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In the last video we talked about ionization energy, or
the energy required to remove an electron.
And we saw the general trend in the periodic table, that
when you're in the bottom left-hand side close to
cesium, cesium really wants to give up electrons.
It's a big atom.
It only has one extra electron in its sixth shell.
It can just give it up, and then it'll have
five complete shells.
So it really wants to give it away, so it requires very
little energy to ionize.
On the complete other side of the spectrum, helium requires
a lot of energy to ionize.
It's very happy.
I has a full shell at the first shell.
It's a very small atom.
The electrons are very close to the protons.
So the coulomb force is super-duper-duper strong.
So it takes a lot of energy to remove that incremental
electron, and we learned that.
And the one thing I want to cover before moving on to
other types of trends or properties amongst the
different atoms is the idea of a second ionization energy.
And I want to do this because sometimes it's covered on some
chemistry exams or some chemistry standardized tests.
And it's just the idea that ionization energy is the
energy required to remove the first electron, to go from a
neutral state to popping one electron off of it.
The second ionization energy is, then, the energy required
to remove the very next electron.
And the reason why this is interesting, is sometimes
they'll say, OK, what elements have a very high second
And your temptation would be, OK, high ionization energy,
that also probably means high second ionization energy.
And that might be true.
For example, neon has a very high ionization energy, It
really wants to keep that 10th electron, because it fills out
the second shell.
And then, of course, even if you were able to remove that
electron, to remove the ninth electron, when now its
configuration looks a lot like fluorine,
that's still very difficult.
So you'd say its second ionization energy
is still very high.
But if you think about it, the elements with the highest
second ionization energies are going to be some of the
elements with the lowest ionization energy.
So, think about it.
And that might be kind of confusing.
Lithium, for example.
Very low ionization energy.
It's got that extra electron.
It just wants to give it away.
But once it gives it away, it's in a very stable
situation, Then its electron
configuration looks like helium.
So to remove that second electron is
So lithium has a very high second ionization energy.
And so you might run into a question where they're like,
which of these elements has the biggest difference between
their ionization energy and their second ionization
energy, where their second ionization energy is higher
than their ionization energy.
And lithium, or anything in group one, that would be true,
because as soon as you remove one electron, its electron
configuration becomes super stable, so removing that
second one is super-duper difficult.
And you also see this in this chart.
This is, of course, the first ionization energies.
But let's say, the case with lithium, you
removed that electron.
It was very easy.
You only needed five electron volts to do it.
But then your configuration looks a lot like helium.
So that second ionization energy is going to look a lot
like helium's first ionization energy.
Anyway, I don't want to confuse you too much.
But that's an interesting point that might pop up every
now and then.
Now another property, which is in a lot of ways, in my mind,
related is the idea of electronegativity.
The concept came up by Linus Pauling.
I always remember him.
He was a famous chemist. What I always remember is that he
was famous for being convinced that Vitamin C was kind of the
key to living forever.
And he would take huge doses of Vitamin C.
I should probably read up on that again.
I don't want to spread lies about Linus Pauling.
But I remember reading that when I was in high school.
But anyway, he came up with the idea of electronegativity.
And the idea is that when two atoms form covalent bonds--
and I haven't taught you what a covalent bond is, and I was
planning on doing that in a couple of videos from now--
but the idea of a covalent bond is really just atoms
Let me draw that out.
So if I have oxygen, oxygen looks something like this.
I could draw it like that.
I could also draw oxygen like this, just because I'm going
to use these extra electrons to bond.
And if you take oxygen like that and you add it to two
hydrogens-- hydrogen has one electron--
what's going to happen?
You might not know yet, if you haven't seen a covalent bond.
But the atoms will actually share electrons.
So this oxygen, you put it in the center.
You have these, over here.
Let me draw it like that.
The electrons from oxygen I'll do in green.
And then hydrogen, I'll just do it in this orange color.
So we have two of these hydrogens.
So one hydrogen will be there.
And then the other hydrogen will be there.
