Let's figure out the electron configuration for nickel,

right there.

28 electrons.

We just have to figure out what shells

and orbitals they go in.

28 electrons.

So the way we've learned to do it is,

we defined this as the s-block.

And we can just remember that helium actually belongs here

when we talk about orbitals in the s-block.

This is the d-block.

This is the p-block.

And so we could start with the lowest energy electrons.

We could either work forward or work backwards.

If we work forwards, first we fill up

the first two electrons going to 1s2.

So remember we're doing nickel.

So we fill up 1s2 first with two electrons.

Then we go to 2s2.

And remember this little small superscript 2 just means

we're putting two electrons into that subshell

or into that orbital.

Actually, let me do each shell in a different color.

So 2s2.

Then we fill out 2p6.

We fill out all of these, right there.

So 2p6.

Let's see, so far we've filled out 10 electrons.

We've configured 10. You can do it that way.

Now we're on the third shell. The third shell.

So now we go to 3s2.

Remember, we're dealing with nickel,

so we go to 3s2.

Then we fill out in the third shell the p orbital.

So 3p6.

We're in the third period, so that's 3p6, right there.

There's six of them.

And then we go to the fourth shell.

I'll do it in yellow.

So we do 4s2. 4s2.

And now we're in the d-block.

And so we're filling in one, two, three, four,

five, six, seven, eight in this d-block.

So it's going to say d8.

And remember, it's not going to be 4d8.

We're going to go and backfill the third shell.

So it will be 3d8.

So we could write 3d8 here.

So this is the order in which we fill,

from lowest energy state electrons to highest energy state.

But notice the highest energy state electrons,

which are these that we filled in, in the end,

these eight, these went into the third shell.

So when you're filling the d-block,

you take the period that you're in minus one.

So we were in the fourth period in the periodic table,

but we subtracted one, right?

This is 4 minus 1.

So this is the electron configuration for nickel.

And of course if we remember,

if we care about the valence electrons,

which electrons are in the outermost shell,

then you would look at these right here.

These are the electrons that will react,

although these are in a higher energy state.

And these react because they're the furthest.

Or at least, the way I visualize them is that

they have a higher probability of being further

from the nucleus than these right here.

Now, another way to figure out

the electron configuration for nickel--

and this is covered in some chemistry classes,

although I like the way we just did it

because you look at the periodic table

and you gain a familiarity with it,

which is important, because then

you'll start having an intuition

for how different elements react with each other

-- is to just say, oK, nickel has 28 electrons,

if it's neutral. It has 28 electrons,

because that's the same number of protons, which is the atomic number.

Remember, 28 just tells you how many protons there are.

This is the number of protons.

We're assuming it's neutral.

So it has the same number of electrons.

That's not always going to be the case.

But when you do these electron configurations,

that tends to be the case.

So if we say nickel has 28, has an atomic number of 28,

so it's electron configuration we can do it this way, too.

We can write the energy shells.

So one, two, three, four.

And then on the top we write s, p, d.

Well we're not going to get to f.

But you could write f and g and h and keep going.

What's going to happen is you're going to fill this one first,

then you're going to fill this one, then that one,

then this one, then this one.

Let me actually draw it.

So what you do is, these are the shells that exist, period.

These are the shells that exist, in green.

What I'm drawing now isn't the order that you fill them.

This is just, they exist. So there is a 3d subshell.

There's not a 3f subshell.

There is a 4f subshell.

Let me draw a line here,

just so it becomes a little bit neater.

And the way you fill them is you make these diagonals.

So first you fill this s shell like that,

then you fill this one like that.

Then you do this diagonal down like that.

Then you do this diagonal down like that.

And then this diagonal down like that.

And you just have to know that there's only two can fit in s,

six in p, in this case, 10 in d.

And we can worry about f in the future,

but if you look at the f-block on a periodic table,

you know how many there are in f.

So you fill it like that.

So first you just say, OK.

For nickel, 28 electrons.

