- Shells, subshells, and orbitals
- Introduction to electron configurations
- Noble gas configuration
- Electron configurations for the first period
- Electron configurations for the second period
- Electron configurations for the third and fourth periods
- Electron configurations of the 3d transition metals
- Electron configurations
- Paramagnetism and diamagnetism
- The Aufbau principle
- Valence electrons
- Valence electrons and ionic compounds
- Valence electrons and ionic compounds
- Atomic structure and electron configuration
- Introduction to photoelectron spectroscopy
- Photoelectron spectroscopy
- Photoelectron spectroscopy
The Aufbau principle
The Aufbau principle states that electrons fill lower-energy atomic orbitals before filling higher-energy ones (Aufbau is German for "building-up"). By following this rule, we can predict the electron configurations for atoms or ions. The Aufbau principle is most useful for the first 20 elements: from Sc on, the Aufbau principle does not accurately predict the order of electron filling in atoms. Created by Sal Khan.
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- Why does the 4s orbital fill before the 3d orbital ( in general)?(8 votes)
- The Aufbau principle states that electrons are allocated to atomic orbitals in order of increasing energetic content. To calculate the order of energy we use the Madelung rule:
a) In H atoms, all subshells of same n have the same energy.
b) in atoms of 2+ electrons, subshells increase in energy as the value of n+l increases, and for subshells of same n+l value, subshells with a lower n are lower in energy.
In this example, 4s orbitals will have an energy of n+l=4 and 3d orbitals have an energy of n+l=5, that is why 4s orbitals fill before 3d orbitals.
Instead of calculating everything, you can refer yourself to the Madelung diagram, that will give you the full order of orbitals filling. There are only a few expectations to these rules, with more complex reasons.(14 votes)
I just read somewhere that in the electron configuration of Cu the 4s subshell is filled after the 3d subshell, what can be the reason for that?(7 votes)
- We observe that 3d4 and 3d9 do not exist as electronic configurations due to instability issues. In general any d4 and d9 configuration will be changed to d5 and d10 respectively by an electron from the next s subshell being excited to the d subshell . This is what happens with Chromium as well.
Chromium's electronic configuration is supposed to be :
1s2 2s2 2p6 3s2 3p6 3d4 4s2 . But d5 is a situation when the entire d subshell is half filled and this is a more stable scenario than d4. So an electron get's excited from the 4s subshell to the 3d subshell to achieve the half filled subshell stability. Finally, this is the actual electronic configuration of Chromium :
1s2 2s2 2p6 3s2 3p6 3d5 4s1 .
So as you can see this special scenario repeats with Copper when a 3d9 gets changed to 3d10 .
So instead of
1s2 2s2 2p6 3s2 3p6 3d9 4s2
1s2 2s2 2p6 3s2 3p6 3d10 4s1(12 votes)
- Why is the electron configuration of Calcium [Ar]4s2 and not [Ar] 3d2? Shouldn't the electrons in Calcium first fill energy level 3?(4 votes)
- Electrons fill into atomic orbitals according to the aufbau principle, from the lowest energies to the highest energies. When comparing the energies of the 4s and 3d orbitals, the 4s is lower in energy than the 3d. Even though the 3d is in a lower shell number, it still possesses slightly more energy so it is filled after the 4s. Hope that helps.(10 votes)
- Does it matter the order the notation is written in? At5:173d is written before 4s, is it ok to switch the order of those?(3 votes)
- There is some debate which way is better, either 4s3d or 3d4s, for describing the electron configurations of transition metals. Short answer, technically the more correct way would be to write 3d4s, but it is still acceptable to write 4s3d.
The difference arises due to the energy levels of the two subshells. When writing electron configurations the aufbau principal says that electron subshells with lower energy should be filled first. So this would mean writing the lower energy subshells further to the left and the higher ones toward the right. So the two ways of writing the electron configurations are having a disagreement about which subshell is lower in energy.
The problem is when we start filling them for transition metals. So when the 4s and 3d subshells are unoccupied by electrons the 4s is lower than the 3d subshell so we would fill the 4s first and write it first giving us 4s3d. However once filled, the 3d subshell actually becomes less energetic and has lower energy than the filled 4s orbital. Now when filled the 3d subshell would be written first since it has lower energy giving us 3d4s.
So technically the electron configurations for elements like scandium or titanium should be written like 3d4s since they are filling those subshells causing the new energy difference as compared to the subshells being unoccupied. However I find it easier to write them as 4s3d since it follows the order of the pyramid Sal has in the upper left, so I use that format instead. The important part is just to be aware of the different energy levels of the subshells when they're unoccupied versus occupied.
