Electron Configurations and Periodicity

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Electron Configurations and Periodicity Group Trends Transition Metals and Nonmetals [Page 1 of 2] We’ve been exploring the periodic table and finding the relationship of elements within the same family and the similar types of chemical reactivity that they have. Let’s now turn to the transition metals. Transition metals, again, are metals like the alkali metals and alkaline earths, but they have very different properties. In particular, physically, in their elemental forms—meaning the elements by themselves—we find that transition metals have much, much higher melting points and are much harder, in general, than the alkaline metals or the alkaline earths.

Let’s look at the electronic structure of these guys. The transition metals, you’ll recall, are characterized by—in fact, why don’t we back up a second and define what we mean by “transition metals.” A transition metal, strictly speaking, is any element where we’re in the process of filling the d orbitals. By that definition, elements in this family—zinc, cadmium, and mercury—are, in fact, not considered transition metals, in that technically, we’ve just completed filling the d shell for these elements. So although they’re very often grouped in this part of the periodic table as part of the transition metals, strictly speaking, in many ways, they’re not. Anyway, back to these transition metals that are in the middle of filling the d electrons. You’ll recall that in transition metals, we first filled the s electrons prior to the d electrons. Like this would be the s electrons of the next shell. So, for instance, in scandium we had . A characteristic feature, then, of transition metals is that they have two electrons in, let’s say, the n level, and in one lower than that, the n – 1 level, they have some variable number of d electrons—like again, . This is just a generic description that chemists often use to describe electron configurations for transition metals.

Now, my point is this: there are two electrons in an s orbital. In fact, when we start to remove electrons, as we start to form cations for these things, what we find is that the electrons that leave first for the transition metals actually are the s electrons, not the d. Now, that might puzzle you a bit, because as we went through the Aufbau process, remember that we put electrons in the 4s orbital first, because it was lower in energy than the 3d. But now what I’m telling you is the reverse, that when we start to take electrons out, it turns out that the electrons come out of the 4s orbital first. What’s going on here is very subtle, but as we remove electrons, we’re creating positive ions. And those positive ions, having less electron density, cause a shielding. We have less shielding, and as a result of less shielding, the 3d and the 4s orbitals switch back to the order that they were in in hydrogen, for instance. Now the 4s is actually above the energy of the 3d. In other words, when we take out those first two electrons from transition metals, they’ll come out of that 4s orbital. In fact, by far the most common charge of the cation for transition metals is going to be the 2+ form. So again, most commonly, these guys are going to be in a 2+ state, but occasionally, they’ll also be in a 3+ state. In fact, one of the properties that distinguishes transition metals from, for instance, alkaline earths, is that transition metals often have a variety of different charged cations that they can have.

Let me drive on. While we’re talking about transition metals, let’s talk about another class of elements that collectively are referred to as the rare earths. In older nomenclature, these were called lanthanides and actinides. These all have a common feature: they contain partially-filled f orbitals. So whereas this was what we sometimes call the d block where we’re partially filling d orbitals, in the f block, we’re partially filling f orbitals.

Now, again, like the transition metals, we had kind of a strange electron configuration in most of these elements, and that is that they had two electrons in an s orbital in the valence shell, and one electron, at least very often, in a d orbital, and a variable number of electrons in the f orbital. So when these guys form cations—and by the way, let me point out that all of these guys are relatively electron-rich in that they have a tendency to easily give up electrons— very different than these nonmetals that we’ll talk about. When these guys give us electrons to form cations, very often we’re going to lose two electrons from the s orbital and the one electron in the d orbital. So a very common charge for the f block elements is 3+: one electron coming from the d orbital and two from the s, because of that strange electron configuration. And just like the transition metals, when the f block elements lose their electrons, that 4f orbital actually does drop lower in energy than the 5d or the 6s—and I’m talking about lanthanides in particular—but as a result, again, the electrons stay behind in that f orbital. So although it’s not always true, 90% of the time, the charge of the f block ions, if they’re going to form cations, is going to be 3+.

