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In the last few videos we learned that the configuration

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of electrons in an atom aren't in a simple, classical,

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Newtonian orbit configuration.

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And that's the Bohr model of the electron.

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And I'll keep reviewing it, just because I think it's an

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important point.

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If that's the nucleus, remember, it's just a tiny,

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tiny, tiny dot if you think about the entire volume of the

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actual atom.

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And instead of the electron being in orbits around it,

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which would be how a planet orbits the sun.

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Instead of being in orbits around it, it's described by

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orbitals, which are these probability density functions.

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So an orbital-- let's say that's the nucleus-- it would

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describe, if you took any point in space around the

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nucleus, the probability of finding the electron.

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So actually, in any volume of space around the nucleus, it

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would tell you the probability of finding the electron within

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that volume.

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And so if you were to just take a bunch of snapshots of

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electrons-- let's say in the 1s orbital.

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And that's what the 1s orbital looks like.

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You can barely see it there, but it's a sphere around the

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nucleus, and that's the lowest energy state that an

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electron can be in.

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If you were to just take a number of

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snapshots of electrons.

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Let's say you were to take a number of snapshots of helium,

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which has two electrons.

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Both of them are in the 1s orbital.

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It would look like this.

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If you took one snapshot, maybe it'll be there, the next

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snapshot, maybe the electron is there.

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Then the electron is there.

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Then the electron is there.

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Then it's there.

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And if you kept doing the snapshots, you would have a

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bunch of them really close.

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And then it gets a little bit sparser as you get out, as you

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get further and further out away from the electron.

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But as you see, you're much more likely to find the

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electron close to the center of the atom than further out.

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Although you might have had an observation with the electron

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sitting all the way out there, or sitting over here.

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So it really could have been anywhere, but if you take

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multiple observations, you'll see what that probability

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function is describing.

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It's saying look, there's a much lower probability of

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finding the electron out in this little cube of volume

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space than it is in this little cube of volume space.

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And when you see these diagrams that draw this

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orbital like this.

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Let's say they draw it like a shell, like a sphere.

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And I'll try to make it look three-dimensional.

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So let's say this is the outside of it, and the nucleus

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is sitting some place on the inside.

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They're just saying -- they just draw a cut-off -- where

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can I find the electron 90% of the time?

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So they're saying, OK, I can find the electron 90% of the

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time within this circle, if I were to do the cross-section.

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But every now and then the electron can show up outside

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of that, right?

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Because it's all probabilistic.

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So this can still happen.

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You can still find the electron if this is the

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orbital we're talking about out here.

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Right?

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And then we, in the last video, we said, OK, the

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electrons fill up the orbitals from lowest energy state to

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high energy state.

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You could imagine it.

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If I'm playing Tetris-- well I don't know if Tetris is the

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thing-- but if I'm stacking cubes, I lay out cubes from

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low energy, if this is the floor, I put the first cube at

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the lowest energy state.

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And let's say I could put the second cube at a low energy

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state here.

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But I only have this much space to work with.

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So I have to put the third cube at the next highest

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energy state.

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In this case our energy would be described as potential

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energy, right?

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This is just a classical, Newtonian physics example.

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But that's the same idea with electrons.

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Once I have two electrons in this 1s orbital -- so let's

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say the electron configuration of helium is 1s2-- the third

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electron I can't put there anymore, because there's only

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room for two electrons.

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The way I think about it is these two electrons are now

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going to repel the third one I want to add.

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So then I have to go to the 2s orbital.

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And now if I were to plot the 2s orbital on top of this one,

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it would look something like this, where I have a high

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probability of finding the electrons in this shell that's

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essentially around the 1s orbital, right?

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So right now, if maybe I'm dealing with

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lithium right now.

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So I only have one extra electron.

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So this one extra electron, that might be where I observed

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that extra electron.

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But every now and then it could show up there, it could

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show up there, it could show up there, but the high

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probability is there.

