very extensive thank you
I think we're now ready to start talking about
some general trends in the periodic table.
And the first one that normally comes up
in a lot of first-year chemistry classes
and it even shows up
on some chemistry standardized tests
is the notion of ionization energy.
And what this is, is essentially the energy required
to remove an electron
from the neutral version of that atom.
So let's see.
Energy to remove an electron.
And just in case you haven't been exposed to
the idea of what an ion is
I guess this is a good time to explain it.
And then this will make sense.
So an ion is essentially any atom or molecule
and we'll talk about molecules.
Molecules are just combinations or groups of atoms
that are all bonded together in some way.
And we'll talk about bonding
in a few videos from now.
But it's an atom or a molecule
in which the protons don't equal the electrons.
And if the protons don't equal the electrons
then you have some charge.
If you have more protons than electrons
you have a positive charge, right?
So let's say if you have hydrogen.
Let's say you have hydrogen.
Normally it's neutral, right?
It has no charge.
But let's say you were to remove an electron from it.
So now you have hydrogen plus an electron.
Hydrogen only has one electron outside.
So now it just has a proton on the inside
no electrons on the outside.
So now it has a plus one charge.
We could write a minus one there.
And now, what we've done in this,
I guess we could say,
this procedure or this reaction is, we've removed
we've ionized hydrogen.
We've removed an electron.
And in this case, this type of ion that hydrogen is
where your protons are greater than your electrons
and this is just a nice word to know.
This is called a cation.
A cation is an ion with a positive charge.
Now you just as easily could have a situation
where you have a, let's say
I don't know, let me take an element
let's say we start with chlorine.
In a stable form it has seven valence electrons.
Why don't we add an electron to it.
And now this will actually be pretty stable.
It will have eight valence electrons
but it'll have a negative charge.
So this, right here, this is a negative ion.
And that's known as an anion.
The way I remember it is, in a lot of words,
"a" means the opposite or a negation.
The "a" prefix.
So "a", anion, means a negative ion.
And then, obviously, cation is the other.
But ionization energy
it really should maybe be called cationization energy
because it's the energy required
to remove an electron
not remove or add an electron.
So it's really the energy required
to turn something into a cation.
So we've already discussed the periodic table
and we'll get to this chart in a second.
But just based on what we already know
which elements will it be harder
to remove an electron from
and which ones will it be easier?
We already talked about -- let's start with group one,
just because it is group one.
Group one is right here.
And we'll start especially with the alkali metals.
We can put hydrogen aside for now.
But all of these, guys, we've talked about a lot.
In order for them to get to the magic number eight
in their outermost shell
the easiest way for them to do it
is just to get rid of that one valence electron
they have in their outermost shell.
Let's say potassium, right there.
Potassium has one valence electron in its fourth shell.
If it just got rid of it
then it has eight in its third shell.
Then it looks a lot like argon
from an electron configuration point of view,
which would be nice.
So these, guys, really want to give away electrons.
So it requires very little energy to ionize them
or to cationize them, to take away their electron.
So this is low ionization energy.
And I think you see where this is going.
What about these guys?
What about neon?
How hard is it to remove an electron from neon.
Well neon is completely satisfied.
It doesn't want to even deal with any of this
reaction business and bonding business.
It's like, you know what?
I've achieved happiness in life.
Don't mess with my electron.
So it really doesn't want to give away its electrons.
Neon, or krypton, or argon
or any of these noble gases.
So to remove an electron from one of these guys
requires a lot of energy.
So this is a high ionization energy.
So in general
as you go from left to right across the periodic table
it goes from low to high.
And some people memorize this,
but you really don't have to memorize it.
You just say, look,
these guys have one extra electron
that they're always trying to get rid of.
These guys have two.
Ideally they'd want to get rid of, maybe,
both of their electrons.
But the first electron doesn't want to jump off
as much as the first electron here
because the first electron here
you get rid of it
you immediately get to the super stable state.
And that trend just becomes more and more true.
This guy, definitely, under no circumstances
wants to give away an electron.
This guy is so close to being like neon
that he definitely doesn't want to go a step backwards
and look more like oxygen.
So he doesn't want to get rid of an electron.
So the trend is pretty clear.
Anytime you're confused about the trend
just look at the extreme cases.
This guy wants to give away electrons.
This guy wants to get electrons.
So if you say, what's the energy required
to take away an electron?
Well this guy's almost going to give it to you.
While this guy's going to be very hard,
to take away an electron.
Now what happens as you go down?
As you go down in this direction,
let's say you go along a group, right?
We already established that
these alkali metals like to give away electrons.
But as you go down,
the electron cloud gets bigger and bigger.