Hope that helps.(10 votes)
- Why does argon's electron configuration have [Ne] but neon's electron configuration not have [He]?(3 votes)
- Anshul, the reason why Neon's electron configuration does not have
[He]is most likely because its configuration is so short that it can be easily written and read without using the shorthand version.
Once you get to elements further down in the periodic table, that's when using the Noble gas notation as a shorthand makes more sense (like with Argon - as Sal does in the video).
You could still write the
[He]shorthand for Neon =>
Ne: [He] 2s2 2p6( if you wanted to show only its valence electrons for example ), and it would be correct as well, but it's not necessary.
I hope this answered your question. :-)(2 votes)
- The about section of this video says that the Aufbau principle only works for the first 20 elements, how are we supposed to configure the rest of the periodic table??(2 votes)
- It says that, "The Aufbau principle is most useful for the first 20 elements". So it can still be used, but it becomes more complicated than simply filling subshells in order. And Sal explains how you can still use the aufbau principle and the accompanying diagram beginning at3:30to work out the electron configurations for elements, even those beyond calcium.
Hope that helps.(4 votes)
- Hi! So only eight electrons are needed in every energy level to make it stable, but these energy levels are able to hold more electrons? Thanks!(2 votes)
- So the eight electron idea you're referring to is known as the octet rule which is a general rule that says certain elements like having a valence shell with eight electrons to be stable. The valence shell being the shell with electrons which participate in chemical reactions and is often the outermost electron shell. The octet rule applies best to second period elements (second row on the periodic table). The second electron shell holds a maximum of eight electrons and its filling creates a stability which elements like to attain.
The rest of the periodic table is full of exceptions to this octet rule. The simplest being the first period elements (hydrogen and helium). They only really make use of the first electron shell which holds a maximum of two electrons. So their valence shell is considered full and stable with just two electrons and so follow what is known as the duet rule.
The third period elements follow the octet for the most part, but you'll find plenty of examples of elements in that row exceeding eight valence electrons. Chlorine in the chlorate ion for example has twelve valence electrons. Elements in the second period don't exceed that eight valance electron limit because their valence shell only holds a maximum of eight electrons. However elements in the third period have an electron shell which can potentially accommodate eighteen electrons. So third period elements would like to follow the octet rule, but they're not limited by it like the second period elements are.
Hope that helps.(3 votes)
- Does it matter how your write the configuration?
For example: Sc:[Ag]3d1 4s2 ; can it also be written as Sc: [Ag]4s2 3d1 ?(0 votes)
- Yes, it actually matters, the reason being that in the case of Scandium, the 4s2 is written afterwards as those electron as Scandium has three valence electrons. The aufbau principle tells us they should add to the lowest available energy sublevel. Therefore, the first electron goes into 3d.
However, it turns out to be more favorable for the second to enter 4s. This is because when there are 20 protons or more, the 3d orbitals are more compact than 4s. The higher repulsion energy of a second electron entering 3d makes it unfavorable compared with the second electron entering 4s. Likewise, the third electron goes into 4s because it is the lowest cost energy option.(6 votes)
- Why does the electron configuration for Bismuth have an F group in it?(2 votes)
- Because it has 4f electrons. Note that the actinides are part of the 6th row of the periodic table.(2 votes)
- Why is it for Ne: 2p^6? - after 1s^2, 2s^2 -- Does this mean that the numbers for 2p can be any number like 2p^8 or 2p^10? I know that all those valence electrons (the ^#) add up to the element number for atomic number but I just don't understand. Does there need to be a specific number then for every shell or can it be any number that comes after ^__ part? Like why 6 specifically? Couldn't you have done 2p^2 and then 3s4? Or something like that? (along with 1s^2, 2s^2)
Sorry if my question is unclear. I hope you can answer it.(1 vote)
- So I take your question is why does the 2p subshell hold specifically 6 electrons. To answer that we have to be family with the quantum numbers of electrons. These are the numbers we can assign to individual electrons in orbit around a nucleus to distinguish them from every other electron. They include the principal quantum number, the angular quantum number, the magnetic quantum number, and the spin quantum number.