Now what just happened?
Well if this hydrogen can pretend that both of these
electrons, it has to kind of share this green
one with the oxygen.
And the deal is, hey, I share the green one and you let me
borrow the green one, and I'll let you borrow the orange one,
we both can kind of feel like we have a stable electron
Hydrogen feels good because the one s-shell
is completely filled.
Oxygen fills great because it's valence shell is
completely filled with eight electrons, two
of which are borrowed.
So it feels great.
This is a covalent bond, where the
atoms are sharing electrons.
And so this sometimes will be drawn like this.
Those are the extra electron pairs of oxygen.
And they'll just draw a line like that.
And that line implicitly is saying, look, there's two
atoms on either end.
There's the oxygen electron there.
And then you have the hydrogen electron there.
And they're kind of shared.
These two things mean the same thing.
But that line just means a covalent bond.
Now my whole point behind talking about covalent bonds a
little prematurely is so that I can touch on
And the idea that Linus Pauling came up with is that
in these covalent bonds, the sharing is not equal.
That some of the atoms will hog the
electrons a little more.
So in this case, oxygen.
We learned about oxygen.
Oxygen is way over here.
It loves to grab electrons.
It has a very high ionization energy.
It's only two away from having an electron configuration
similar to neon and being super-duper happy.
So oxygen loves electrons.
Hydrogen is a little bit here or there.
It could gain an electron and then it'll
have a stable 1s orbital.
Or it could lose an electron and it will essentially just
turn into a positive ion.
It can go either way.
So it's a little bit more ambivalent about what happens
relative to the electrons.
But oxygen really wants the electrons so
that it can get completed.
So in this relationship between oxygen and hydrogen,
oxygen is more electronegative.
It's more electronegative, which means it kind of hogs
the electrons a little bit more.
So if you were to draw this relationship here, it might
look something-- If you were to draw this bond.
This is all abstract.
Maybe you would draw it a little bit
heavier on that side.
And this is not really at all a convention, but I
just made that up.
Or if you just drew the hydrogen and the oxygen part
of it, maybe the electrons spend most of their time
around-- this is a probability distribution-- and less of
their time around hydrogen.
And that would be true for the other hydrogen.
They spend less of their time around the hydrogen and a lot
more of their time around the oxygen.
The idea of electronegativity is just that one atom is going
to hog the electrons more when you form a covalent bond.
Now if we wanted to figure out the trend of electronegativity
on the periodic table, what do you think's going to happen?
Which elements are likely to hog electrons?
Well, the ones that love electrons.
The ones that it's very hard to take
electrons away from them.
The ones that are super-close to completing a full eight
valence electrons in their outermost shell.
So the most electronegative atoms are
going to be right here.
They're going to be the halogens, especially the
fluorine, because the small ones want the electrons even
more because they're a small atom.
The electrons are going to get closer to the nucleus.
And the reason why I'm not talking about the noble gases
here is because these don't form covalent bonds.
They're always happy.
They're all these inert gases.
Inert just means that they don't do anything.
A similar word is inertia.
Inertia means the tendency to want to stay at rest, not do
anything, or stay in motion, but I won't go
into that too much.
But these are inert.
They don't do anything.
So these guys react.
They form covalent bonds up here.
And when they form covalent bonds, they hog the atoms [correction: electrons].
Likewise, when these guys down here form covalent bonds,
they're like, you know what, you can have the atoms [correction: electrons]. I
don't need them.
I'm actually happier without them altogether.
In fact, sometimes these guys actually just
give away the atom [correction: electron].
They don't even form a covalent bond.
It's called an ionic bond.
We'll talk about that in the next video.
But as you can see, the trend is the same as it is for
These guys, a lot of energy required
to remove an electron.
That's because they love electrons.
So these guys are also very electronegative.
They're going to hog the electrons in a covalent bond.
These guys, very low ionization energy.
Very easy to take an electron away from them.
And that's why they have very low electronegativity.
They're very unlikely to hog an electron in a bond.
