Periodic table (electron configurations)

Transcription

In the last couple of videos we figured out the electron configurations for atoms that only had electrons in the s and p subshells. And so we have this obvious problem. We also have the d subshell, which we'll talk about, here, these bizarre shapes. And then eventually you even get into the f subshells, which are these really kind of exotic-looking shapes. And the shapes, they're interesting to look at and think about, but they're not as important for actually figuring out the configuration. So the question arises, what happens when we start going to the d and f subshells? So the general way to think about it is the energy shell you're in is equivalent to the period. We are in the periodic table. So, just so it all fits on one page. The periods were written out here to left, but then I wouldn't be able to finish the whole table. So this is period one. Let me write this in a darker color. So period one, two, three, four, five, six. I think I barely am fitting on the page. Right? So each row is a period. And then for the purposes of figuring out electron configuration-- we did this in the last video-- we want to put helium-- let me just copy and paste exactly helium-- we want to put helium in the s-block. So we want to put helium right there. The reason why, just in case you're curious of why helium is put there in the periodic table, it's because it has very similar properties to the other elements in this column or this group. Each column is called a group. And we'll talk about valence electrons and why that leads to different properties. But for electron configuration purposes, we can put it in the s-block. That's not too hard to remember because it's just one element. And it kind of makes sense. 1s1, 1s2, et cetera. And what you do is you draw blocks around them. So I've said multiple times already that this right here is the s-block. This over here on the right is the p-block. That's the p-block. And then this in the middle right here is the d-block. And so, if you want to figure out the electron configuration of any atom, the way you think about it, they fill in this order, but when you go from calcium, calcium would have filled out the 4s2, right? 4s1. 4s2. So if I just do its fourth energy shell, it looks like this. Calcium is 4s2. And then you start filling the d-block, right? What did I say? I wanted to do-- so that's calcium. So now I want to figure out the electron configuration for iron, right? Which is in the d-block. So it turns out-- and this is kind of an artifact and I'll do a little bit more of a detailed video on this in the future-- that it actually goes and backfills the third energy shell, because all of a sudden the d orbitals can kind of fit in the gaps of the third energy shell. So what you do is, you go one energy shell above it. So whatever period you're in the d-block, you go that period minus one to figure out what energy shell the d-block is filling. So iron has one, two, three, four, five, six elements in the d-block. So it's going to have d6. But it's not going to be 4d6. It's going to be 3d6. And I figured that out because it's in the fourth period and I subtracted 1 from that. So this is kind of the highest energy eight electrons in iron, right? 4s2. 3d6. If I said, what are the electrons that are in the outermost energy shell? I would say that there are two electrons in the outermost energy shell for iron. But if I were to say, which energy shell has the highest energy electrons, it would be these. Let me actually do the whole electron configuration. Let me pick up another one. Let me take, I don't know, copper, right here. Let me do copper. So the highest energy electrons it has are going to be one, two, three, four, five, six, seven, eight, nine. Actually, let me not do copper, because copper does something very interesting in real life. So it actually is one of the few things that kind of is a special case. Let me do a different one. Let me do the whole thing for iron. Sorry to be waffling around so much. If you wanted to do the entire electron configuration for iron, it would be 1s2. That's the first energy shell. soon. Let me do that in magenta, right there. 1s2. And then in, let's say, orange, then you have 2s2. And then you have six in the p section right there. So 2p6. Now we're in the third energy shell. Let me go switch to this bluish color. So then I fill up 3s2. Remember, this is the s-block. Then I fill out 3p6. Fill out those, right there. Right? One, two, three, four, five, six. And now I'm going to add these electrons. Let me pick a nice green. So then I go to 4s2. So it's 4s2. And now this was the interesting thing, that this whole d-block is interesting. Now I fill out another d-block. Or my first d-block. One, two, three, four, five, six. But it won't be in the fourth energy shell. It'll be in the fourth minus one energy shell. It'll be in the third energy shell. So this will go to 3d6, just like we did at the beginning of the video. And so it's in the third energy shell, so I would actually write it here. 3d6. So if I wanted to write things in order of which energy shell they are, I could have written it this way. If I wanted to write it in order of the highest energy electrons. Remember, the shells are kind of the best way to visualize how far away we are from the nucleus. So in this case, these higher energy electrons are going to be further in the nucleus even though it's a higher energy state to be in. If I did it in terms of energy state, I could rearrange these two. But in most of chemistry, what matters is what's in the outer shell. So it's interesting that although we filled our 4s2 here, and then we kept adding more and more electrons, those electrons were just filling a lower energy shell. So in this atom, in the case of iron, when we talk about the electrons in the outer energy shell, and those are valence electrons, and these are the ones that react. So that these are called-- let me do this in a better color-- valence electrons. This iron has two valence electrons, because the outer shell is 4s2. Even though it had these. Even after filling 4s2, it had six more electrons, but those kind of backfilled the third energy shell. So that's one way. And then, so you might say, well what happens when we go to the f shell or the f-block? And so that's these down here. So in a lot of periodic tables you see these lanthanides and actinides down here. And they're supposed to fill in the gap right here. And that might be a little hard to visualize. And I'll show you why they do that. You could have just as easily made a periodic table that looks like this. Where you insert them in, where you push everything to the right, and you insert these in. But obviously this type of periodic table is a lot harder to fit in. You could have done the same thing with the d-block, actually. So in this one, this is the s-block, this is the f-block, and this is the d-block. And then this is the p-block, right here. When we were dealing with the f-block-- so let's say we wanted to figure out, I don't even know what element this is, electron configuration for this atomic symbol La. So it's filling out this last incremental electron. It fills the f-block. Maybe I should do it in lower case. So it has one in the f orbital and this is the sixth period, but with the f-block you subtract 2. So you subtract 2 from it. So it will be 4f1. And then 6s2. Right? The s-block you just look at the period. 6s2. And then if you were to keep going back, you would then go to 5p6. So then it would be 5p6. And then it would fill out these 10 in the d-block, right there, that are in the fifth period, but remember you subtract 1 from the d-block. So it would be 4d10. And then it's 5s2. And you just keep going back that way. And it seems complicated at first, but just remember, when you're in the s- or the p-block you just look at the period you're in, but then when you start filling the d-block, it fills in a-- this is the d-block-- it fills in a subshell that's one lower. And when you start filling the f-block, which are these really large elements, you start filling a subshell that is two lower. And so maybe in the next video I'll do a couple of these electron configurations, because I think I'm already out of time. And I'll actually show you another way to figure this out that's often covered in some chemistry classes. See you