Electron Configuration - Hi. It's Paul Andersen and welcome to chemistry essentials video 5. This is on electron configurations. When I took chemistry I remember having to do electron configurations. I remember this chart and these different orbitals. And this diagram. And I learned how to do it. But I never really knew what I was doing. And so in this video I'm going to show you how to do electron configurations. We'll do that at the end. But I also really want to tell you what's going on behind the scenes. And so all electron configuration is is simply the distribution of electrons. So where the electrons are found in atoms or in ions. And so a good way to figure that out is to look at their ionization energy which is going to be the amount of energy it takes to pull an electron away. And so we can quantify that by using Coulomb's Law. And so what we can do is we can work our way out from the electrons on the inside to those on the outside. And when we're done we have a pretty good picture of where all of those electrons are. Because in most atoms, they're going to be multi electron. They have many electrons. What we'll find is there they're going to be organized into shells, subshells and then orbitals. And all of those are going to have similar ionization energy. And there's a chart on ionization energy that's the most important one. Now if we look at these orbitals and subshells what we'll find is that the ones on the inside are going to be, or the inner electrons are going to be called core electrons. And the ones on the outside are going to be called valence electrons. And what's interesting is that those on the inside, once they're filled will actually shield the valence electrons from the power of the nucleus itself. And so let's look at the first and simplest of electrons. We've got hydrogen right here. One proton. One electron. And so Coulomb's Law allows us to quantify the force between the two. And so it basically comes down to the charges and then the distance between the charges. And so if we were to look at ionization energy, it's the amount of energy it takes to pull that electron away. How do we figure it out? You're simply going to multiply the two charges, so let's say the proton has a positive charge. Let's call that plus 1. And then electron has a negative charge, let's call that negative 1. And the only other thing we really need is the radius, the distance between the two. And so let me ask you a quick question. Which do you think would have a higher ionization energy? Hydrogen. It's got one proton, one electron. Or helium. It's going to have two of each. And two neutrons as well. So make a guess. And let me show what helium looks like. Helium, you can see it just got a little bit smaller and the reason why is since it has two protons in the middle there's going to be more pull or there's going to be a larger energy of ionization in helium. And that's because there are a greater amount of positive charge on the inside. So we would say helium has a higher ionization energy. What about the next one. What if we were to compare helium then with lithium? What do you think happens to the ionization energy there? Well you might immediately think it's going to be bigger, but remember the electrons are organized with only two electrons in that first shell. And then when I learned it was going to be eight in that next subshell. And so if we look at that lithium, what we'll find is we've increase the number of protons. So there's more of a charge here. But that electron is going to be in the next shell. And so since it's in the next shell it's shielded from the power of that proton, the protons in the nucleus itself. And so it's going to have a really small ionization energy here just because of the shielding effect and the fact that we have another shell. And so what do you think about beryllium? Well hold on to that and we'll take a look at the pattern. And so what this is is the chart that really allows you to understand what's going on with the electrons. And so let me lay it out for you. We have atomic number across the bottom. So this would be hydrogen at 1. Helium at 2. All the way out. And then we're going to have ionization energy, which is the amount of energy it takes to pull that electron away. And what you can see are patterns immediately. And so these big spikes that you can see are ending with these noble gases, those are going to be those different shells. You can see that they're broken down into subshells. And then we're going to get to orbitals in just a second. And so let's kind of lay this out for just a second. So a shell is going to be, let's look at the first one. So here we've got hydrogen. So hydrogen is going to be right here. And we said it had a high ionization energy and that's because it's really close to that nucleus. So it's close to the protons inside there. And so what we can do is we can fill out this box. So this is going to be the 1s, we would call that 1 is just going to have a shell and this is going to be the 1s shell. But the reason we have boxes here is that it was later discovered that electrons will fill what are called orbitals. And those orbitals can only have two electrons in each one. So let's look at hydrogen. Hydrogen we said has an ionization energy we'll say of on this chart around 12. If we go to helium it's going to have a higher ionization energy. Why is that? It's going to be because remember we have those protons on the inside that are holding it in. But let's watch what happens when we go to lithium. What's happened is we've gone into this whole next orbital. And so there's this shielding effect of this filled orbital here. And