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.