Groups of the periodic table | Periodic table | Chemistry | Khan Academy

So let's talk a little bit about groups of the periodic table.

Now, in a very simple way to think about groups is that they just are the columns of the periodic table, and a 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, nine, ten, eleven, twelve, thirteen, fourteen, fifteen, sixteen, seventeen, and eighteen.

I know some of you 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, all right words—and make room for these F block elements. But the convention is that we don't number them.

What's interesting is: why do we go through the trouble of calling one of these columns a group? Well, this is what's interesting about the periodic table: 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 and electrons in the outermost shell tend to coincide, although there's a slightly different variation in the valence electrons. These are the reactive electrons that are going to react while, which tend to be the outermost shell electrons. But there are exceptions to that, and actually, there are a lot of interesting exceptions that happen in the transition metals in the D-block. But we're not going to 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.

If we go to Group 1, 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. So, 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 2.

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 metals. Now, why do all of these have very similar reactions? Why do they have very similar properties?

Well, to think about that, you 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. 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. All of these elements in orange, right over here, have one valence electron, and they're trying to get to the octet rule—this kind of stable nirvana for atoms. You could imagine 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, and actually, they have very similar properties. They're shiny and soft, and 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. If we move one over to the right, this group, too, right over here, these are called the alkaline earth metals. Alkaline earth metals—once again, they have very similar properties, and that's because they have two valence electrons—two electrons in their outermost shell.

Also, for them, not quite as reactive as the alkali metals, but let me write this: alkaline earth metals. For them, it's easier to lose two electrons than to try to gain six to get to eight. 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. 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 is going to be the same as argon.

It's going to be argon; the Aufbau principle would tell us that the electron configuration would have the 4s—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.

When people think about the Aufbau principle, they imagine all of these D-block elements as somehow filling the D-block. 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.

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. Then you're going to fill your S-block: 2s2 and then 2p2. So how many valence electrons does it have? Well, in its second shell, its outermost shell has two plus two; it has four valence electrons, and that's going to be true for the things in this group.

Because of that, carbon has similar bonding behavior to silicon, and we can keep going on. For example, oxygen and sulfur 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. 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.