All right, our last topic for this chapter about atoms and elements is talking about a couple periodic trends. And what do we mean about periodic trends? It means things that happen a s we go across the periodic table. What happens to size? What happens to other things? An so we're first going to talk about size. It it makes a little bit of sense. Let's talk about lithium. What do we know about lithium? It's configuration, we decided, is 2, 1 versus sodium was 2, 8, 1 and so if we draw these levels ... right, we draw them as circles around the nucleus. And there's two electrons in here for lithium and one electron in here for lithium and we can do the same thing for our sodium. And we can draw the configuration for sodium. Same thing 1, 2, & 3 but for sodium it's 2, 8, (counting) 1, 2, 3, 4, 5, 6, 7, 8 and 1. Right, and I didn't draw this very well, but if I drew this very well you could kind of see that since the outermost electron on lithium is only down at level 2 and the outermost electron on sodium is at level 3, and those shells get bigger and bigger that sodium is just going to be bigger than lithium because there's an electron further away. Because how we measure atomic size is we measure to the outermost electron. From the nucleus to the outermost electron. And so this is a bigger distance there than this is there. And we say that means sodium is bigger than lithium as you can see from this picture. So as we go down the periodic table that happens in every case. As you go down the periodic table; size goes up ... Size goes up as we go down the periodic table ... And again that's because we're adding that extra shell. So what happens as we go across the periodic table? Well, as we go across the periodic table it's a little bit different. And so let's look at, say, carbon and nitrogen. Looking at that picture you're like 'I can't really tell the difference between those two much they look pretty much the same,' but let's go ahead and draw them. We're gonna draw carbon ... with the configuration of 2, 4. And we're gonna draw nitrogen with the configuration of 2, 5. And I'm just drawing those electrons anywhere I feel like. Ok. So we're drawing that here and we see those electrons there. And you know on the first one we said when we get to a new shell right and gets bigger but here we're not getting to a new shell we're actually adding them to the same thing. And so you know if you just start seating more people in the same row in the theater, it's not like the furthest person's any further away. So at first guess you might be like, you know what, maybe the size doesn't change as we go across the periodic table. Because we're not adding to a new row, we're just adding people in the same row. It doesn't make things any bigger and that's actually a great first guess. It turns out there's a second thing going on in here. In our little nucleus. Our little nucleus right here is different in each one of those. What is it? In carbon our nucleus has 6 protons and in nitrogen our nucleus has 7 protons. So if you're one of these electrons way out here, you know this electron here, and you look at the nucleus in carbon, what do you see? You see 6 protons and you want to be near those 6 protons because you're a negatively charged electron and those are positively charged protons, you want to be near them. What do you see if you are a proton in nitrogen? Right, what do you see? Well you see seven protons. Okay. Seven protons! That's better than six protons. I want to be close to those seven protons. And so they kind of look around and when no one's looking they dance just a little bit closer, and they hop a little bit closer, making that ring just a tiny bit smaller than it would have been. And so this guy here wants to be closer ... because it's being pulled in by seven protons instead of just six. And so the six might make it like this. But when you have seven, they squeeze in a little bit because there's more pulling those negative electrons in. And so what we find is that as we go across the periodic table ... as we go this way, size goes down. Atoms actually get smaller as you go across the periodic table. So nitrogen is smaller than carbon, and fluorine is smaller than carbon, and neon is smaller than carbon. But as you go down the periodic table like we said, they get bigger because you're adding electrons to a new shell. They're just getting bigger, and bigger, and bigger. So that's the trend for atomic size on the periodic table. Yeah, what's our other trend? They're going to look at we're gonna look at a trend called ionization energy. What is ionization energy? It's the energy required ... to remove ... the easiest electron. And what do I mean by the easiest electron? Well, remember they're in these different levels, right? Anybody ever played football? What do you do when you catch the ball? When you catch the ball, you don't hold it like way out in front of you and run down the field like this. You're tuck it in tight because when you tuck it in tight it's really hard for people to get. When you leave it right out here it's likely to fall, it's likely to get away. So it turns out when we're talking about ionization energy, we're always talking about the energy of removing one of the farthest, farthest out electrons. Now when we think about something, right, we've got our positively charged nucleus here. And, whoopes, we've got, let's see, our positively charged nucleus here. And let's say we had an electron in that first shell and we're trying to remove that. It's really close to the nucleus, right? So it's going to be kind of hard to remove. And in another example we've got lots of different shells. In fact we're going to talk about an electron out here and like the third or fourth shell. It's much further away from the nucleus. It's the football being held way out here. So is it going to be easier or harder to get that one? Yeah it's gonna be easier, because it's further away. We don't have to pull it as far. Imagine magnets, right, when they're close together they're hard to pull away. When they're a little bit further they're a lot easier to pull away. So it still takes energy, we're still pulling a negative electron from something it wants to be close to. But we're now pulling it a little bit easier. So this guy is going to have a lower ... ionization energy ... And this one's gonna have a higher ... ionization energy. Now, you'll notice that my conversation here and my reasoning depends all on size. Right? And what do we learn about size? As you go down the periodic table these get bigger, and bigger, and bigger. So lithium, small, sodium's bigger, potassium bigger, rubidium is bigger. But as they get bigger should it get easier or harder and take its electrons away? You're right, it should get easier. And so it's easier to take an electron from a from rubidium than it is to take an electron from lithium. Because rubidium is huge, it's farthest electron is really far away. You just kind of walk over and you pick it off. Lithiums electrons are tucked in, they're really close to that nucleus it's really hard to pull them off. Okay so as we go down the periodic table size gets bigger, but ionization energy decreases. And that's what it says here. Ionization energy decreases as you go down the periodic table. They're opposite trends. What did we see in terms of size as we went across the periodic table? If you went from carbon to nitrogen to oxygen to fluorine we saw that the atoms got smaller and smaller and smaller and smaller. So what do you think happens to ionization energy as they get smaller and smaller and smaller? It's going to get slightly harder and harder and harder to remove those electrons. Because they're closer to the nucleus you have to pull them further away. And so as we go across the periodic table, ionization energy increases. So they're opposite trends. Size increases as you go down, decreases as you go across. Ionization energy increases as you go across, and decreases as you go down. So if you remember one, I just remember the other ones the opposite, you're all good. All right, thanks so much.