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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.