Groups of the periodic table

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- Let's talk a little bit about groups

of the periodic table.

Now, a very simple way
to think about groups

is that they just are the
columns of the periodic table,

and the standard convention
is to number them.

This is the first column,
so that's group one,

second column,

third group, fourth,
fifth, sixth, seventh,

eighth, group nine, group 10,

11, 12, 13, 14

15, 16, 17, and 18.

As some of ya'll might be thinking,

what about these F block
elements over here?

If we were to properly
do the periodic table

we would shift all of these

that everything from
the D block and P block

all are right words

and make room for these F block elements,

but the convention is is
that we don't number them.

But what's interesting?

Why do we go to the trouble about

calling one of these columns,

about calling these columns a group?

This is what's interesting
about the periodic table

is that all of the elements in a column,

for the most part, and
there's tons of exceptions,

but for the most part the
elements in the column

have very, very, very similar properties.

That's because the elements in a column,

or the elements in a group
tend to have the same number of

electrons in their outermost shell.

They tend to have the same
number of valence electrons.

And valence electrons are
electrons in the outermost shell

they tend to coincide, although

there's a slightly different variation.

The valence electrons, these are

the electrons that are going to react,

which tend to be the
outermost shell electrons,

but there are exceptions to that.

There's actually a lot
of interesting exceptions

that happen in the transition
metals in the D block.

But we're not gonna go into those details.

Let's just think a little about

some of the groups that
you will hear about

and why they react in very similar ways.

If we go with group one, group one ...

And hydrogen is a little
bit of a strange character

because hydrogen isn't trying

to get to eight valence electrons.

Hydrogen in that first shell
just wants to get to two

valence electrons like helium has.

Hydrogen is kind of ...

It doesn't share as much in common

with everything else in
group one as you might expect

for, say, all of the things in group two.

Group one, if you put hydrogen aside,

these are referred to
as the alkali metals.

And hydrogen is not
considered an alkali metal.

These right over here are the alkali.

Alkali metals.

Now why do all of these
have very similar reactions?

Why do they have very similar properties?

Well, to think about that

you just have to think about
their electron configurations.

For example, the electron
configuration for lithium

is going to be the same

as the electron configuration of helium,

of helium.

Then you're going to go
to your second shell,

2s1.

It has one valence electron.

It has one electron in
its outermost shell.

What about sodium?

Well, sodium is going to have the same

electron configuration as neon.

Then it's going to go 3s1.

Once again, it has one valence electron,

one electron in its outermost shell.

All of these elements in
orange right over here,

they have one valence electron

and they're trying to
get to the octet rule,

this kind of stable nirvana for atoms.

You could imagine is that
they're very reactive

and when they react they tend to lose

this electron in their outermost shell.

That is the case.

These alkali metals are
very, very reactive.

Actually they have very
similar properties.

They're shiny and soft.

Because they're so reactive
it's hard to find them

where they haven't
reacted with other things.

Let's keep looking at the other groups.

If we move one over to the right

this group two right over here,

these are called the
alkaline earth metals.

Alkaline,

alkaline earth metals.

Once again, they have very similar ...

They have very similar properties

and that's because they
have two valence electrons,

two electrons in their outermost shell.

Also for them, not as quite as reactive

as the alkaline metals.

Let me write this out,
alkaline earth metals.

But for them it's easier
to lose two electrons

than to try to gain six to get to eight.

And so these tend to also
be reasonably reactive

and they react by losing
those two outer electrons.

Now something interesting happens

as you go to the D block.

We studied this when we looked
at electron configurations,

but if you look at the
electron configuration

for say scandium right over here,

the electron, let me do it in magenta,

the electron configuration for scandium,

so scandium,

scandium's electron configuration

is going to be the same as argon.

It's going to be argon.

Then you're going to fill it in

we're in the one, two,
three, fourth period.

It's going to be 4s2.

Then we start filling the D block.

These are the D block elements here.

You have to remember, the D block

you backfill.

In the D block, this is going to

be now 3s1.

How many electrons does it
have in its outermost shell?

Once again its outermost
shell is its fourth shell,

is its fourth shell.

These are, you could argue,
higher energy electrons

that fills this ...

These are filled before that,

and there are exceptions to this

especially that we see
a lot in the D block.

This is what's, I guess you
could say to some degree,

is defining its reactivity.

Although in the transition
metals, the D block elements,

I'm sorry, I made a little mistake there.

This is 4s2 3d1.

Let me emphasize that.

We're backfilling the D block.

But these, their outermost
electrons are in ...

They still have two of
those outermost electrons.

There, once again, are

exceptions in these
transition metals right here

that for the most part

are going in backfilling that D block.

Once you've kind of
backfilled those D blocks

then you come over here

and you start filling the P block.

For example, if you look at

the electron configuration for,

let's say carbon,

carbon is going to have the
same electron configuration

as helium, as helium.

Then you're going to
fill your S block, 2s2,

and then 2p one two.

So 2p2.

How many valence electrons does it have?

Well, in its second shell,
its outermost shell,

it has two plus two.

It has four valence electrons.

That's going to be true for
the things in this group.

And because of that, carbon has similar

bonding behavior to silicone,

to the other things in its group.

We could keep going on,

for example, oxygen and sulfur.

These would both want
to take two electrons

from someone else because they
have six valence electrons

and they want to get to eight.

They have similar bonding behavior.

You go to this yellow
group right over here.

These are the halogens.

There's special name for them.

These are the halogens.

These are highly reactive
because they have seven

valence electrons.

They would love nothing more than to get

one more valence electron.

They love to react.

In fact, they especially love to react

with the alkali metals over here.

Then finally you get to
kind of your atomic nirvana

in the noble gases here.

The noble gases, that's the
other name for the group,

18 elements, noble gases.

They all have the very similar property

of not being reactive.

Why don't they react?

Because they have eight valence electrons.

They have filled their outermost shell.

They don't find the need.

They're noble.

They're kind of above the fray.

They don't find the need to
have to react with anyone else.