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So when you say where is it going to be 90% of the time?

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It'll be like this shell that's around the center.

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Remember, when it's three-dimensional you would

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kind of cover it up.

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So it would be this shell.

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So that's what they drew here.

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They do the 1s.

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It's just a red shell.

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And then the 2s.

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The second energy shell is just this blue shell over it.

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And you can see it a little bit better in, actually, the

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higher energy orbits, the higher energy shells, where

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the seventh s energy shell is this red area.

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Then you have the blue area, then the red, and the blue.

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And so I think you get the idea that each of those are

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energy shells.

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So you kind of keep overlaying the s energy orbitals around

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each other.

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But you probably see this other stuff here.

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And the general principle, remember, is that the

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electrons fill up the orbital from lowest energy orbital to

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higher energy orbital.

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So the first one that's filled up is the 1s.

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This is the 1.

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This is the s.

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So this is the 1s.

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It can fit two electrons.

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Then the next one that's filled up is 2s.

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It can fill two more electrons.

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And then the next one, and this is where it gets

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interesting, you fill up the 2p orbital.

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That's this, right here.

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2p orbitals.

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And notice the p orbitals have something, p sub z, p

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sub x, p sub y.

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What does that mean?

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Well, if you look at the p-orbitals, they have these

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dumbbell shapes.

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They look a little unnatural, but I think in future videos

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we'll show you how they're analogous to standing waves.

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But if you look at these, there's three ways that you

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can configure these dumbbells.

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One in the z direction, up and down.

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One in the x direction, left or right.

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And then one in the y direction, this way, forward

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and backwards, right?

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And so if you were to draw-- let's say you wanted to draw

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the p-orbitals.

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So this is what you fill next.

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And actually, you fill one electron here, another

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electron here, then another electron there.

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Then you fill another electron, and we'll talk about

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spin and things like that in the future.

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But, there, there, and there.

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And that's actually called Hund's rule.

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Maybe I'll do a whole video on Hund's rule, but that's not

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relevant to a first-year chemistry lecture.

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But it fills in that order, and once again, I want you to

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have the intuition of what this would look like.

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Look.

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I should put look in quotation marks,

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because it's very abstract.

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But if you wanted to visualize the p orbitals-- let's say

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we're looking at the electron configuration

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for, let's say, carbon.

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So the electron configuration for carbon, the first two

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electrons go into, so, 1s1, 1s2.

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So then it fills-- sorry, you can't see everything.

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So it fills the 1s2, so carbon's configuration.

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It fills 1s1 then 1s2.

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And this is just the configuration for helium.

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And then it goes to the second shell, which is the second

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period, right?

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That's why it's called the periodic table.

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We'll talk about periods and groups in the future.

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And then you go here.

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So this is filling the 2s.

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We're in the second period right here.

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That's the second period.

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One, two.

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Have to go off, so you can see everything.

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So it fills these two.

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So 2s2.

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And then it starts filling up the p orbitals.

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So then it starts filling 1p and then 2p.

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And we're still on the second shell, so 2s2, 2p2.

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So the question is what would this look like if we just

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wanted to visualize this orbital

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right here, the p orbitals?

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So we have two electrons.

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So one electron is going to be in a-- Let's say if this is,

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I'll try to draw some axes.

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That's too thin.

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So if I draw a three-dimensional

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volume kind of axes.

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If I were to make a bunch of observations of, say, one of

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the electrons in the p orbitals, let's say in the pz

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dimension, sometimes it might be here, sometimes it might be

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there, sometimes it might be there.

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And then if you keep taking a bunch of observations, you're

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going to have something that looks like this bell shape,

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this barbell shape right there.

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And then for the other electron that's maybe in the x

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direction, you make a bunch of observations.

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Let me do it in a different, in a

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noticeably different, color.

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It will look like this.

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You take a bunch of observations, and you say,

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wow, it's a lot more likely to find that electron in kind of

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the dumbell, in that dumbbell shape.

