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 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 in the periodic table. So if we were to do, just so it all fits on one page, the periods were written out here to the left, but then I wouldn't be able to finish the whole table.
So this is period 1. Let me write this in a darker color. So period 1, 2, 3, 4, 5. Six. I think I barely am fitting on the page. 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 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 is 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. And 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 let's see, I've said multiple times already that this is 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. This is the d block. And so if you want to figure out the electron configuration of any atom, you just have to see. The way you think about it, they fill in this order, but when you say, let's say, 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. It's 4s2. And then you start filling the d block, right?
Oh, what did I say? I wanted to do. So that's calcium. This is how I wanted 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, of that it actually goes and back fills 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 in the d block. You go that period minus 1 to figure out what energy shell the d block is filling. So iron has 1, 2, 3, 4, 5, 6 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 the highest energy eight electrons in iron. 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 cell has the highest energy electrons, it would be these. Let me actually do the whole electron configuration for, let me pick up another one just so it's a For, let me take, I don't know, what was this? This is copper right here. Let me do copper right there.
So the highest energy electrons it has are going to be 1, 2, 3, 4, 5, 6, 7, 8, 9. Actually, let me not do copper, because copper does something very interesting in real life. So actually, it's one of the few things that kind of is a special case. Let me do a different one. Let me do, well, let me do.
I don't know. 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. Now the second, let me do that in magenta right there. 1s2. And then in, let's say, in orange, then you have 2s2.
And then you have 6 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. 1, 2, 3, 4, 5, 6. And now I'm going to add these electrons. Then I add, 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, it's one, 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'll go to three d six, 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. I could write it, if I wanted to write three d six. 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 it, 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 called valence electrons, and these are the ones that react. So these are called 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 6 more electrons, but those kind of backfilled the third energy shell.
So that's one way. And then, so you might say, oh, well, what happens when we go to the F-shell, or the F-block? So that's these down here. So a lot of periodic tables, you see these lanthanoids and actanoids 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. 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. This is the p block.
And when you're dealing with the f block, So let's say we wanted to figure out, I don't even know what element this is, the electron configuration for this atomic symbol, LA. So it's filling out this last incremental electron, fills the f block. So it's f, 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'll 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 subcell.
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 soon.