Groups of the periodic table (video) | Khan Academy

Video transcript

– [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.