We're ready to begin our next learning objective in which we're going to learn to write the electron configuration of cations and anions. The rules for doing this are not difficult at all. You just have to keep straight where the electrons started in terms of the atom, and then what it becomes after you either add or take away electrons. Let's begin with our anions. When you do the electron configuration of anions, you will do it for the atom, and then just place the additional electrons in the next available orbitals, so according to the rules that you already know. So for example, if we were going to do O2 minus, we would begin with oxygen, okay, so, electron configuration of oxygen. You're going to want to look at the periodic table and verify that this is what you would have come up with, helium 2s2, 2p4. Now you see here, that it has six valence electrons, and if you look at oxygen's group number, it has got a six associated with it. So we know that it matches up. So now where are you going to put the next electrons? Well, there's still two spots available, in the p orbital. So those two electrons can go into the oxygen, in, into the 2p subshell there, and be up to 2p6. Now you can't get any bigger than six, but that fills up that valence shell. And that gives oxygen, with a two minus charge, the same electron configuration as neon. So you could also write, for the configuration of O2 minus, you could write neon and that would be correct as well. Look at your periodic table. See how neon is two to the right of oxygen? Well that tells you that neon would be an appropriate electron configuration, as well. Now usually, when you're talking about an anion, they are formed from non-metals. The negative charge gives the atom a noble gas configuration. So for, when we look at a periodic table, we can start determining what these non-metals' charges would be if they became an ion. It's how many electrons they would want to gain in order to become like a noble gas, as far as this electron consig, configuration is concerned. Now let's look at cations. Students want to take electrons out and just do electron configuration for the fewer number of electrons, but don't do that. It says in red always, because you really need to do this, because you will find that it doesn't always have to happen, but you won't know the difference. So always write it for the metal, and then take the electrons out. Since it's a cation, we know it's positively charged. We're moving electrons. Take it out of the highest n, p's before s's. So that's the rules that we'll follow. So if we were going to do electron configuration of barium 2 plus, I am telling you always with red letters, always do the electron configuration of just barium. So find barium on your periodic table. Find the noble gas that comes before it. And the electron configuration of barium is xenon, 6s2. Okay? So barium 6s2, now we're ready to remove two electrons. You're removing them from the highest n. Well here we only have one choice of n. We have an n of six. We're going to take those two electrons. So when we get the plus two charge here that we want to consider, we're going to remove these two electrons right here, and that's going to give me barium two plus, just being xenon. So, for this one, if you just did it for two fewer electrons, you'd get xenon. So you would get into this lulled into this false sense that that always works, and it just does not. So let's have you try one. You do the electron configuration of Fe3 plus. Remember what I said with the always. Did you pick argon 3d5? Well if you didn't, that you might've gotten messed up with not doing the always. So what would be the electron configuration of just Fe? Okay? It would be argon, okay? Then it would be 4s2 and then 3d6. So that's the electron configuration of iron. Now we're ready to remove the electrons. We need to remove three electrons. So if you picked number three. You put em in, instead of taking away. That plus sign is not having you add. You're making it positively charged by removing electrons. You need to remove three. If you removed, if you chose this one, you probably removed three from here, but that doesn't follow the rules. The rules say, remove from the highest n. Well which is bigger? Four is bigger than three. So we gotta remove these first. That removes two of them. And then we can remove one of these. And that's how we get to argon 3d5. So you need to do iron first, remove the electrons from the highest n, p before s. This case, we didn't have any p's to remove. There's a term that we use when were talking about electron configurations, and the term is isoelectronic. Isoelectronics are, isoelectronic means atoms and ions which have the same electron configuration. Now they're going to have to have the same number of electrons too, but that's not the definition. You can have two things, maybe with the same elec, number of electrons, and not have the same electron configuration. But if they have the same electron configuration, then it's the term isoelectronic. So Fe minus has this electron configuration, because Fe would be 2s2, 2p5. We put one more electron in, it's 2s2, 2p6. In a previous slide for this learning objective, we did oxygen two minus and it was helium 2s2, 2p6. So those were two things that had the exact same electron configuration. They do have the same number of electrons, obviously you can't have the same electron configuration without that. And, they are isoelectronic. Okay, so now let's have you answer this question. Which of the following is a cation that is isoelectronic with those? If you chose number six, you would be correct. Maybe you chose two, maybe you chose three, but the truth of the matter is that both of those have the exact same electron configuration together. So the answer is two and three. Now you might not have recognized that. We know that F minus and O2 minus are helium 2s2, 2p6, right? And you might have come up with sodium and magnesium, being just neon. But, this is the same as neon. Okay, if you take helium and you work your way over 2s2, 2p6, and you had six electrons in. You're sitting there, where neon is. So you add electrons to obtain a neon noble gas configuration. Sodium, the element was this. And when we remove one electron to become sodium plus, we have neon. Magnesium, the element, was this. And when we remove two electrons, we remove those two and we have neon. You might have said, well why isn't that an answer? All of these have electron configuration of neon, neon would have an electron configuration of neon, and that's just because of this word right here. It says, what is the cation that has the same electron configuration as those two? And neon's not a cation. Okay, so here's the next one. Is the electron configuration of vanadium the same as the electron configuration of iron 3 plus? Now you did iron three plus a little bit ago. They both have 23 electrons, so are they the same? The answer is no. Okay? So why not? Let's make sure we understand why it would be no. Let's pull up a periodic table here. Now let's do the electron configuration of vanadium. Here's vanadium. The noble gas that comes before it is argon, okay? We work our way over to vanadium. We have 4s2, 3d3. Was that the same as your iron 3 plus? Well, it's not. Remember argon, I mean iron is argon 4s2, 3d6, and iron 3 plus, we remove these electrons and give, and take one of those away and we're left with 3d5. So this is an example of two things that have the same number of electrons, but are not isoelectronic. So they have a different electron configuration. So this concludes this learning objective in which we've learned to do the electron configuration of cations and anions, and we've learned what isoelectronic means.
