Chemistry: Periodic Relationships

This lesson has not been reviewed.

Please purchase the lesson to review.

We have arrived. We're now ready to make sense of the periodic table based on what we know about electron configurations. Let's start out at the far left, looking at a series of elements that we call alkaline metals.
What's in common with these alkaline metals? Well, if we look at lithium, sodium, potassium, rubidium, and cesium, we see that they all have one thing in common, and that is one electron in an s orbital in their valence shell. And so we might start to predict that these elements will behave chemically in a very similar fashion. In fact, we're going to see a demonstration of that very shortly. But again, not only would we predict these guys may be a little more willing than other elements to lose that electron, we would predict that they would form compounds with other elements in a similar ratio because they all have the one electron to give.
If we go from the alkaline metals to the alkaline earths, again you're notice a very important common feature of these metals, and that is that they all have two electrons in a valence s orbital. Let's go back to the periodic table for a moment and look at the alkaline earths, and you'll see beryllium, magnesium, calcium--again, all of these materials have a common feature that they have two electrons in their valence orbital, and in particular two electrons in an s orbital of their valence shell.
And so we see that the periodic table is in fact a very handy organization of elements based on their electron configurations. Once again, all of these elements and these elements have their last electrons--meaning their highest energy electrons--in s orbitals. All of these elements turn out to have their last electrons in the p orbitals. In fact, now what I'm going to do is just switch a little bit the periodic table we're looking at to better help us see these relationships. So same periodic table as before but now instead of the elements, what I've highlighted are what the last electrons--what are the highest energy electrons, the valence electrons--in these elements. And again, what's in common with this first block is that the valence shell is the s orbital--whether it's the 2s, 3s, 4s, 5s, so on. Over here, it's the p orbitals that are being filled.
Down here in the transition metals, we talked about that when we went from potassium, calcium, to scandium, that we started to fill the d orbitals. And so what is characteristic of all of the elements in this part of the periodic table is that the last orbitals to fill are the d orbitals, whether they're the 3d orbitals or 4d orbitals or 5d orbitals, and that again gives them common relationships with each other as we go up and down in the periodic table.
What's going on here we haven't described yet, but just as we talked about shielding affecting the d orbitals, p orbitals and s orbitals to differing degrees, shielding also affects the f orbitals differently. And in fact what we find is that the four f orbitals are so shielded that they don't fill until after the 6s orbitals do. Now, that's a huge difference but you can get a sense now of just how much the shielding is playing a role. In fact, the 6s, the 5d and the 4f orbitals are all essentially the same energy. So when we go to our periodic table, we find that for cesium and barium, the 6s orbitals are filling. Then we take a jump. We go from barium actually to lanthanum. This element--actually, people don't know quite where to put it because it has one electron now in the 5d orbitals, only one. And when we go to one more electron, we go into the 4f orbital. So you see again, these three orbitals are all just about the same energy. Very difficult to predict exactly where those electrons are going to be. But the periodic table helps tell us to some degree.
These are the 6s orbitals being filled. Then we jump over and fill the 4f orbitals. And some people--just be aware. Sometimes people will take those first two elements and leave them up here. So just that you're aware of that, that there are actually two different versions of the periodic table. But then the f orbitals are filled. Then we go back and fill the d orbitals. Then finally the p orbitals.
Now, let me go back and say that more specifically: 6s, 5d^1, then 4f, then we fill the rest of the 5d, then finally we fill the 6p. So again, you can see how the blocks of the periodic table correspond to what the valence orbitals are that are being filled. So this series here in green, the upper series, is called the lanthanides. And the same thing is going to happen to us one more time when we get into the actinides. The actinides don't happen until after we fill the 7s orbitals. Then we go into the 6d and 5f, and that's where this happens. This is now filling of the 5f orbitals. Then we go back into the 6d orbitals, and finally we go into the 7p, except that there are no elements here. At least they haven't been discovered yet. We get to a certain point, as we'll see later, where the nucleus just becomes too heavy, if you will. It becomes too many protons and neutrons and it's no longer stable beyond this point. So if these elements were ever to be discovered--and these are going to have to be man-made elements--then those would be the 7p orbitals that we would be filling.
All right, let's return once again to the notion that elements within a family of the periodic table share a common relationship; that their electronic properties will dictate that they will have chemical properties that will be similar as well. We talked about the alkaline metals all having one electron beyond the filled shell. And it's going to be a natural tendency of these metals to lose that one electron, forming monocadions and salts. In other words, as we use this term soon, the oxidation state of these materials will tend to be +1.
Now, for the alkaline earth metals, these will be two electrons beyond a closed-shell configuration. They will tend to lose two electrons now. And so a common charge for these things would be a 2+. In other words, we say that their oxidation state will be most commonly 2+.
Let's move all the way to the other side of the periodic table and what we'll see is that the halides are all going to be one electron shy of having a closed shell, at least with regard to their s and p orbitals being filled. And so the halides, we might predict--and we would be right--would tend to pick up one more electron in order to have a closed shell. Now, let me make this a little more clear by writing out explicitly what I'm saying.
So let's start with Chloride, for instance. We can describe the core of chlorine as--not chloride; I'm sorry. Chlorine, just the atom chlorine--as a neon core, 3s2, 3p5, okay? So again, one electron shy of filling the core. So notice that if we made chloride, if we give it one more electron, the electron configuration that we would write would be a neon core, 3s2, 3p--now, we need to add one more electron to make the monoanion, and that's going to be 3p6. And we notice that that is the same exact electronic configuration as the noble gas, argon. Neon 3s2, 3p6. So again, our prediction would be that chlorine, missing one electron from a noble gas configuration, would have a tendency to accept one more electron, thus filling its valence shell and arriving at the electron configuration of a noble gas.
Likewise, if we go one step beyond there, think about potassium. Potassium will have--and I'm going to write it a little differently than I normally would. I'm going to write it as a neon core, 3s2, 3p6--just to remind you that those electrons are there; in the future I might write potassium as having an argon core, but I'm going to just write it out a little bit more explicitly for now--and then 4s1. Recall that the s orbital is filled prior to the 3d orbital. So we can understand why potassium would want to lose an electron, because if we removed that one electron and formed potassium+, we end up with the same electron configuration as argon again. So we're going to find, as we look at periodic relationships and start to explore how atoms react with each other to form compounds, that elements that are very close to the electron configuration of a noble gas that have one extra electron will have a tendency to lose that electron or electrons, whereas elements that are very close to having a noble gas configuration but have one too few will tend to pick up that extra electron or two, once again to try to arrive at this noble gas configuration.
Just as a comment here: There's nothing magical about a noble gas configuration other than it describes a situation where there are no other empty vacancies to put electrons in the valence shell, and where we reach the maximum in our effective nuclear charge for a given period of the periodic table. So it's very hard to remove an electron; it's very difficult to put an extra electron in. And so as a result, the chemical reactivity of other materials in the periodic table, other elements, will tend in general to be towards trying to reach stable electronic configuration of these noble gases. And that's just one more time, the alkaline metals will tend to lose one electron and become monocations; the alkaline earths will tend to lose two electrons. On the other side, the Halides will tend to gain an electron; the chalconides will tend to accept two electrons, all of which give us the noble gas configuration.
Electron Configurations and Periodicity
Electron Spin and Pauli Exclusion Principle
Periodic Relationships Page [1 of 2]