And you could say this 55th electron
there's 55 protons,
there's also in a neutral cesium atom,
there's also 55 electrons
that the 55th electron is
a lot further away from the nucleus of this atom
than the third electron is in the case of lithium.
So this, the 55th electron,
not only does it want to be given away,
but it has even a weaker attraction to the nucleus
than the third electron does in lithium.
So because you're getting larger and larger
as you go down a group,
and the electrons are getting
further and further away from the nucleus,
this guy wants to give away his electrons
even more than lithium does.
So ionization energy decreases as you go down.
Even though xenon really wants to keep his electrons
he's a little bit more willing to give them away
than neon is.
So in general, the ionization energy,
or the energy required to ionize an atom,
will increase as you go up.
And if you ever forget it, don't memorize these things.
Because that might be useful for just one test
but then later in life
when you're 42 and someone asks you,
hey, what has a higher ionization energy,
cesium or fluorine?
You might have forgotten it.
But then if you look at a periodic table, you'd say,
you know, cesium has this one electron
that it's just dying to give away.
It's super far away from the nucleus.
That 55th electron just really wants to leave.
While fluorine, that ninth electron
just needs one more electron.
Well let's say neon.
Neon is super happy.
All the electrons are in the stable configuration.
They're close to the nucleus.
There's a lot of attraction
with the protons in the nucleus.
They definitely don't want to
give away their electrons.
So if you talk about the
energy required to remove them
very low energy at cesium
very high energy at helium
So that's the trend.
And this is just, this trend we'll see often
this is your willingness to give electrons.
And this is how much you want to
hog electrons or keep them to yourself.
And this right here, this is actually
I got this off Wikipedia
this is the actual, measured ionization energies.
And people looked at patterns like this.
I'm not sure if they actually
looked at the ionization energy,
but they looked at patterns like this,
and this is actually
where they came up with the periodic table.
Because they said, look,
as we increment up
the number of protons that we have in an atom,
and likewise, the number of electrons,
we see these repeated patterns, or these periods,
in the elements.
So this is hydrogen.
So hydrogen's ionization energy
is around 13 1/2 electron volts.
Which is a unit of energy.
It can be converted into joules, if you like.
But then all of a sudden, helium is a lot more stable.
It takes almost double the energy
to remove that second electron from helium
because it's so stable.
But as soon as you do that,
as soon as you go from helium,
this point right here,
that point right here is lithium.
Lithium is atomic number of three.
Let me put that in.
This is lithium.
All of a sudden, to remove an electron from lithium
it only requires five electron volts.
So less than half of what it required for hydrogen.
Then as you go to the right of the periodic table,
the ionization energy keeps increasing.
These little divots are interesting.
We could talk about that, maybe, in a future video.
But the ionization energy increases
all the way to neon.
And then you get to neon
and then you add one more
see, neon is a noble gas, it's completely happy
then you add one more electron
and you get to sodium.
And you say, oh, now sodium,
it's really easy to take away that electron.
And the ionization energy drops.
Not only did it drop,
but here it dropped slightly below lithium.
So you see this general trend.
So these are the noble gases right here.
Let me make sure
you can see them all in the video screen.
Very hard to remove that eighth electron.
It's very hard, but a little easier than helium.
And that's because neon is a little bit bigger.
The electrons are a little bit
further away from the nucleus.
Same pattern, but it's actually interesting.
Argon is actually not that different than hydrogen.
It's hard to remove that electron.
But it's actually no harder
to remove the electron on krypton
than on hydrogen.
And you go all the way out to radon.
And you can see, you know,
why is the distance between these increasing?
Well if you remember the periodic table,
all of a sudden when you go after argon,
then you have all of the d-block elements that go in.
So after you get to argon, notice.
This is where you only had s- and p-block elements.
Now all of a sudden the d-block elements show up.
So you have more subshells to fill up.
And now all of a sudden, before you get to radon,
you're starting to fill out the f-block, too.
And that's why that distance is increasing,
because you have to also fill out the f elements.
But the general trends that we just discussed
they apply here.
That as you go to the right along the periodic table,
it becomes harder and harder to remove an electron
because everyone wants to
get to the magic number eight.
And it's very hard to remove them from a noble gas.
But then as you go one more above that,
this guy would love to give away his electron
and get back to a configuration like neon.
And then the other trend is, as you go down a group
in this case, these are the noble gases
the ionization energy decreases.
And that's because
the overall radius of the atom increases,
so the valence electrons in radon are
a little bit less strongly attracted
than the valence electrons in helium.
Anyway, I think that's it for ionization energy.
I'll continue this in another video.
We'll talk about
metallic character and electronegativity.
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