More on orbitals and electron configuration<br>More intuition on orbitals. Touching on electron configuration.

In the last few videos we learned that the configuration of electrons in an atom aren't in a simple, classical, Newtonian orbit configuration. And that's the Bohr model of the electron.

And I'll keep reviewing it, just because I think it's an important point. If that's the nucleus, remember, it's just a tiny, tiny, tiny dot if you think about the entire volume of the actual atom.

And instead of the electron being in orbits around it, which would be how a planet orbits the sun. Instead of being in orbits around it, it's described by orbitals, which are these probability density functions.

So an orbital-- let's say that's the nucleus-- it would describe, if you took any point in space around the nucleus, the probability of finding the electron.

So actually, in any volume of space around the nucleus, it would tell you the probability of finding the electron within that volume.

And so if you were to just take a bunch of snapshots of electrons-- let's say in the 1s orbital. And that's what the 1s orbital looks like.

You can barely see it there, but it's a sphere around the nucleus, and that's the lowest energy state that an electron can be in.

If you were to just take a number of snapshots of electrons. Let's say you were to take a number of snapshots of helium, which has two electrons.

Both of them are in the 1s orbital. It would look like this.

If you took one snapshot, maybe it'll be there, the next snapshot, maybe the electron is there. Then the electron is there. Then the electron is there. Then it's there.

And if you kept doing the snapshots, you would have a bunch of them really close. And then it gets a little bit sparser as you get out, as you get further and further out away from the electron.

But as you see, you're much more likely to find the electron close to the center of the atom than further out.

Although you might have had an observation with the electron sitting all the way out there, or sitting over here.

So it really could have been anywhere, but if you take multiple observations, you'll see what that probability function is describing.

It's saying look, there's a much lower probability of finding the electron out in this little cube of volume space than it is in this little cube of volume space.

And when you see these diagrams that draw this orbital like this. Let's say they draw it like a shell, like a sphere. And I'll try to make it look three-dimensional.

So let's say this is the outside of it, and the nucleus is sitting some place on the inside.

They're just saying -- they just draw a cut-off -- where can I find the electron 90% of the time?

So they're saying, OK, I can find the electron 90% of the time within this circle, if I were to do the cross-section.

But every now and then the electron can show up outside of that, right? Because it's all probabilistic. So this can still happen.

You can still find the electron if this is the orbital we're talking about out here. Right?

And then we, in the last video, we said, OK, the electrons fill up the orbitals from lowest energy state to high energy state. You could imagine it.

If I'm playing Tetris-- well I don't know if Tetris is the thing-- but if I'm stacking cubes, I lay out cubes from low energy, if this is the floor, I put the first cube at the lowest energy state.

And let's say I could put the second cube at a low energy state here. But I only have this much space to work with.

So I have to put the third cube at the next highest energy state.

In this case our energy would be described as potential energy, right? This is just a classical, Newtonian physics example. But that's the same idea with electrons.

Once I have two electrons in this 1s orbital -- so let's say the electron configuration of helium is 1s2-- the third electron I can't put there anymore, because there's only room for two electrons.

The way I think about it is these two electrons are now going to repel the third one I want to add.

So then I have to go to the 2s orbital. And now if I were to plot the 2s orbital on top of this one, it would look something like this, where I have a high probability of finding the electrons in this shell that's essentially around the 1s orbital, right?

So right now, if maybe I'm dealing with lithium right now. So I only have one extra electron. So this one extra electron, that might be where I observed that extra electron.

But every now and then it could show up there, it could show up there, it could show up there, but the high probability is there.

So when you say where is it going to be 90% of the time? It'll be like this shell that's around the center.

Remember, when it's three-dimensional you would kind of cover it up. So it would be this shell.

So that's what they drew here. They do the 1s. It's just a red shell. And then the 2s. The second energy shell is just this blue shell over it.

And you can see it a little bit better in, actually, the higher energy orbits, the higher energy shells, where the seventh s energy shell is this red area. Then you have the blue area, then the red, and the blue.

And so I think you get the idea that each of those are energy shells. So you kind of keep overlaying the s energy orbitals around each other.

