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- [Instructor] So let's talk
a little bit about groups

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of the periodic table.

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Now, a very simple way
to think about groups

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

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and standard convention is to number them.

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

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second column, third group,
fourth, fifth, sixth,

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seventh, eighth, group
nine, group 10, 11, 12,

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13, 14, 15, 16, 17, and 18.

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And I know some of
y'all might be thinking,

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what about these f-block
elements over here?

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If we were to properly
do the periodic table,

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we would shift all of these,

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everything from the d-block
and p-block rightwards,

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and make room for these f-block elements,

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

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But what's interesting, why
do we go through the trouble

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about calling one of these columns,

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of calling these columns a group?

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Well, this is what's interesting
about the periodic table,

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is that all of the elements in a column,

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

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but for the most part,
the elements in the column

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have very very very similar properties,

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and that's because the
elements in a column,

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or the elements in a group,

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tend to have the same number of electrons

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in their outermost shell.

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They tend to have the same
number of valence electrons,

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and valence electrons and
electrons in the outermost shell,

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they tend to coincide, although,

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there's a slightly different variation.

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The valence electrons,
these are the electrons

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that are going to react,

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which tend to be the
outermost shell electrons,

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but there are exceptions to that,

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and there's actually a lot
of interesting exceptions

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that happen in the transition
metals, in the D block,

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but we're not gonna go into those details.

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Let's just think a little bit about

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some of the groups that
you will hear about,

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and why they react in very similar ways.

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So if we go with group one,

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group one, and hydrogen is a little bit

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of a strange character,

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because hydrogen isn't trying to get

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to eight valence electrons,

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hydrogen in that first shell

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just wants to get to two valence
electrons, like helium has,

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and so hydrogen is kind of,

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it's not, it doesn't
share as much in common

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with everything else in group one

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as you might expect for, say,

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all of the things in group two.

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Group one, if you put hydrogen aside,

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these are referred to
as the alkali metals,

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and hydrogen is not
considered an alkali metal,

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so these right over here are the alkali,

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alkali metals.

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Now why do all of these
have very similar reactions?

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Why do they have very similar properties?

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Well, to think about that,

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you just have to think about
their electron configurations.

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So, for example, the electron
configuration for lithium

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is going to be the same

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as the electron configuration of helium,

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of helium, and then,

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you're going to go to
your second shell, 2s1.

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It has one valence electron.

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It has one electron in
its outermost shell.

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What about sodium?

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Well, sodium is going to have the same

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electron configuration as neon,

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and then it's going to go 3s1,

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so once again, it has
one valence electron,

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one electron in its outermost shell.

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So all of these elements
in orange right over here,

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they have one valence electron,

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

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this kind of stable nirvana for atoms,

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and so you can imagine is
that they're very reactive,

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and when they react, they tend to lose

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this electron in the outermost
shell, and that is the case.

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These alkali metals
are very very reactive,

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and actually, they have
very similar properties.

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They're shiny and soft, and actually,

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because they're so reactive,

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it's hard to find them where they haven't

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reacted with other things.

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Well, let's keep looking
at the other groups.

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Well, if we move one over to the right,

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this group two right over here,

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these are called the
alkaline earth metals.

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Alkaline, alkaline earth metals.

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And once again, they have
very similar properties,

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and that's because they
have two valence electrons,

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two electrons in their outermost shell,

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and also for them, not quite as reactive

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as the alkali metals,

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but let me write this,
alkaline earth metals,

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but for them it's easier
to lose two electrons

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than to try to gain six to get to eight,

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and so these tend to also
be reasonably reactive,

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and they react by losing
those two outer electrons.

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Now something interesting
happens as you go to the D-block,

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and we studied this when we looked

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at electron configurations,

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but if you look at the
electron configuration

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for say, scandium right over here,

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the electron, let me do it in magenta,

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the electron configuration for scandium,

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so scandium,

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scandium's electron configuration

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is going to be the same as argon,

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it's going to be argon.

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The aufbau principle would tell us

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that the electron configuration,

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we would have the 4s2 just like calcium,

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but by the aufbau principle,

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we would also have one electron in 3d.

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So it would be argon, then 3d1 4s2.

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And to get things in the
right order for our shells,

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let me put the 3d1 before the 4s2.

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And so when people think
about the aufbau principle,

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they imagine all of these d-block elements

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as somehow filling the d-block.

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Now as we know in other videos,
that's not exactly true,

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but when you're conceptualizing
the electron configuration

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it might be useful.

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Then you come over here and
you start filling the p-block.

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So for example, if you look
at the electron configuration

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for, let's say carbon,

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carbon is going to have the
same electron configuration

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as helium, as helium,

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and then you're going to
fill your s-block 2s2,

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and then 2p one 2.

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So 2p2.

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So how many valence
electrons does it have?

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Well, in its second shell,
its outermost shell,

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it has two plus two, it
has four valence electrons,

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and that's going to be true
for the things in this group,

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and because of that,

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carbon has similar bonding
behavior to silicon,

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to the other things in its group.

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And we can keep going on, you know,

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for example, oxygen, oxygen and sulfur,

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these would both want
to take two electrons

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from someone else because they
have six valence electrons,

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they want to get to eight,

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so they have similar bonding behavior.

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You go to this yellow
group right over here,

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these are the halogens.

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So there's a special name for them.

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These are the halogens.

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And these are highly reactive,

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because they have seven valence electrons.

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They would love nothing more

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than to get one more valence electron,

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so they love to react, in fact,

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they especially love to react

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with the alkali metals over here.

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And then finally, you get to
kind of your atomic nirvana

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in the noble gases here.

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And so the noble gases,
that's the other name

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for the group 18 elements, noble gases.

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And they all have the
very similar property

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of not being reactive.

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Why don't they react?

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They have filled their outermost shell.

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They don't find the need, they're noble,

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they're kind of above the fray,

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