Now the other trend that some people sometimes talk about is
the metallic nature of an element.
And so, there's a lot of things that, in my mind, I
imagine when someone talks about metallic nature, I
imagine it should conduct electricity, it should be
shiny, it should be malleable.
I can bend it without it cracking.
That's how I imagine metallic nature.
But when people talk about it in chemistry, they're really
just talking about a willingness
to give away electrons.
That's metallic nature.
And that is important.
If you talk about something that's going to conduct
electricity or be malleable or have this sea of electrons
available that the atoms can sit in.
But the same trend.
Which atoms are very likely to give away electrons?
Well, the bottom left, right?
As you go down, the atom gets bigger, so the electrons are
further away from the nucleus.
So the coulomb force is weaker, so those electrons are
more weakly bound.
And also, if you just have one extra electron here or two
extra electrons there in your outermost shell, you're just
like, hey, let me get rid of them and then I'll have a
complete outer shell.
So these guys want to give away electrons.
So they have a very high metallic nature.
These guys want to keep electrons.
And they want to take more.
So they have a very low metallic nature.
In fact these are completely non-metallic in any way.
And if you were to say, within a group, the trend-- I mean, I
did the diagonal, but that's in general true-- is that the
further you go down a group, the size of the atom is
increasing and the outer electrons are
further from the nucleus.
So the electron force is going to be weaker-- or the coulomb
force is going to be weaker.
So you're more likely to give away electrons.
So your metallic nature will increase as you go down.
And your metallic nature will increase as you go to the
left, because when you only have a couple of electrons in
your outermost shell, you want to give them away.
So metallic nature, it goes in the opposite direction.
It goes like that.
But for the same reason.
These guys love to hog electrons.
These guys love to give them away.
So ionization energy increased to the top right.
Electronegativity increased to the top right.
Metallic nature increased to the bottom left.
The last trend we could talk about is just atomic radius.
And there's a lot of different ways to actually measure this.
And there's no one best way, because obviously, we already
talked about it, an atom doesn't have a fixed radius.
The electron could show up pretty much anywhere.
So you could just kind of do a hard boundary.
OK, 90% chance of finding the electron.
That's your sphere of the atom.
Or you could say, OK, if this atom bonds with another atom,
what is half the distance between the two nucleuses.
If you make a bond like that.
This is the distance between the two nucleuses and then you
can say the atomic radius is that.
So there's a lot of ways.
But I think you get the general idea.
It's just the size of the atom.
And you could already imagine that as you go down any one
group, the size of the atom increases.
You're adding on more and more energy levels,
more and more shells.
The atom is just getting larger and larger.
In fact, we've used that as an argument as to why, as you go
down, ionization energy goes down, or
electronegativity goes down.
So the atoms become larger as you go down.
Now, the one thing that might be a little un-intuitive is
what happens as you go to the right?
You're adding electrons as you go to the right, but you're
adding them all in the same shell, right?
So if this is the nucleus, right there, and you're in
some shell, some orbital shell.
And obviously, they're not all spheres.
But let's say you're in some orbital shell.
As you go to the right in a period, you just keep adding
electrons to that shell.
This is a super-gross oversimplification.
And as you go to the right, you have more
protons in the nucleus.
So this is only getting more and more positively charged.
So what happens is that these electrons get pulled inwards.
They get pulled inwards.
So as you move to the right on the
periodic table, size decreases.
And then you say, OK, but what about when you
go to the next period?
You're getting more protons there.
Won't that decrease?
But at the same time, you're now adding the electrons in a
new shell that's further from them.
So it gets larger when you go to the new period.
So electron size, as you go down, large.
And as you go to the left, you get larger.
So electron size goes from the bottom right to the top left.
Although in general, the things that are in a lower
period will have a larger size than most things in a higher
period, regardless of what group it's in.
But the general trend within a group, the higher the number,
the larger the atom.
Within a period, the more protons you have,
the smaller the atom.
Anyway, I hope you found those interesting.
In the next few videos we'll start with bonding.
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