
So what's going on that's different here? Well, the general explanation of why there's this abnormality to our general trend is that by the time we've gotten to fluorine or oxygen, the effective nuclear charge is so high that those 2s and 2p orbitals are pulled way inside--very, very tight compared to how they were over in this part of the periodic table. These atoms, remember, are much smaller than over here because of this higher charge. And the generally accepted explanation is that although there's a space here and fluorine is indeed willing to take that extra electron, that it indeed is in a volume that's relatively small and where there's a lot of congestion, a lot of other electrons. And that it's the repulsion of these electrons that starts to take away some of the value in bringing in that extra electron. So it still is a good deal. Fluorine is happy to take that extra electron, but it's not quite as good as in chlorine, where there's a little bit more room for the electron density to spread out.
Now, as you go below chlorine, the trend starts to behave itself again, and electron affinity drops as we go down in the periodic table.
Electron Configurations and Periodicity
Electron Affinity Page [2 of 3]

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