The Spin Of Electrons -- Key To Understanding "Paired Electrons"
 

Spin? What's that?
 

Well, it's an additional property that electrons (and other particles) possess. Here's an analogy: think about the Earth orbiting the sun--
 

Haven't you just been pounding into my head that electrons don't orbit like planets?
 

It's true, they don't--and yet that picture remains helpful and illuminating in many contexts. So bear with me for a moment: think about the Earth. Not only does it orbit around the sun once a year, it's also spinning once a day on its own axis...
 

And that's what spin is! Although I suppose you're going to tell me that electrons don't really spin, any more than they really orbit.
 

You catch on quickly. They don't spin--but it's tremendously useful to think about them as if they did, and for most practical purposes, you can. In the present case, think of the two electrons in that lowest energy level as spinning in opposite directions. It's often said that one has "spin up" and the other "spin down."
 

So each level has room for a spin up and a spin down--that makes sense. But you haven't explained s, p, and d yet; there must be more complications to worry about.
 

Indeed...
 

Take a look at the elements beyond lithium and see what you can discover.
 

Let's see...the fourth element, beryllium (Be), has a second electron in the higher energy level, which means that level now has both a spin up and a spin down. So the next element should begin a third energy level--is that right?
 

See for yourself. Try clicking on boron (B), the fifth element--you know, the one Bruce Willis was so excited about.
 

Hey, what's going on? The fifth electron has a slightly higher energy than the other yellow ones, but it's not directly above them; it's in that column labeled "p."
 

Ah. Do you have any ideas about what might be happening?
 

Well...my best guess is that the colored rows, pink and yellow, represent the main energy levels, and s, p, and d are like smaller sublevels of them.
 

Very good. So what do you think is going to happen as you keep going along that row of the periodic table?
 

Hmm...carbon (C) has a second electron in the p column, so now s and p in the yellow row each have a spin up and a spin down. The next electron must start a whole new energy level, or maybe it goes into the d column, if that's the next higher sublevel.
 

Sounds very logical...but now look at nitrogen (N).
 

Hey, the seventh electron went into the p column too! How can there be three with the same energy?
 

It gets worse. Go on.
 

Oxygen (O), fluorine (F), neon (Ne)--more electrons just keep getting stuffed into that same state. What happened to the exclusion principle? This makes no sense at all!
 

It makes perfect sense, once you know the rules.
 

I wish you'd stop being so mysterious and explain these "rules" you keep referring to.
 

I was just about to do that. I have to warn you, though, that the rules may seem arbitrary to you, and I won't be giving any satisfactory explanation of the reasons for them. Partly that's because I want to spare you a lot of complicated math, and partly it's because this is just the way nature is. I'll just ask you to have faith that all this numerology comes out of a sophisticated mathematical theory, and has been upheld time after time by experiment.
 

I can, if you like, tell you about quantum numbers; they provide a more quantitative way of understanding these rules.

For now I think I'll be satisfied if you can tell me how to predict those electron arrangements you've been showing me.
 

I can do that. First of all, you were correct when you guessed that those colored rows in the chart correspond to the "main energy levels"; they're often called primary energy levels, incidentally. Usually, a higher row means a higher energy, and energy gaps between rows tend to be quite large, in comparison with the gap between, say, s and p.
 

Are you ever going to explain what s and p mean?
 

I'll do that right now. As you surmised, the s, p, and d columns represent smaller "sublevels" of the primary rows...
 

Then why not just call them A, B, and C, or something else at least vaguely logical?
 

This is a bit of archaic notation left over from nineteenth century spectroscopy--rather silly, but everyone uses it, so we're stuck with it. If you must know, s stands for "sharp," p is for "principal," and d is for "diffuse"--supposedly they refer to the appearance of various spectral lines. The next one is called f, for "fundamental"; mercifully, the subsequent ones just go alphabetically: g, h, etc.

Okay, so why is it that s only holds two electrons, but p has room for six?
 

Well...remember that each electron is existing in one of those strange probability clouds, which, as you've seen, can have widely varying shapes and sizes. Another statement of the Pauli exclusion principle I mentioned is this: no two electrons in an atom can be in the same type of cloud with the same spin.
 