So first I fill this one out.

So that's 1s2. 1s2.

Then I go, there's no 1p, so then I go to 2s2.

Let me do this in a different color.

So then I go right here, 2s2.

That's that right there.

Then I go up to this diagonal, and I come back down.

And then there's 2p6.

And you have to keep track of how many electrons

you're dealing with, in this case.

So we're up to 10 now.

So we used that one up.

Then the arrow tells us to go down here,

so now we do the third energy shell.

So 3s2.

And then where do we go next?

3s2.

Then we follow the arrow.

We start there, there's nothing there,

there's something here.

So we go to 3p6.

And then the next thing we fill out is 4s2.

So then we go to 4s2.

And then what's the very next thing we fill out?

We have to go back to the top.

We come here and then we fill out 3d.

And then how many electrons do we have left to fill out?

So we're going to be in 3d. So 3d.

And how many have we used so far?

2 plus 2 is 4.

4 plus 6 is 10.

10 plus 2 is 12.

18.

20.

We've used 20, so we have 8 more electrons to configure.

And the 3d subshell can fit the 8 we need,

so we have 3d8.

And there you go, you've got the exact same answer that

we had when we used the first method.

Now I like the first method

because you're looking at the periodic table the whole time,

so you kind of understand an intuition

of where all the elements are.

And you also don't have to keep remembering,

OK, how many have I used up as I filled the shells? Right?

Here you have to say, i used two, then I used two more.

And you have to draw this kind of elaborate diagram.

Here you can just use the periodic table.

And the important thing is you can work backwards.

Here there's no way of just eyeballing this and saying,

OK, our most energetic electrons are going to be

3d8,

and our highest energy shell is going to be 4s2.

There's no way you could get that out of this

without going through this fairly involved process.

But when do you use this method, you can immediately say,

OK, if I'm worried about element Zr, right here.

If I'm worried about element Zr.

I could go through the whole exercise of

filling out the entire electron configuration.

But usually the highest shell, or the highest energy electrons,

are the ones that matter the most.

So you immediately say, OK,

I'm filling in 2 d there,

but remember, d, you go one period below.

So this is 4d2.

Right? Because the period is five.

So you say, 4d2. 4d2.

And then, before that, you filled out the 5s2 electrons.

The 5s2 electrons. And then you could keep going backwards.

And you filled out the 4p6. 4p6.

And then, before you filled out the 4p6.

then you had 10 in the d here.

But what is that?

It's in the fourth period,

but d you subtract one from it, so this is 3d10.

So 3d10.

And then you had 4s2.

This is getting messy. Let me just write that.

So you have 4d2.

That's those two there.

Then you have 5s2. 5s2.

Then we had 4p6.

That's over here.

Then we had 3d10.

Remember, 4 minus 1, so 3d10.

And then you had 4s2.

And you just keep going backwards like that.

But what's nice about going backwards is you immediately know,

OK, what electrons are in my highest energy shell?

Well I have this five as the highest energy shell I'm at.

And these two that I filled right there, those are

actually the electrons in the highest energy shell.

They're not the highest energy electrons. These are.

But these are kind of the ones that have

the highest probability of being furthest away from the nucleus.

So these are the ones that are going to react.

And these are the ones that matter

for most chemistry purposes.

And just a little touchpoint here,

and this isn't covered a lot,

but we like to think that electrons are filling these buckets,

and they stay in these buckets.

But once you fill up an atom with electrons,

they're not just staying in this nice, well-behaved way.

They're all jumping between orbitals,

and mixing together,

and doing all sorts of crazy, unpredictable things.

But this method is what allows us to at least get a sense

of what's happening in the electron.

For most purposes, they do tend to

react or behave in ways

that these orbitals kind of stay to themselves.

But anyway, the main point of here is really just to

teach you how to do electron configurations,

because that's really useful for later on

knowing how things will interact.

And what's especially useful is to know

what electrons are in the outermost shell,

or what are the valence electrons.