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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.

- Let's talk a little bit about groups

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