But you probably see this other stuff here. And the general principle, remember, is that the electrons fill up the orbital from lowest energy orbital to higher energy orbital.

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

It can fit two electrons. Then the next one that's filled up is 2s. It can fill two more electrons.

And then the next one, and this is where it gets interesting, you fill up the 2p orbital. That's this, right here. 2p orbitals.

And notice the p orbitals have something, p sub z, p sub x, p sub y. What does that mean?

Well, if you look at the p-orbitals, they have these dumbbell shapes. They look a little unnatural, but I think in future videos we'll show you how they're analogous to standing waves.

But if you look at these, there's three ways that you can configure these dumbbells.

One in the z direction, up and down. One in the x direction, left or right.

And then one in the y direction, this way, forward and backwards, right?

And so if you were to draw-- let's say you wanted to draw the p-orbitals. So this is what you fill next.

And actually, you fill one electron here, another electron here, then another electron there.

Then you fill another electron, and we'll talk about spin and things like that in the future. But, there, there, and there. And that's actually called Hund's rule.

Maybe I'll do a whole video on Hund's rule, but that's not relevant to a first-year chemistry lecture.

But it fills in that order, and once again, I want you to have the intuition of what this would look like. Look.

I should put look in quotation marks, because it's very abstract.

But if you wanted to visualize the p orbitals-- let's say we're looking at the electron configuration for, let's say, carbon.

So the electron configuration for carbon, the first two electrons go into, so, 1s1, 1s2. So then it fills-- sorry, you can't see everything.

So it fills the 1s2, so carbon's configuration. It fills 1s1 then 1s2. And this is just the configuration for helium.

And then it goes to the second shell, which is the second period, right? That's why it's called the periodic table. We'll talk about periods and groups in the future.

And then you go here. So this is filling the 2s. We're in the second period right here. That's the second period. One, two. Have to go off, so you can see everything. So it fills these two. So 2s2.

And then it starts filling up the p orbitals. So then it starts filling 1p and then 2p. And we're still on the second shell, so 2s2, 2p2.

So the question is what would this look like if we just wanted to visualize this orbital right here, the p orbitals?

So we have two electrons. So one electron is going to be in a-- Let's say if this is,

I'll try to draw some axes. That's too thin. So if I draw a three-dimensional volume kind of axes.

If I were to make a bunch of observations of, say, one of the electrons in the p orbitals, let's say in the pz dimension, sometimes it might be here, sometimes it might be there, sometimes it might be there.

And then if you keep taking a bunch of observations, you're going to have something that looks like this bell shape, this barbell shape right there.

And then for the other electron that's maybe in the x direction, you make a bunch of observations.

Let me do it in a different, in a noticeably different, color. It will look like this.

You take a bunch of observations, and you say, wow, it's a lot more likely to find that electron in kind of the dumbell, in that dumbbell shape.

But you could find it out there. You could find it there. You could find it there. This is just a much higher probability of finding it in here than out here.

And that's the best way I can think of to visualize it. Now what we were doing here, this is called an electron configuration.

And the way to do it-- and there's multiple ways that are taught in chemistry class, but the way I like to do it-- is you take the periodic table and you say, these groups, and when I say groups I mean the columns, these are going to fill the s subshell or the s orbitals.

You can just write s up here, just right there. These over here are going to fill the p orbitals. Actually, let me take helium out of the picture. The p orbitals.

Let me just do that. Let me take helium out of the picture. These take the p orbitals.

And actually, for the sake of figuring out these, you should take helium and throw it right over there. Right?

The periodic table is just a way to organize things so it makes sense, but in terms of trying to figure out orbitals, you could take helium. Let me do that.

The magic of computers. Cut it out, and then let me paste it right over there. Right?

And now you see that helium, you get 1s and then you get 2s, so helium's configuration is-- Sorry, you get 1s1, then 1s2. We're in the first energy shell. Right?

So the configuration of hydrogen is 1s1. You only have one electron in the s subshell of the first energy shell.

The configuration of helium is 1s2. And then you start filling the second energy shell.

The configuration of lithium is 1s2. That's where the first two electrons go. And then the third one goes into 2s1, right? And then I think you start to see the pattern.

