Hi. It's Paul Andersen and welcome
to chemistry essentials video 5. This is on electron configurations. When I took chemistry
I remember having to do electron configurations. I remember this chart and these different
orbitals. And this diagram. And I learned how to do it. But I never really knew what
I was doing. And so in this video I'm going to show you how to do electron configurations.
We'll do that at the end. But I also really want to tell you what's going on behind the
scenes. And so all electron configuration is is simply the distribution of electrons.
So where the electrons are found in atoms or in ions. And so a good way to figure that
out is to look at their ionization energy which is going to be the amount of energy
it takes to pull an electron away. And so we can quantify that by using Coulomb's Law.
And so what we can do is we can work our way out from the electrons on the inside to those
on the outside. And when we're done we have a pretty good picture of where all of those
electrons are. Because in most atoms, they're going to be multi electron. They have many
electrons. What we'll find is there they're going to be organized into shells, subshells
and then orbitals. And all of those are going to have similar ionization energy. And there's
a chart on ionization energy that's the most important one. Now if we look at these orbitals
and subshells what we'll find is that the ones on the inside are going to be, or the
inner electrons are going to be called core electrons. And the ones on the outside are
going to be called valence electrons. And what's interesting is that those on the inside,
once they're filled will actually shield the valence electrons from the power of the nucleus
itself. And so let's look at the first and simplest of electrons. We've got hydrogen
right here. One proton. One electron. And so Coulomb's Law allows us to quantify the
force between the two. And so it basically comes down to the charges and then the distance
between the charges. And so if we were to look at ionization energy, it's the amount
of energy it takes to pull that electron away. How do we figure it out? You're simply going
to multiply the two charges, so let's say the proton has a positive charge. Let's call
that plus 1. And then electron has a negative charge, let's call that negative 1. And the
only other thing we really need is the radius, the distance between the two. And so let me
ask you a quick question. Which do you think would have a higher ionization energy? Hydrogen.
It's got one proton, one electron. Or helium. It's going to have two of each. And two neutrons
as well. So make a guess. And let me show what helium looks like. Helium, you can see
it just got a little bit smaller and the reason why is since it has two protons in the middle
there's going to be more pull or there's going to be a larger energy of ionization in helium.
And that's because there are a greater amount of positive charge on the inside. So we would
say helium has a higher ionization energy. What about the next one. What if we were to
compare helium then with lithium? What do you think happens to the ionization energy
there? Well you might immediately think it's going to be bigger, but remember the electrons
are organized with only two electrons in that first shell. And then when I learned it was
going to be eight in that next subshell. And so if we look at that lithium, what we'll
find is we've increase the number of protons. So there's more of a charge here. But that
electron is going to be in the next shell. And so since it's in the next shell it's shielded
from the power of that proton, the protons in the nucleus itself. And so it's going to
have a really small ionization energy here just because of the shielding effect and the
fact that we have another shell. And so what do you think about beryllium? Well hold on
to that and we'll take a look at the pattern. And so what this is is the chart that really
allows you to understand what's going on with the electrons. And so let me lay it out for
you. We have atomic number across the bottom. So this would be hydrogen at 1. Helium at
2. All the way out. And then we're going to have ionization energy, which is the amount
of energy it takes to pull that electron away. And what you can see are patterns immediately.
And so these big spikes that you can see are ending with these noble gases, those are going
to be those different shells. You can see that they're broken down into subshells. And
then we're going to get to orbitals in just a second. And so let's kind of lay this out
for just a second. So a shell is going to be, let's look at the first one. So here we've
got hydrogen. So hydrogen is going to be right here. And we said it had a high ionization
energy and that's because it's really close to that nucleus. So it's close to the protons
inside there. And so what we can do is we can fill out this box. So this is going to
be the 1s, we would call that 1 is just going to have a shell and this is going to be the
1s shell. But the reason we have boxes here is that it was later discovered that electrons
will fill what are called orbitals. And those orbitals can only have two electrons in each
one. So let's look at hydrogen. Hydrogen we said has an ionization energy we'll say of
on this chart around 12. If we go to helium it's going to have a higher ionization energy.
