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In the last couple of videos we figured out the electron

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configurations for atoms that only had electrons in the s

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and p subshells.

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And so we have this obvious problem.

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We also have the d subshell, which we'll talk about, here,

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these bizarre shapes.

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And then eventually you even get into the f subshells,

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which are these really kind of exotic-looking shapes.

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And the shapes, they're interesting to look at and

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think about, but they're not as important for actually

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figuring out the configuration.

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So the question arises, what happens when we start going to

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the d and f subshells?

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So the general way to think about it is the energy shell

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you're in is equivalent to the period.

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We are in the periodic table.

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So, just so it all fits on one page.

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The periods were written out here to left, but then I

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wouldn't be able to finish the whole table.

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So this is period one.

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Let me write this in a darker color.

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So period one, two, three, four, five, six.

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I think I barely am fitting on the page.

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Right?

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So each row is a period.

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And then for the purposes of figuring out electron

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configuration-- we did this in the last video-- we want to

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put helium-- let me just copy and paste exactly helium-- we

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want to put helium in the s-block.

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So we want to put helium right there.

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The reason why, just in case you're curious of why helium

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is put there in the periodic table, it's because it has

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very similar properties to the other elements in this column

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or this group.

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Each column is called a group.

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And we'll talk about valence electrons and why that leads

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to different properties.

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But for electron configuration purposes, we can

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put it in the s-block.

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That's not too hard to remember because

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it's just one element.

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And it kind of makes sense.

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1s1, 1s2, et cetera.

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And what you do is you draw blocks around them.

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So I've said multiple times already that this right here

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is the s-block.

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This over here on the right is the p-block.

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That's the p-block.

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And then this in the middle right here is the d-block.

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And so, if you want to figure out the electron configuration

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of any atom, the way you think about it, they fill in this

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order, but when you go from calcium, calcium would have

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filled out the 4s2, right?

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4s1.

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4s2.

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So if I just do its fourth energy shell,

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it looks like this.

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Calcium is 4s2.

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And then you start filling the d-block, right?

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What did I say?

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I wanted to do-- so that's calcium.

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So now I want to figure out the electron configuration for

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iron, right?

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Which is in the d-block.

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So it turns out-- and this is kind of an artifact and I'll

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do a little bit more of a detailed video on this in the

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future-- that it actually goes and backfills the third energy

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shell, because all of a sudden the d orbitals can kind of fit

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in the gaps of the third energy shell.

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So what you do is, you go one energy shell above it.

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So whatever period you're in the d-block, you go that

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period minus one to figure out what energy shell the d-block

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is filling.

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So iron has one, two, three, four, five, six elements in

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the d-block.

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So it's going to have d6.

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But it's not going to be 4d6.

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It's going to be 3d6.

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And I figured that out because it's in the fourth period and

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I subtracted 1 from that.

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So this is kind of the highest energy eight electrons in

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iron, right?

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4s2.

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3d6.

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If I said, what are the electrons that are in the

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outermost energy shell?

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I would say that there are two electrons in the outermost

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energy shell for iron.

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But if I were to say, which energy shell has the highest

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energy electrons, it would be these.

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Let me actually do the whole electron configuration.

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Let me pick up another one.

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Let me take, I don't know, copper, right here.

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Let me do copper.

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So the highest energy electrons it has are going to

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be one, two, three, four, five, six, seven, eight, nine.

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Actually, let me not do copper, because copper does

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something very interesting in real life.

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So it actually is one of the few things that kind of is a

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special case.

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Let me do a different one.

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Let me do the whole thing for iron.

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Sorry to be waffling around so much.

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If you wanted to do the entire electron configuration for

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iron, it would be 1s2.

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That's the first energy shell. soon.

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Let me do that in magenta, right there.

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1s2.

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And then in, let's say, orange, then you have 2s2.

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And then you have six in the p section right there.

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So 2p6.

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Now we're in the third energy shell.

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Let me go switch to this bluish color.

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So then I fill up 3s2.

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Remember, this is the s-block.

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Then I fill out 3p6.

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Fill out those, right there.

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Right?

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One, two, three, four, five, six.

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And now I'm going to add these electrons.

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Let me pick a nice green.

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So then I go to 4s2.

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So it's 4s2.

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And now this was the interesting thing, that this

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whole d-block is interesting.

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Now I fill out another d-block.

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Or my first d-block.

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One, two, three, four, five, six.

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But it won't be in the fourth energy shell.

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It'll be in the fourth minus one energy shell.

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It'll be in the third energy shell.

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So this will go to 3d6, just like we did at the beginning

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of the video.

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And so it's in the third energy shell, so I would

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actually write it here.

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3d6.

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So if I wanted to write things in order of which energy shell

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they are, I could have written it this way.

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If I wanted to write it in order of the

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highest energy electrons.

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Remember, the shells are kind of the best way to visualize

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how far away we are from the nucleus.

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So in this case, these higher energy electrons are going to

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be further in the nucleus even though it's a higher energy

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state to be in.

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If I did it in terms of energy state, I could

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rearrange these two.

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But in most of chemistry, what matters is what's

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in the outer shell.

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So it's interesting that although we filled our 4s2

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here, and then we kept adding more and more electrons, those

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electrons were just filling a lower energy shell.

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So in this atom, in the case of iron, when we talk about

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the electrons in the outer energy shell, and those are

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valence electrons, and these are the ones that react.

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So that these are called-- let me do this in a better color--

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valence electrons.

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This iron has two valence electrons, because the outer

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shell is 4s2.

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Even though it had these.

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Even after filling 4s2, it had six more electrons, but those

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kind of backfilled the third energy shell.

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So that's one way.

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And then, so you might say, well what happens when we go

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to the f shell or the f-block?

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And so that's these down here.

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So in a lot of periodic tables you see these lanthanides and

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actinides down here.

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And they're supposed to fill in the gap right here.

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And that might be a little hard to visualize.

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And I'll show you why they do that.

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You could have just as easily made a periodic table that

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looks like this.

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Where you insert them in, where you push everything to

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the right, and you insert these in.

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But obviously this type of periodic table is a lot

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harder to fit in.

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You could have done the same thing with

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the d-block, actually.

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So in this one, this is the s-block, this is the f-block,

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and this is the d-block.

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And then this is the p-block, right here.

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When we were dealing with the f-block-- so let's say we

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wanted to figure out, I don't even know what element this

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is, electron configuration for this atomic symbol La.

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So it's filling out this last incremental electron.

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It fills the f-block.

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Maybe I should do it in lower case.

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So it has one in the f orbital and this is the sixth period,

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but with the f-block you subtract 2.

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So you subtract 2 from it.

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So it will be 4f1.

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And then 6s2.

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Right?

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The s-block you just look at the period.

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6s2.

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And then if you were to keep going back, you

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would then go to 5p6.

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So then it would be 5p6.

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And then it would fill out these 10 in the d-block, right

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there, that are in the fifth period, but remember you

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subtract 1 from the d-block.

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So it would be 4d10.

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And then it's 5s2.

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And you just keep going back that way.

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And it seems complicated at first, but just remember, when

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you're in the s- or the p-block you just look at the

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period you're in, but then when you start filling the

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d-block, it fills in a-- this is the d-block-- it fills in a

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subshell that's one lower.

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And when you start filling the f-block, which are these

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really large elements, you start filling a subshell that

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is two lower.

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And so maybe in the next video I'll do a couple of these

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electron configurations, because I think I'm already

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out of time.

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And I'll actually show you another way to figure this out

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that's often covered in some chemistry classes.

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See you