And then when you go to nitrogen you say, OK, it has three in the p sub-orbital.

So you can almost start backwards, right? So we're in period two, right? So this is 2p3.

Let me write that down. So I could write that down first. 2p3. So that's where the last three electrons go into the p orbital.

Then it'll have these two that go into the 2s2 orbital. And then the first two, or the electrons in the lowest energy state, will be 1s2.

So this is the electron configuration, right here, of nitrogen.

And just to make sure you did your configuration right, what you do is you count the number of electrons.

So 2 plus 2 is 4 plus 3 is 7. And we're talking about neutral atoms, so the electrons should equal the number of protons.

The atomic number is the number of protons. So we're good. Seven protons.

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

And if I wanted to figure out the configuration of silicon, right there, what is it?

Well, we're in the third period. One, two, three. That's just the third row. And this is the p-block right here.

So this is the second row in the p-block, right? One, two, three, four, five, six. Right.

We're in the second row of the p-block, so we start off with 3p2. And then we have 3s2.

And then it filled up all of this p-block over here. So it's 2p6. And then here, 2s2.

And then, of course, it filled up at the first shell before it could fill up these other shells. So, 1s2.

So this is the electron configuration for silicon.

And we can confirm that we should have 14 electrons. 2 plus 2 is 4, plus 6 is 10. 10 plus 2 is 12 plus 2 more is 14. So we're good with silicon.

I think I'm running low on time right now, so in the next video we'll start addressing what happens when you go to these elements, or the d-block.

And you can kind of already guess what happens. We're going to start filling up these d orbitals here that have even more bizarre shapes.

And the way I think about this, not to waste too much time, is that as you go further and further out from the nucleus, there's more space in between the lower energy orbitals to fill in more of these bizarro-shaped orbitals.

But these are kind of the balance -- I will talk about standing waves in the future-- but these are kind of a balance between trying to get close to the nucleus and the proton and those positive charges, because the electron charges are attracted to them, while at the same time avoiding the other electron charges, or at least their mass distribution functions.

Anyway, see you in the next video.

In the last few videos we learned that the configuration of electrons in an atom aren't in a simple, classical, Newtonian orbit configuration. And that's the Bohr model of the electron.

And I'll keep reviewing it, just because I think it's an important point. If that's the nucleus, remember, it's just a tiny, tiny, tiny dot if you think about the entire volume of the actual atom.

And instead of the electron being in orbits around it, which would be how a planet orbits the sun. Instead of being in orbits around it, it's described by orbitals, which are these probability density functions.

So an orbital-- let's say that's the nucleus-- it would describe, if you took any point in space around the nucleus, the probability of finding the electron.

So actually, in any volume of space around the nucleus, it would tell you the probability of finding the electron within that volume.

And so if you were to just take a bunch of snapshots of electrons-- let's say in the 1s orbital. And that's what the 1s orbital looks like.

You can barely see it there, but it's a sphere around the nucleus, and that's the lowest energy state that an electron can be in.

If you were to just take a number of snapshots of electrons. Let's say you were to take a number of snapshots of helium, which has two electrons.

Both of them are in the 1s orbital. It would look like this.

If you took one snapshot, maybe it'll be there, the next snapshot, maybe the electron is there. Then the electron is there. Then the electron is there. Then it's there.

And if you kept doing the snapshots, you would have a bunch of them really close. And then it gets a little bit sparser as you get out, as you get further and further out away from the electron.

But as you see, you're much more likely to find the electron close to the center of the atom than further out.

Although you might have had an observation with the electron sitting all the way out there, or sitting over here.

So it really could have been anywhere, but if you take multiple observations, you'll see what that probability function is describing.

It's saying look, there's a much lower probability of finding the electron out in this little cube of volume space than it is in this little cube of volume space.

And when you see these diagrams that draw this orbital like this. Let's say they draw it like a shell, like a sphere. And I'll try to make it look three-dimensional.

So let's say this is the outside of it, and the nucleus is sitting some place on the inside.

They're just saying -- they just draw a cut-off -- where can I find the electron 90% of the time?

