# Electron configurationsÂ 2

## Video transcript

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. 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. 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. 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. And how many have
we used so far? 2 plus 2 is 4. 4 plus 6 is 10. 10 plus two 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 2d there, but remember, d, you go
one period below. So this is 4d2. Right? Because the period is five. So you say, 4d2. And then, before that,
you filled out the five s2 electrons. And then you could keep
going backwards. And you filled out the 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. 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.