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But you could find it out there.

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You could find it there.

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You could find it there.

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This is just a much higher probability of finding it in

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here than out here.

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And that's the best way I can think of to visualize it.

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Now what we were doing here, this is called an electron

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configuration.

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And the way to do it-- and there's multiple ways that are

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taught in chemistry class, but the way I like to do it-- is

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you take the periodic table and you say, these groups, and

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when I say groups I mean the columns, these are going to

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fill the s subshell or the s orbitals.

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You can just write s up here, just right there.

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These over here are going to fill the p orbitals.

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Actually, let me take helium out of the picture.

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The p orbitals.

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Let me just do that.

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Let me take helium out of the picture.

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These take the p orbitals.

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And actually, for the sake of figuring out these, you should

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take helium and throw it right over there.

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Right?

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The periodic table is just a way to organize things so it

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makes sense, but in terms of trying to figure out orbitals,

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you could take helium.

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Let me do that.

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The magic of computers.

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Cut it out, and then let me paste it right over there.

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Right?

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And now you see that helium, you get 1s and then you get

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2s, so helium's configuration is-- Sorry, you

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get 1s1, then 1s2.

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We're in the first energy shell.

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Right?

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So the configuration of hydrogen is 1s1.

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You only have one electron in the s subshell of the first

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energy shell.

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The configuration of helium is 1s2.

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And then you start filling the second energy shell.

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The configuration of lithium is 1s2.

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That's where the first two electrons go.

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And then the third one goes into 2s1, right?

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And then I think you start to see the pattern.

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And then when you go to nitrogen you say, OK, it has

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three in the p sub-orbital.

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So you can almost start backwards, right?

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So we're in period two, right?

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So this is 2p3.

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Let me write that down.

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So I could write that down first. 2p3.

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So that's where the last three electrons

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go into the p orbital.

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Then it'll have these two that go into the 2s2 orbital.

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And then the first two, or the electrons in the lowest energy

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state, will be 1s2.

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So this is the electron configuration,

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right here, of nitrogen.

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And just to make sure you did your configuration right, what

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you do is you count the number of electrons.

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So 2 plus 2 is 4 plus 3 is 7.

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And we're talking about neutral atoms, so the

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electrons should equal the number of protons.

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The atomic number is the number of protons.

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So we're good.

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Seven protons.

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So this is, so far, when we're dealing just with the s's and

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the p's, this is pretty straightforward.

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And if I wanted to figure out the configuration of silicon,

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right there, what is it?

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Well, we're in the third period.

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One, two, three.

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That's just the third row.

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And this is the p-block right here.

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So this is the second row in the p-block, right?

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

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Right.

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We're in the second row of the p-block, so we

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start off with 3p2.

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And then we have 3s2.

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And then it filled up all of this p-block over here.

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So it's 2p6.

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And then here, 2s2.

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And then, of course, it filled up at the first shell before

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it could fill up these other shells.

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So, 1s2.

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So this is the electron configuration for silicon.

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And we can confirm that we should have 14 electrons.

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2 plus 2 is 4, plus 6 is 10.

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10 plus 2 is 12 plus 2 more is 14.

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So we're good with silicon.

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I think I'm running low on time right now, so in the next

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video we'll start addressing what happens when you go to

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these elements, or the d-block.

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And you can kind of already guess what happens.

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We're going to start filling up these d orbitals here that

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have even more bizarre shapes.

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And the way I think about this, not to waste too much

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time, is that as you go further and further out from

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the nucleus, there's more space in between the lower

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energy orbitals to fill in more of these

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bizarro-shaped orbitals.

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But these are kind of the balance -- I will talk about

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standing waves in the future-- but these are kind of a

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balance between trying to get close to the nucleus and the

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proton and those positive charges, because the electron

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charges are attracted to them, while at the same time

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avoiding the other electron charges, or at least their

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mass distribution functions.

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Anyway, see you in the next video.

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