So they're saying, OK, I can find the electron 90% of the time within this circle, if I were to do the cross-section.

But every now and then the electron can show up outside of that, right? Because it's all probabilistic. So this can still happen.

You can still find the electron if this is the orbital we're talking about out here. Right?

And then we, in the last video, we said, OK, the electrons fill up the orbitals from lowest energy state to high energy state. You could imagine it.

If I'm playing Tetris-- well I don't know if Tetris is the thing-- but if I'm stacking cubes, I lay out cubes from low energy, if this is the floor, I put the first cube at the lowest energy state.

And let's say I could put the second cube at a low energy state here. But I only have this much space to work with.

So I have to put the third cube at the next highest energy state.

In this case our energy would be described as potential energy, right? This is just a classical, Newtonian physics example. But that's the same idea with electrons.

Once I have two electrons in this 1s orbital -- so let's say the electron configuration of helium is 1s2-- the third electron I can't put there anymore, because there's only room for two electrons.

The way I think about it is these two electrons are now going to repel the third one I want to add.

So then I have to go to the 2s orbital. And now if I were to plot the 2s orbital on top of this one, it would look something like this, where I have a high probability of finding the electrons in this shell that's essentially around the 1s orbital, right?

So right now, if maybe I'm dealing with lithium right now. So I only have one extra electron. So this one extra electron, that might be where I observed that extra electron.

But every now and then it could show up there, it could show up there, it could show up there, but the high probability is there.

So when you say where is it going to be 90% of the time? It'll be like this shell that's around the center.

Remember, when it's three-dimensional you would kind of cover it up. So it would be this shell.

So that's what they drew here. They do the 1s. It's just a red shell. And then the 2s. The second energy shell is just this blue shell over it.

And you can see it a little bit better in, actually, the higher energy orbits, the higher energy shells, where the seventh s energy shell is this red area. Then you have the blue area, then the red, and the blue.

And so I think you get the idea that each of those are energy shells. So you kind of keep overlaying the s energy orbitals around each other.

But you probably see this other stuff here. And the general principle, remember, is that the electrons fill up the orbital from lowest energy orbital to higher energy orbital.

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

It can fit two electrons. Then the next one that's filled up is 2s. It can fill two more electrons.

And then the next one, and this is where it gets interesting, you fill up the 2p orbital. That's this, right here. 2p orbitals.

And notice the p orbitals have something, p sub z, p sub x, p sub y. What does that mean?

Well, if you look at the p-orbitals, they have these dumbbell shapes. They look a little unnatural, but I think in future videos we'll show you how they're analogous to standing waves.

But if you look at these, there's three ways that you can configure these dumbbells.

One in the z direction, up and down. One in the x direction, left or right.

And then one in the y direction, this way, forward and backwards, right?

And so if you were to draw-- let's say you wanted to draw the p-orbitals. So this is what you fill next.

And actually, you fill one electron here, another electron here, then another electron there.

Then you fill another electron, and we'll talk about spin and things like that in the future. But, there, there, and there. And that's actually called Hund's rule.

Maybe I'll do a whole video on Hund's rule, but that's not relevant to a first-year chemistry lecture.

But it fills in that order, and once again, I want you to have the intuition of what this would look like. Look.

I should put look in quotation marks, because it's very abstract.

But if you wanted to visualize the p orbitals-- let's say we're looking at the electron configuration for, let's say, carbon.

So the electron configuration for carbon, the first two electrons go into, so, 1s1, 1s2. So then it fills-- sorry, you can't see everything.

So it fills the 1s2, so carbon's configuration. It fills 1s1 then 1s2. And this is just the configuration for helium.

And then it goes to the second shell, which is the second period, right? That's why it's called the periodic table. We'll talk about periods and groups in the future.

And then you go here. So this is filling the 2s. We're in the second period right here. That's the second period. One, two. Have to go off, so you can see everything. So it fills these two. So 2s2.

And then it starts filling up the p orbitals. So then it starts filling 1p and then 2p. And we're still on the second shell, so 2s2, 2p2.

So the question is what would this look like if we just wanted to visualize this orbital right here, the p orbitals?

