Shells, subshells, and orbitals | Atomic structure and properties | AP Chemistry | Khan Academy

We've learned in other videos that the atom is, in fact, made up of even smaller constituent particles, which is pretty amazing because atoms are already unimaginably small. Those particles are the protons, which have a positive charge; you have your neutrons, which have a neutral charge or no charge; and then you have your electrons, which have a negative charge.

Now, the big question physicists and chemists were facing over 100 years ago is: how are these things configured? They realized that the positive charge is concentrated at the center of the atom. In fact, most of the mass, which is made up of the protons and the neutrons, is concentrated at the center.

So, an early model for how an atom worked was: maybe you have your protons and neutrons in the center. Let's say we're talking about a helium atom. A helium atom has two protons in the nucleus, and a typical helium atom would have two neutrons as well. So, the nucleus might look something like that.

Early physicists and chemists said, "All right, well look, the protons have a positive charge; electrons have a negative charge, so they'll be attracted to each other. Opposite signs, opposite charges attract; the same charge repels each other." So, maybe the electron, which has a negative charge, orbits around the nucleus the way that a planet would orbit around its star. So maybe it orbits something like this.

Maybe one electron has an orbit that looks something like that, and then another electron, if you're talking about a neutral helium atom, would have two electrons and two protons. Well, maybe the other one orbits something like this. I'm trying to just draw an elliptic or a circular-looking orbit. This is the idea that the electrons are in orbits.

Now, it turns out that this is not exactly the case; electrons are not in these well-defined circular or elliptical orbits. In fact, at any given point in time, it's not necessarily exactly right there. It could be there, but there's some probability it's here, there's some probability it's there, there's some probability it's over there.

To describe where electrons are likely to be found, physicists and chemists introduced the idea of an orbital. The best way to think about orbitals is to think about a hydrogen atom. Actually, the math for orbitals has been best completed for the hydrogen atom, as it is the simplest atom.

In a hydrogen atom, especially the typical isotope of hydrogen found on Earth, the nucleus actually has no neutrons; you just have a single proton at the center. If you have a neutral hydrogen atom, that one electron, instead of being in orbit around that one proton like that, we can really just think about the probabilities of where it might be.

It could be here; it could be here at any given moment; it could be there; it could be off the screen at some moment. But it's more likely to be in certain regions of space around the nucleus than others. We can visualize where it's most likely to be by saying, "All right, it looks like 90 percent of the time it's in a sphere that looks something like that."

Once again, it could be here, it could be there, it could be out here; it could be anywhere. We're just saying where it happens to be 90 percent of the time; that's the visualization.

Now, an interesting question is: what if you were to give that electron a little bit more energy? Well, what does energy mean? If you think about planets, or a rocket, or a satellite orbiting around, if you were to give it a little bit more energy—if you were to give it a little push—it could have a larger orbit. It would look something like that.

But quantum mechanics isn't about things happening gradually. Sometimes people think quantum means small or something like that. No, it really means that you're talking about discrete packets. So, in quantum physics, in quantum chemistry, if you add a certain amount of energy to an electron, instead of having a 90 percent chance of being found in this first shell, this first energy level, it could then be found. It would then jump into the next energy level or the next shell.

Now it might be more 90 percent of the time it's going to be found in this shell right over here. Then, if you were to give it the right boost of energy, once again, just a little bit won't do; you have to give it enough so that it jumps into the next energy level. Then it might form these weird patterns that look kind of like dumbbells, where 90 percent of the time you can view it as it's on the orbital that looks kind of like that dumbbell shape.

I just did it in kind of the horizontal direction; you could have it in the vertical direction. You could also have it in the in-and-out direction of this page. If you're wondering where these shapes come from, if you keep adding more and more energy, you get these more and more exotic shapes for orbitals. Think about standing waves—that's my best hint I can give you.

At the quantum level, actually at all levels, but especially at the quantum level, you see things like electrons have both particle and wave-like properties. So imagine something like a standing wave, where if I were to just take a rope and shake it, I might get standing waves that look like that.

If I were to take some type of a membrane in two dimensions and push on one side right here, if I were to drum on that, you might get it so this part dips down and then that part dips up. So when you get to three dimensions, you end up getting this dumbbell shape when you add more energy, and then you get more and more exotic shapes.

Just to imagine what some of the first orbitals look like, rendered by a computer, you see it right over here. If you have your lowest energy electron, you are in what is called an s orbital right over here. We would call this 1s because it is at the first shell, the one closest to the nucleus.

If you give even more energy, then that electron might jump into the second energy level or the second shell. The orbital in that second shell, which would be the default if it's the lowest energy in the second shell, would be the 2s orbital. Once again, you have this spherical orbital; it's just more likely to be found further out than when it was just in the one shell.

Once again, if you add even more energy, you'll still be in the second shell, but you will be in one of these orbitals that have higher energies. You could view this as the 2p orbital that is in the x dimension. This could be the 2p orbital that is in the y dimension.

So some people call that 2px; some people would call that 2py. This you could view as the in-and-out of the page, so you could view that as the z dimension—that is 2pz. The orbitals keep going; there is a d orbital once you get to the third shell, and once you get to the fourth shell, there is an f orbital.

All we've talked about right now is in hydrogen. If we keep giving more energy to that one electron, what happens to it? What is the shape of the probabilities of where it might be in three-dimensional space? As you can imagine, if you have two electrons, it's not exactly the same, but this is a pretty good approximation.

You can actually put two electrons in this one s orbital, but after that, you can imagine the electrons are repelling each other, so another electron doesn't want to go there. The third electron that you add is going to end up in the 2s orbital; it's going to be at that higher energy level, and then that can fit two.

You can fit up to four electrons between the 1s and then the 2s, and then the fifth one is going to have to go into one of these p orbitals. Now, one last point is to make sure you understand the terminology of orbitals and shells.

First of all, you have this idea of shells, and sometimes the word shell will be used interchangeably with energy level. Energy levels, and so in this diagram or this visualization right over here, I've depicted the one shell and then I've also depicted the two shell. This is a shell right over here; this is another shell.

Now you'll also hear the term subshell. Subshell, or sometimes people will say sublevels; that's where they're talking about s or p or d and eventually f. So if I circle this, I'm talking about that first shell.

Now the first shell only contains one subshell, and that's the 1s subshell, and the 1s subshell only has one orbital—once again, the 1s orbital. So for the first shell, the shell, the subshell, and the orbital are all referring to the same thing.

But as we get to the second shell, it's a little bit different. If we talk about the subshells in the second shell, there’s s and p. So this is a subshell, and then this is another subshell right over here. There are actually three orbitals in the p subshell.

So I'll leave you there. In the next video, we'll actually look at various atoms and think about their electron configurations—where do their electrons sit in which of these shells, subshells, and orbitals?