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Groups of the periodic table

The s-, p-, and d-block elements of the periodic table are arranged into 18 numbered columns, or groups. The elements in each group have the same number of valence electrons. As a result, elements in the same group often display similar properties and reactivity. Created by Sal Khan.

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  • spunky sam blue style avatar for user NB Marina
    why are the Group 2 elements are called alkali "earth" metals, why earth??
    (34 votes)
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  • piceratops ultimate style avatar for user Uzayr Ali
    I don't get it.
    What are the 3s1 and 3s2. Could someone explain all these to me properly?
    (17 votes)
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  • male robot johnny style avatar for user Ahmed Yaqoob
    Why does Sal use 18 groups, we were taught to use 8 groups in our school, is our school wrong or is Sal wrong..?
    (8 votes)
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    • duskpin ultimate style avatar for user Cheyanne Izon
      Sometimes people leave out the transitional metals as it can help to see how many valence electrons are in each group. For example, in group 17 there are 7 valence electrons so if you forget about the "ten" part its just group 7. Or in other words the Noble Gases are group 8 instead of 18.
      (21 votes)
  • duskpin sapling style avatar for user AbbieRose
    i don’t understand all of the 3d4 and 2p2 like what does that mean?
    (9 votes)
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    • piceratops ultimate style avatar for user FrozenPhoenix45
      Okay, this is a long, long explanation. You may already know some of this stuff but I'll go over all of it anyway.

      Basically, every atom has orbitals around their nucleus. An orbital is basically a given space that an electron can exist in. For example, if the orbital is in the shape of a sphere, the electron can exist in any point within that sphere, but it cannot go outside that sphere. Each orbital also represents a different energy level. The farther out the orbital from the nucleus, the more energy the electron has. Finally, each orbital can hold only two electrons and no more.

      Okay, now, there are different types of orbitals that each take different shapes. s orbitals are sphere shaped, and there can only be one at each energy level. p orbitals are kind of dumbbell shaped and there can be up to three of them at each energy level. d orbitals are a really complicated shape and there can be up to five. And I really don't know much about f orbitals (I've never had one that large) but I believe there can be seven

      The numbers that you are talking about are called electron configurations, and each element has its own unique one. Their electron configuration is determined by the number of electrons their neutral form has and the orbitals in which those electrons exist. For example, let's start with hydrogen. Hydrogen has an electron configuration of 1s1. This means that

      1 = the first energy level
      s = s orbital
      1 = 1 electron in that s orbital

      And if you were to go a step farther with helium, you would get 1s2, because it has two electrons in that s orbital.

      Now for a more complicated atom. Let's do oxygen. Okay, so we know oxygen has 8 electrons, and only two electrons can exist in each orbital. So on the first energy level, we have 1s2, because it has at least as many electrons as helium.

      Now, there can be no p orbitals in the first energy level, so then we jump straight to the second. The s orbital there is also filled completely for 2s2. However, we still only have half the electrons. So now for p orbitals. Do you remember me saying that up to three p orbitals can exist in each energy level? Well that means that there can be up to six electrons in each energy level for the p orbitals. so then we would say 2p4

      2 = the energy level we are on
      p = the type of orbital
      4 = the number of electrons needed to complete oxygen's number of 8

      So oxygen's electron configuration is: 1s2 2s2 2p4

      Hopefully you're following me, but the best way to really get the hang of this is to practice it yourself. I'l give you one more. Let's do iron. Iron has an atomic number of 26, so that mean's it'll have 26 electrons in its neutral form.

      1s2 2s2 2p6 3s2 3p6 4s2 3d6

      That is the electron configuration for iron. If you add up all the numbers after the orbital letters, you'll find they equal twenty-six.

      On a side note, yes, I know 3d orbitals comes after 4s. It's confusing, but technically it has more energy, so it comes after

      One last thing. The periodic table is organized into blocks to represent the orbitals. The first two groups (1A and 2A) along with helium are called the s block. If you're counting the electrons one by one, one electron will go into an s orbital for each element in the s block you count. The same is true for the p block (groups 3A - 8A) and the d block (the transition metals--the little dip in the middle of the table). You'll notice that each block has as many elements in a row as the number of electrons it can fit into those orbitals in that energy level.

      Another side note: the group that is floating out of the rest of the table is the f block

      So I think that's it. This is without a doubt the longest answer I've written and might hold that record for a while, but I hope you understand a little better. I encourage you to try to find the electron configurations of other elements on your own, as practice really helps with this. And if you're still confused...





      here are some Khan Academy videos and articles to help you.

      I hope this helped! Comment if you have any questions; I'll answer to the best of my ability.

      (I also urge you to read the first comment to this answer. Another user has made a few minor corrections to my answer that are worth reading)
      (33 votes)
  • male robot hal style avatar for user R3hall
    Sal introduces Group 11 at and I can't help but notice that this is the column that contains copper, silver, and gold. We know their colors: reddish-brown, grey, and yellow. Roentgenium [Rg] aka #111 is also in Group 11. What does that mean and what color is #111?
    (10 votes)
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  • piceratops seedling style avatar for user Abhishek singh
    why hydrogen have two valenced electron?
    (7 votes)
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  • piceratops sapling style avatar for user Moulee Veera
    isnt hydrogen a halogen? Why is it marked as a mettaloid (green)
    (1 vote)
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    • spunky sam blue style avatar for user Alfred
      The term Halogens is from the Greek word for salt forming. For this reason, Hydrogen is not considered a halogen even though it has many of the same properties of the halogens. Hydrogen is often grouped with the 7A elements because it is lacking an electron.
      (2 votes)
  • blobby green style avatar for user arrittmo
    would the alkali medals react to water?
    (5 votes)
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    • blobby green style avatar for user mercy
      Yes, alkali metals react very strongly with water. Because their valence shells only need to lose one electron to become stable, they will easily try to form bonds. If it's hard to visualize, look up videos of people throwing sodium stored in mineral oil into large lakes; its really interesting!
      (6 votes)
  • male robot donald style avatar for user Eshanth Mendez
    What are the colors in the periodic table meant for?
    (4 votes)
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  • blobby green style avatar for user lakshyav500
    I am extremely confused about how Sal gets 2s1 and 3s1. How am I supposed to find the electron configuration in the first place? Can someone please explain?
    (4 votes)
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    • leaf red style avatar for user Richard
      Electrons follow the aufbau principle when filling into orbitals. This means that they are placed into the lowest energy orbitals first and then build upon those to fill higher energy orbitals. To know the exact filling order, we can use the diagonal rule. Essentially we write out the subshells for each electron shell and draw diagonal lines through the list starting at the top; this tells us the order of the orbitals we fill.

      1: 1s
      2: 2s 2p
      3: 3s 3p 3d
      4: 4s 4p 4d 4f
      5: 5s 5p 5d 5f
      6: 6s 6p 6d
      7: 7s 7p

      Using this rule, we expect the 1s orbital to be filled first, then the 2s, then the 2p, then the 3s, then the 3p, then the 4s, then the 3d, then the 4p, and so on. As for the number of electrons for each element, they should equal the number of protons or atomic number of the element. Each subshell also contains a maximum number of electrons. s subshells hold only 2 electrons, p subshells hold 6, d subshells hold 10, and f subshells hold 14.

      Hope that helps.
      (6 votes)

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.