so we have much less ionization energy there. Let's watch what happens with beryllium. Well that kind of makes sense. We're increasing the protons. And so that should go up. But then what happens with the next one. Well, when we go to boron what's happening is it's actually going into a new orbital. And so it's going into an orbital of the second subshell. And so if we didn't have these orbitals, what would happen is it would just be a consistent all the way up to neon. But you can see we kind of have this jag as we move up. If we go to the next one, that makes sense. Carbon's going to have more ionization energy. We're adding more protons. Nitrogen same way. But now let's look at this. Why does the ionization energy go here after we leave nitrogen? Well the reason why is that these electrons are added. And they're added one at a time in each orbital. But they'll have a specific spin. And so you can only have two electrons each way. But what they'll do is they'll fill them kind of like seats in a bus. And so they're all going to get on one seat per bus, but then when we get more electrons they're going to have to double up. And so what happens is when we go to oxygen what we're doing is we're making that electron sit in that orbital. It would rather not be there. There's repulsion between those electrons, and so it lowers the ionization energy. We go to fluorine, we go to neon. Once we have that shield effect we jump all the way down to sodium again. And so once I really understood what was going on in this chart, then all of these subshells made sense. And so where do they come from? It's quantum mechanics. And it's quantum numbers that determine that. But if we're looking at periodicity, in other words if we're looking at the periodic table, these things would be subshells. And so we're going to have a shell which is going to be the 1, 2, 3, 4, 5. But we're also going to have subshells which are smaller. And then within each of those we're going to have orbitals if that makes sense, if we go in size. And so let's get to the electron configuration. And so what we have here are the s, p, the d and the f. They come from quantum mechanics. What's interesting about the s is that it only has one orbital in it. The p has three orbitals. So you could have 6 electrons. d has 5 orbitals so you could have 10 electrons. And f is going to have 7 orbitals so you could have 14 electrons. So if you want to write this out, you could just write on a chart. Start with the 1s. And just write all the way down to 7. Then you want to write right to the write of that 1s or 2s, you're going to write the 2. And then you're just going to write p's all the way down. And then you go down to the 3. We'll write a 3 over here and then you're going to write d's all the way down. It could go all of the way to 7, but we'll never really need that. And then you go to the 4. You write a 4 over here, and then you write f's all the way down. And so once we have that we can put diagonal lines in like this and then we can do the electron configuration for anything. So you wouldn't even need a periodic table. You could just do this and you could figure out electron configurations. Let's start with hydrogen. What's going to be the electron configuration of hydrogen? You're simply going to start at the top. And so it's the 1s shell or orbital or subshell. And so what we would write it as is 1s1. And that 1 represents the 1 electron that we have inside that orbital. If we go to boron then, how are we going to write boron? Well you're going to start here. Go here, go here. Boron is going to be a 1s2 because we can only put 2 electrons in there. We then go to the 2s. It's going to be 2s2. Because we can only put 2 electrons in that. And then we're going to go to the 2p and we only have 1 electron left. So we're going to put it there. So that would be boron. What about neon? And again you could pause the video and always try it out. Neon is going to be this. So it's 1s2 2s2 2p remember could 6. And so we're going to write 6 there. Let's go to the next one. What about sodium which is number 11. We've got 11 electrons. Well we could write it out like this. 1s2 2s2 2p6 and then the next one goes in the 3s1. But you can see how long these electron configurations are going to get. And so we can abbreviate that. And so we could just say it's neon, put that in brackets with a 3s1. So that allows us to write the nobel gas in there and then we can add what's after the nobel gas. And so here's a couple for you to try. Could you try chlorine which as 17 electrons. Or silver which has 47. I'll put the answers in the video descriptions down below. But give it a try. So what did we learn? We learned that an electron configuration is the distribution of what? Again, could you pause the video and figure out what's in the blanks? Distribution of electrons in atoms or ions. They each have different ionization energy. We could quantify that through Coulomb's Law. Remember most atoms are going to have multi electrons. And those are organized in subshells. And then this property is important as well. Those inner or core electrons are going to shield the valence or outside electrons from the power of the nucleus. And so what you should have learned is how the energies of electrons vary within shells of atoms. And so I could point you to this ionization energy chart. And as we walked through that, that should help. And then the other thing you should have learned is we could use Coulomb's Law to analyze measured energies of electrons. And that works great, but it doesn't explain orbitals. And so we have to modify our theory a little bit. And I hope that was helpful.