- [Instructor] So 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 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.
And I know some of y'all 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,
everything from the d-block and p-block rightwards,
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 through the trouble
about calling one of these columns,
of calling these columns a group?
Well, 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,
and 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 and 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,
and 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 bit about
some of the groups that you will hear about,
and why they react in very similar ways.
So 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,
and so hydrogen is kind of,
it's not, 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,
so 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.
So, for example, the electron configuration for lithium
is going to be the same
as the electron configuration of helium,
of helium, and 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,
and then it's going to go 3s1,
so once again, it has one valence electron,
one electron in its outermost shell.
So 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,
and so you can imagine is that they're very reactive,
and when they react, they tend to lose
this electron in the outermost shell, and that is the case.
These alkali metals are very very reactive,
and actually, they have very similar properties.
They're shiny and soft, and actually,
because they're so reactive,
it's hard to find them where they haven't
reacted with other things.
Well, let's keep looking at the other groups.
Well, 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.
And once again, they have very similar properties,
and that's because they have two valence electrons,
two electrons in their outermost shell,
and also for them, not quite as reactive
as the alkali metals,
but let me write this, 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,
and 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.
The aufbau principle would tell us
that the electron configuration,
we would have the 4s2 just like calcium,
but by the aufbau principle,
we would also have one electron in 3d.
So it would be argon, then 3d1 4s2.
And to get things in the right order for our shells,
let me put the 3d1 before the 4s2.
And so when people think about the aufbau principle,
they imagine all of these d-block elements
as somehow filling the d-block.
Now as we know in other videos, that's not exactly true,
but when you're conceptualizing the electron configuration
it might be useful.
Then you come over here and you start filling the p-block.
So 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,
and then you're going to fill your s-block 2s2,
and then 2p one 2.
So 2p2.
So 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,
and that's going to be true for the things in this group,
and because of that,
carbon has similar bonding behavior to silicon,
to the other things in its group.
And we can keep going on, you know,
for example, oxygen, oxygen and sulfur,
these would both want to take two electrons
from someone else because they have six valence electrons,
they want to get to eight,
so they have similar bonding behavior.
You go to this yellow group right over here,
these are the halogens.
So there's a special name for them.
These are the halogens.
And these are highly reactive,
because they have seven valence electrons.
They would love nothing more
than to get one more valence electron,
so they love to react, in fact,
they especially love to react
with the alkali metals over here.
And then finally, you get to kind of your atomic nirvana
in the noble gases here.
And so the noble gases, that's the other name
for the group 18 elements, noble gases.
And they all have the very similar property
of not being reactive.
Why don't they react?
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.