Now let’s switch our attention to this piece of the periodic table: the nonmetals. In the nonmetals’ case, let’s first point out the fact that as we go from the noble gases—remember what we said about the noble gases, that they’re unreactive to halogens, to chalcogens, and so on—that one predictable change is how they react with hydrogen. So whereas hydrogen fluoride is a combination of one part hydrogen to one part fluorine, we know that oxygen reacts with hydrogen in a 1-to-2 ratio to form water, whereas nitrogen reacts with hydrogen in a 1-to-3 ratio to form ammonia. Copyright © Thinkwell Corp. All rights reserved. www.thinkwell.com

Electron Configurations and Periodicity Group Trends Transition Metals and Nonmetals [Page 2 of 2] And carbon, moving still one further, reacts with hydrogen in a 1-to-4 ratio to form methane. As you might expect, within a family, we’re going to find similar types of reactivity. For instance, just as carbon can form a compound with hydrogen in a 1-to-4 ratio, so silicon can to make silane, . Just as ammonia can form and is a stable compound, so is phosphine, , or arsine, . And of course, just like water, there is , hydrogen disulfide, very similar in a lot of aspects to water, except it smells a heck of a lot worse. But we’ll get back to that.

Turning now to the far right of the periodic table, there are the halogens. Let’s take a moment and look at the halogens. We pretty much can start to predict what’s going to happen now. The halogens all share the common property that they’re one electron shy of completing their shell, so they all have this desire to accept an additional electron to, again, reach the noble gas configuration. We tend to see these guys in a 1-to-1 ratio with alkali metals such as sodium. Remember, those are one-electron givers, if you will, the halogens being one-electron takers. We’ll see a 1-to-2 ratio with the alkaline earths; we’ve got now a 2-electron giver and two 1-electron takers, if you will.

We might ask, “Of all the halogens, which is going to be the most reactive in this type of a reaction?” If you think back, we were talking about the alkali metals, and where we found the most reactive alkali metals was down here in this part of the periodic table. Because of the way the alkali metals reacted—in that they wanted to give up an electron, or it was rather easy to take an electron from them—the ones that were most easily taken advantage of would be at the bottom of the periodic table where the electrons were held the least tightly. When we look at the reactivity of the halogens, remember, they want to react in such a way that they take electrons. Now we want to look at the top of the periodic table, because that’s where—this area, at least is where—electron affinity, the energy released as we get an electron—is the highest, in this portion of the periodic table. We also remember we’ve used the term “electronegative” to describe this portion. This is the highest electronegativity of elements. So we would expect fluorine to be more reactive than chlorine, and more reactive than bromine or iodine, and in fact, that is the case. Many a lab has been blown to bits, in fact, by fluorine, because of its incredible reactivity. It asks no questions when it takes electrons.

Finally, let’s turn to the very end of the periodic table: the noble gases. Just to remind you, one of the most common places we see noble gases is in discharge tubes. We know these things as neon signs. This gas happens to be helium, but neon, as we know, gives this more distinctive orange glow to it. So noble gases are great as far as passing electricity and getting light out of them, but they’re pretty lousy as far as reactions go. They have stable electron configurations. They’re not interested in taking electrons, and they’re not interested in giving electrons. For the most part, these guys show no reactivity at all—hence their name, “noble gases.” They’re found only in their elemental form with very few exceptions. Let’s close by thinking about what that exception is likely to be. Let’s think for a moment about these noble gases and the different ways they could react. They could accept another electron. Well, in all of these guys, to accept another electron, they’re going to have to go to a higher energy level, and the effective nuclear charge is going to be zero, because we have equal numbers of protons and electrons screening them and the next available space is much further away, so the screening is very complete. There’s no way we’re going to get these guys interested in taking another electron.

The other thing that we could do is ask them to give up an electron. If we want to combine them with a partner that’s going to take electrons and show no mercy, the best bet is going to be fluorine. Fluorine, again, is going to have the highest electronegativity overall, so if we’re predicting a reaction with fluorine, which of the noble gases would you expect is going to be the most likely to react with fluorine? Well, the answer would be xenon. (And I’m not counting radon here. Radon would be even more reactive, but it’s pretty radioactive, so we normally don’t do much with it.) Xenon, of the rest of these noble gases, is going to be the most reactive because of how it’s going to react. It’s going to react by giving up electrons. Because it’s being asked to give up electrons, those electrons are the furthest away from the nucleus in xenon, compared to these others. They’re the highest-energy electrons, in other words. Those are going to be, again, the easiest to get away. So indeed, the only reactions that we know of with noble gases are reactions with xenon and krypton, and most of these involve fluorine.

Overall, we’re starting to get a glimpse that we can actually predict chemical reactivity by knowing about electronic structure. It becomes more clear to us why we’ve spent so much time—why you’ve spent so much time—learning about all of these atoms and their electron configurations, and much more than you thought you’d ever want to know, because by knowing where the electrons are, you’re going to be able to predict how they react. And that is the power of chemistry: being able to predict reactions and what the outcomes are going to be.