So you're saying that p electrons have more cloud shapes available to them?
 

Precisely. It happens that s clouds are always spherical; the spheres just get bigger as the primary energy level increases.

However, p and d states are more interesting: there can be several different-shaped clouds at the same energy. For example, here are two p states from the second primary level:
 

It turns out that there are three kinds of p clouds in each primary level.
 

...and each one can hold a spin up and a spin down, so that's why six electrons fit into the p column!
 

Very good. You can think of each electron's "quantum state," its full, unique description, as being the sum of a particular probability cloud plus a spin:
 

	+
A certain d state in the third primary level		spin down

I understand the rules so far, but I still have a lot of questions. First of all, I can see from the chart that the lowest row only has s electrons, but the next row has electrons in both the s level and the p level, or whatever you're supposed to call them...
 

"Sublevel" is fine; you can also speak of electrons' being in s or p "states." You'll frequently hear these states called orbitals, especially in chemistry. Of course, that term can be a little confusing--
 

Yeah, yeah, I know--because electrons don't really orbit.
 

Good, you've been paying attention. Incidentally, the s states in the first primary level are called 1s orbitals, those in the second row 2s, and so forth.

Now, to return to your question: another rule you'll have to remember is that the number of sublevels increases with each primary energy level. The first row has just s orbitals, the second has s and p, the third s, p, and d, and so forth.
 

The way the periodic table is arranged is starting to make sense to me now. The rows of the table match up with the primary energy levels--that's why the first row only has two elements. In the second row, lithium and beryllium are filling the two 2s spots; then there's a big space because the next six elements are filling the 2p orbitals.
 

That's exactly right. Now look at the third row; is it the way you'd expect it to be?
 

Hmm...sodium (Na) and magnesium (Mg) add the two 3s electrons, and then the next six elements, up to argon (Ar), fill the 3p orbitals. That's all fine, but why does the row end there? This row should have a d sublevel, too, if you were telling the truth.
 

I never lie--but I won't deny that there are more subtleties I haven't yet revealed to you. Click on potassium (K), the first element in the fourth row.
 

Hey! Now there's an electron in an 4s state, and still nothing in the 3d orbitals. How do you explain that?
 

I've never come right out and said this, but I'm sure you've noticed that the energy levels get filled in order from lowest energy to highest; when you add a new electron, it goes into the lowest-energy state that's available.
 

Sure, that's pretty obvious. That's why I thought the next electron for potassium would go into a d orbital instead of up to a whole new row--you said that a higher primary level always means a higher energy.
 

I said usually, not always. The energy levels aren't always so well-behaved as one might like; it sometimes happens that the first orbitals in a "higher" primary level actually have less energy than the top orbitals of the level below. Click on the advanced button to find out why this occurs.

So in potassium, the 4s orbitals end up with less energy than the 3d states--that's why potassium starts a new row in the periodic table. I bet calcium (Ca) puts an electron in that second 4s orbital...yup, it does.
 

Okay, now click on scandium (Sc).
 

Now the 3d orbitals are getting some attention! All those elements from scandium through zinc (Zn) are just filling the ten green d spaces. Then gallium (Ga) goes back up and starts the 4p orbital.
 

This periodic table ends at krypton, when the 4p orbitals are filled--but of course one can keep going. Click here to open an extended version of the periodic table, which shows the electron configurations for the elements up to 103. I'll let you play around with this and figure out the order in which the various orbitals are filled.

All right, I think I understand now why the periodic table is laid out the way it is. The rows go with the primary energy levels...
 

...and there are different sections of the table that go with the different sublevels.
 

That's right; you can think of the table as being divided into "blocks" like this:
 

That makes so much sense! Now I can see why someone had the idea of putting the elements in a table like this; once you know how the electron configurations work, this arrangement wouldn't be so hard to come up with.
 

Actually, the first designer of the periodic table knew nothing about energy levels or even electrons; at the time, hardly anything was known about the structure of atoms.
 

Really? Then how would anyone know which elements to put where?
 

Well, that's quite an interesting story. You see, there's more to the periodic table than just electron configurations. Up until now I've been giving you a physicist's view--from the bottom up, so to speak. Now we'll look at the periodic table from a chemist's perspective, from the top down.