 Let's think a little bit about the notion of atomic size or atomic radius in this video. And at first thought, you might think, well, this might be a fairly straightforward thing. If I'm trying to calculate the radius of some type of circular object, I'm just thinking about, well, what's the distance between the center of that circular object and the edge of it? So the length of this line right over here, that would be the radius. And so a lot of people, when they conceptualize an atom, they imagine a positive nucleus with the protons in the center right over here. And then they imagine the electrons on these fixed orbits around that nucleus. So they might imagine some electrons in this orbit right over here, just kind of orbiting around. And then there might be a few more on this orbit out here orbiting around, orbiting around out here. And you might say, OK, well, that's easy to figure out the atomic radius. I just figure out the distance between the nucleus and the outermost electron. And we could call that the radius. And that would work except for the fact that this is not the right way to conceptualize how electrons, or how they move, or how they are distributed around a nucleus. Electrons are not in orbits the way that planets are in orbit around the sun. And we've talked about this in previous videos. They are in orbitoles, which are really just probability distributions of where the electrons can be. But they're not that well-defined. So you might have an orbitole. And I'm just showing you in two dimensions. It would actually be in three dimensions. Or maybe there's a high probability that the electrons where I'm drawing it in kind of this more shaded in green. But there's some probability that the electrons are in this area right over here. And some probability that the electrons are in this area over here. And let's say even a lower probability that the electrons are over here, like this over here. And so you might say, well, OK, at a moment, the electron's there. The outermost electron, let's say, is there. You might say, well, that's the radius. But in the next moment, there's some probability. It might be more likely that it ends up here. But there's some probability that it's going to be over there. And then the radius could be there. And so electrons, these orbitals, these diffuse probability distributions, they don't have a hard edge. So how can you say what the size of an atom actually is? And there's several techniques for thinking about this. One technique for thinking about this is saying, OK, if you have two of the same atoms of the same element that are not connected to each other, that are not bonded to each other, that are not part of the same molecule, and you were able to determine somehow the closest that you could get them to each other. So without them bonding. So you were kind of seeing, what's the closest that they can kind of get to each other? So let's say that's one of them, and then this is the other one right over here. And if you could figure out that distance, that closest, that minimum distance, without some type of really strong influence happening here, but just the minimum distance that you might see between these two, and then you could take half of that. So that's one notion, and that's actually called the van der Waals radius. Another way is, well, what about if you have two atoms of the same element that are bonded to each other? They're bonded to each other through a covalent bond. So a covalent bond, we've seen this in the past. The most famous of covalent bonds is a covalent bond. You essentially have two atoms. So that's the nucleus of one. That's the nucleus of the other. And they're sharing electrons. So their electron clouds actually overlap with each other. So in a covalent bond, the electrons in that bond can spend some of their time on this atom and some of their time on this atom right over here. And so when you have a covalent bond like this, you can then find the distance between the two nuclei and take half of that and call that the atomic radius. So these are all different ways of thinking about it. With that out of the way, let's think about what the trends for atomic size or atomic radii would be in the periodic table. So the first thing to think about is what you think will be the trend for atomic radii as we move through a period. So let's say we're in the fourth period and we were to go from potassium to to krypton. What do you think is going to be the trend here? And if you want to think about the extremes, how do you think potassium is going to compare to krypton in terms of atomic radius? I encourage you to pause this video and think about that on your own. Well, when you're in the fourth period, the outermost electrons are going to be in your fourth shell. Now, what's going on there? Well, when you're at potassium, you have 19 electrons, 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, 19. You have 19 protons, and you have 19 electrons. I'll just do it as 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18. But you only have one electron in that outermost in that fourth shell. So let's just say that's that electron at a moment, just for visual. It doesn't necessarily have to be there, but just to visualize that. So that one electron right over there, you have 19 protons. So you have some, as you could say, Coulomb force that is attracting it, that is keeping it to there. But if you go to krypton, all of a sudden you have much more positive charge in the nucleus. So you have 1, 2, 3, 4, 5, 6, 7, 8. I don't have to do them all. You have 36. You have a positive charge of 36. Let me write that. Plus 36. Here you had plus 19. And you have 36 electrons, and I've lost track of them. But in your outermost shell, in your fourth, you're going to have the 2s, and then you're going to have the 6p. So you have 8 in your outermost shell. So let me 1, 2, 3, 4, 5, 6, 7, 8. So one way to think about it, you have more positive charge in the center, and you have more negative charge on that outer shell. So that's going to bring that outer shell inward. It's going to have more, I guess you could imagine, one way, more Coulomb attraction right over there. And because of that, that outermost shell is going to be drawn in. Krypton is going to be smaller. It's going to have a smaller atomic radius than potassium. And so the trend, as you go to the right, is that you are getting, and the general trend, I would say, is that you are getting smaller as you go to the right in a period. And that's the reason why the smallest atom, the element with the smallest atom, is not hydrogen. It's helium. Helium is actually smaller than hydrogen, depending on what technique you use to measure it. And that's because, if we take the simplest case, hydrogen, you have one proton in the nucleus, and then you have one electron in that 1s shell. And in helium, you have two protons in the nucleus. And I'm not drawing the neutrons, and obviously there's different isotopes, different numbers of neutrons. But you have two electrons now in your outer most shell. And so you have more, I guess you could say, you could have more Coulomb attraction here. This is plus 2. And then these two combined are negative 2. They are going to be drawn inward. So that's the trend. As we go to the right, as we go from the left, to the right of the periodic table, we're getting smaller. Now what do you think is going to happen as we go down the periodic table? As we go down the periodic table in a given group? Well, as we go down a group, each new element down the group, we're in a new period. We're adding a new shell. So you're adding more and more and more shells. Here you have just the first shell, now the second shell. And each shell is getting further and further and further away. So as you go down in the periodic table, you are getting larger. You're having a larger atomic radius depending on how you are measuring it. So what's the general trend? Well, if you're going larger as you go down, that means you're getting smaller as you go up. You get smaller, smaller as you go up. So what's going to be the smallest ones? Well, we've already said helium is the smallest. So what are going to be some of the largest atoms? What's going to be the atoms down here in the bottom left? So these are going to be large. These are going to be small. So large over here, small over here. And the general trend is you go from the bottom left to the top right, you are getting smaller.