So we have two electrons. So one electron is going to be in a-- Let's say if this is,

I'll try to draw some axes. That's too thin. So if I draw a three-dimensional volume kind of axes.

If I were to make a bunch of observations of, say, one of the electrons in the p orbitals, let's say in the pz dimension, sometimes it might be here, sometimes it might be there, sometimes it might be there.

And then if you keep taking a bunch of observations, you're going to have something that looks like this bell shape, this barbell shape right there.

And then for the other electron that's maybe in the x direction, you make a bunch of observations.

Let me do it in a different, in a noticeably different, color. It will look like this.

You take a bunch of observations, and you say, wow, it's a lot more likely to find that electron in kind of the dumbell, in that dumbbell shape.

But you could find it out there. You could find it there. You could find it there. This is just a much higher probability of finding it in here than out here.

And that's the best way I can think of to visualize it. Now what we were doing here, this is called an electron configuration.

And the way to do it-- and there's multiple ways that are taught in chemistry class, but the way I like to do it-- is you take the periodic table and you say, these groups, and when I say groups I mean the columns, these are going to fill the s subshell or the s orbitals.

You can just write s up here, just right there. These over here are going to fill the p orbitals. Actually, let me take helium out of the picture. The p orbitals.

Let me just do that. Let me take helium out of the picture. These take the p orbitals.

And actually, for the sake of figuring out these, you should take helium and throw it right over there. Right?

The periodic table is just a way to organize things so it makes sense, but in terms of trying to figure out orbitals, you could take helium. Let me do that.

The magic of computers. Cut it out, and then let me paste it right over there. Right?

And now you see that helium, you get 1s and then you get 2s, so helium's configuration is-- Sorry, you get 1s1, then 1s2. We're in the first energy shell. Right?

So the configuration of hydrogen is 1s1. You only have one electron in the s subshell of the first energy shell.

The configuration of helium is 1s2. And then you start filling the second energy shell.

The configuration of lithium is 1s2. That's where the first two electrons go. And then the third one goes into 2s1, right? And then I think you start to see the pattern.

And then when you go to nitrogen you say, OK, it has three in the p sub-orbital.

So you can almost start backwards, right? So we're in period two, right? So this is 2p3.

Let me write that down. So I could write that down first. 2p3. So that's where the last three electrons go into the p orbital.

Then it'll have these two that go into the 2s2 orbital. And then the first two, or the electrons in the lowest energy state, will be 1s2.

So this is the electron configuration, right here, of nitrogen.

And just to make sure you did your configuration right, what you do is you count the number of electrons.

So 2 plus 2 is 4 plus 3 is 7. And we're talking about neutral atoms, so the electrons should equal the number of protons.

The atomic number is the number of protons. So we're good. Seven protons.

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

And if I wanted to figure out the configuration of silicon, right there, what is it?

Well, we're in the third period. One, two, three. That's just the third row. And this is the p-block right here.

So this is the second row in the p-block, right? One, two, three, four, five, six. Right.

We're in the second row of the p-block, so we start off with 3p2. And then we have 3s2.

And then it filled up all of this p-block over here. So it's 2p6. And then here, 2s2.

And then, of course, it filled up at the first shell before it could fill up these other shells. So, 1s2.

So this is the electron configuration for silicon.

And we can confirm that we should have 14 electrons. 2 plus 2 is 4, plus 6 is 10. 10 plus 2 is 12 plus 2 more is 14. So we're good with silicon.

I think I'm running low on time right now, so in the next video we'll start addressing what happens when you go to these elements, or the d-block.

And you can kind of already guess what happens. We're going to start filling up these d orbitals here that have even more bizarre shapes.

And the way I think about this, not to waste too much time, is that as you go further and further out from the nucleus, there's more space in between the lower energy orbitals to fill in more of these bizarro-shaped orbitals.

But these are kind of the balance -- I will talk about standing waves in the future-- but these are kind of a balance between trying to get close to the nucleus and the proton and those positive charges, because the electron charges are attracted to them, while at the same time avoiding the other electron charges, or at least their mass distribution functions.

Anyway, see you in the next video.

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