Why is that? It's going to be because remember we have those protons on the inside that
are holding it in. But let's watch what happens when we go to lithium. What's happened is
we've gone into this whole next orbital. And so there's this shielding effect of this filled
orbital here. And so we have much less ionization energy there. Let's watch what happens with
beryllium. Well that kind of makes sense. We're increasing the protons. And so that
should go up. But then what happens with the next one. Well, when we go to boron what's
happening is it's actually going into a new orbital. And so it's going into an orbital
of the second subshell. And so if we didn't have these orbitals, what would happen is
it would just be a consistent all the way up to neon. But you can see we kind of have
this jag as we move up. If we go to the next one, that makes sense. Carbon's going to have
more ionization energy. We're adding more protons. Nitrogen same way. But now let's
look at this. Why does the ionization energy go here after we leave nitrogen? Well the
reason why is that these electrons are added. And they're added one at a time in each orbital.
But they'll have a specific spin. And so you can only have two electrons each way. But
what they'll do is they'll fill them kind of like seats in a bus. And so they're all
going to get on one seat per bus, but then when we get more electrons they're going to
have to double up. And so what happens is when we go to oxygen what we're doing is we're
making that electron sit in that orbital. It would rather not be there. There's repulsion
between those electrons, and so it lowers the ionization energy. We go to fluorine,
we go to neon. Once we have that shield effect we jump all the way down to sodium again.
And so once I really understood what was going on in this chart, then all of these subshells
made sense. And so where do they come from? It's quantum mechanics. And it's quantum numbers
that determine that. But if we're looking at periodicity, in other words if we're looking
at the periodic table, these things would be subshells. And so we're going to have a
shell which is going to be the 1, 2, 3, 4, 5. But we're also going to have subshells
which are smaller. And then within each of those we're going to have orbitals if that
makes sense, if we go in size. And so let's get to the electron configuration. And so
what we have here are the s, p, the d and the f. They come from quantum mechanics. What's
interesting about the s is that it only has one orbital in it. The p has three orbitals.
So you could have 6 electrons. d has 5 orbitals so you could have 10 electrons. And f is going
to have 7 orbitals so you could have 14 electrons. So if you want to write this out, you could
just write on a chart. Start with the 1s. And just write all the way down to 7. Then
you want to write right to the write of that 1s or 2s, you're going to write the 2. And
then you're just going to write p's all the way down. And then you go down to the 3. We'll
write a 3 over here and then you're going to write d's all the way down. It could go
all of the way to 7, but we'll never really need that. And then you go to the 4. You write
a 4 over here, and then you write f's all the way down. And so once we have that we
can put diagonal lines in like this and then we can do the electron configuration for anything.
So you wouldn't even need a periodic table. You could just do this and you could figure
out electron configurations. Let's start with hydrogen. What's going to be the electron
configuration of hydrogen? You're simply going to start at the top. And so it's the 1s shell
or orbital or subshell. And so what we would write it as is 1s1. And that 1 represents
the 1 electron that we have inside that orbital. If we go to boron then, how are we going to
write boron? Well you're going to start here. Go here, go here. Boron is going to be a 1s2
because we can only put 2 electrons in there. We then go to the 2s. It's going to be 2s2.
Because we can only put 2 electrons in that. And then we're going to go to the 2p and we
only have 1 electron left. So we're going to put it there. So that would be boron. What
about neon? And again you could pause the video and always try it out. Neon is going
to be this. So it's 1s2 2s2 2p remember could 6. And so we're going to write 6 there. Let's
go to the next one. What about sodium which is number 11. We've got 11 electrons. Well
we could write it out like this. 1s2 2s2 2p6 and then the next one goes in the 3s1. But
you can see how long these electron configurations are going to get. And so we can abbreviate
that. And so we could just say it's neon, put that in brackets with a 3s1. So that allows
us to write the nobel gas in there and then we can add what's after the nobel gas. And
so here's a couple for you to try. Could you try chlorine which as 17 electrons. Or silver
which has 47. I'll put the answers in the video descriptions down below. But give it
a try. So what did we learn? We learned that an electron configuration is the distribution
of what? Again, could you pause the video and figure out what's in the blanks? Distribution
of electrons in atoms or ions. They each have different ionization energy. We could quantify
that through Coulomb's Law. Remember most atoms are going to have multi electrons. And
those are organized in subshells. And then this property is important as well. Those
inner or core electrons are going to shield the valence or outside electrons from the
power of the nucleus. And so what you should have learned is how the energies of electrons
vary within shells of atoms. And so I could point you to this ionization energy chart.
And as we walked through that, that should help. And then the other thing you should
have learned is we could use Coulomb's Law to analyze measured energies of electrons.
And that works great, but it doesn't explain orbitals. And so we have to modify our theory
a little bit. And I hope that was helpful.
