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  • - [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.

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

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週期表中的各組|週期表|化學|可汗學院 (Groups of the periodic table | Periodic table | Chemistry | Khan Academy)

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    林宜悉 發佈於 2021 年 01 月 14 日
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