For this question we need to focus on the angular quantum number and magnetic quantum number. The angular quantum is usually represented by l and can be integer values starting from 0 (0,1,2,3,etc.). The numbers correspond to the subshell letters. So l = 0 is s, l = 1 is p, etc. The magnetic quantum number, represented by ml, correspond to the number of orbitals in a subshell. The magnetic quantum numbers which are allowed in a subshell are integer numbers that range from -l to +l. So for a p subshell where l = 1, ml could be -1,0, and 1. This means that a p subshell has three orbitals represented by ml values of -1,0, and 1. And we know that an individual orbital holds a maximum of two electrons so if there are three orbitals in a p subshell, a p subshell holds a maximum of six electrons (3x2=6).
So the quantum numbers of subshells show why a p subshell, doesn't have to be just the 2p subshell, holds a maximum of six electrons. That only explains why six in particular is the number, but for a deeper explanation as to why six only fit in a p subshell you would need to delve into quantum mechanics which is a little too much here.
Hope that helps.(2 votes)
- [Instructor] In other videos we introduced ourselves to the idea of orbitals and these are various orbitals in their various subshells that you could find in various shells of an atom. And in this video we're gonna get a little bit more practice with electron configuration. In particular, we're going to expose ourselves to the idea of the Aufbau principle. Now Aufbau comes from German. It means the building principle. It's a very useful way of thinking about electron configurations past calcium. Let's just a get little bit warmed up. What is the electron configuration of neon? Pause this video and think about it and as a hint I will give you the periodic table here. All right, well neon has an atomic number of 10 and if we're talking about a neutral neon atom it's gonna have 10 electrons. And so, the first two will that fill that first shell. So we have 1s2 and then, the next two are going to fill the 2s subshell in your second shell. So then you're gonna have 2s2. And then we have six more electrons to get to 10 and that's now going to fill your 2p subshell, so 2p6. And so what's the order of the subshells that we just filled? Well, first we filled 1s, then we filled 2s, then we filled 2p and you can also see that in the periodic table of elements. In this first row, you're filling that first shell. In this second row or this second period, you are filling that second shell. Now what's going to happen if we were to go to say, argon? So if you're going to go to argon, what will that electron configuration look like? Pause the video and think about that. Well, we can use the noble gas configuration or the noble gas notation. We could say, all right we're going to be building off of neon. So we're gonna have the electron configuration of neon, but then we're going to add electrons into our third shell. So from neon we would then add two electrons into the 3s subshell, 3s2. And then, to get to 18 electrons, we're at 12 right now, we're gonna have six more that are going to be in the 3p subshell, so 3p6. So, on this diagram over here we went from 2s to 2p to fill up neon and then as we went to argon we go to 3s to 3p. Now what would be the electron configuration of calcium? Pause the video and think about it. All right, well calcium has 20 protons. So a neutral calcium would have 20 electrons. So two more electrons than argon. So we can build off of argon and where are those electrons going to go? And this is where the Aufbau principle is interesting. There is indeed a 3d subshell, but in the case of calcium instead of those two electrons being in the 3d subshell, they end up in the 4s subshell. So calcium's electron configuration is the same as argon and instead of it being 3d2 here on top of that it goes straight to 4s2. And so that's why I was drawing this diagram like this and you'll often see that in an introductory chemistry class. You fill 1s first, no surprises. You're filling in that first shell. Then you fill 2s. Then you fill 2p and you filled your second shell. Then you go to 3s, once again no surprises. Then you go to 3p. Now this is the surprise and why this Aufbau diagram is useful. For electron configuration purposes, if you're thinking about potassium or calcium, the extra electrons are now going to go in the 4s subshell. So now let's think about what the electron configuration of the scandium would be. Pause this video and think about that. Well scandium has one more proton than calcium. It has 21 protons and if it is neutral, it's also gonna have one more electron relative to a neutral calcium atom. And so, it could have a similar electron configuration. So we could base it off of argon. We have two electrons in the 4s subshell, so I'll write 4s2 and the Aufbau principle would describe that and the Aufbau principle, this little diagram, would say, all right that other electron is going to be in the 3d subshell, so you do 3d1. And this is indeed an accurate electron configuration for scandium. Now if the Aufbau principle makes you think that you're filling 4s first and then you are starting to fill 3d, if you were actually building up a scandium atom, and that's actually taught in most chemistry books and in most classes, but actually if people start with the scandium nucleus that has 18 electrons. So that would have a positive charge, when they add that first electron it actually does not go to 4s. It goes to 3d. So this electron actually gets added first. If you're actually thinking about building. But I don't want to confuse you too much. In this video we're just thinking about the electron configuration and for that Aufbau can be very useful. Now for electron configuration purposes, 3d, you then go to 4p, and then you then go to 5s, and that's why you might see this type of a diagram, once again, in your traditional first year chemistry books. So the big takeaway here is the Aufbau principle that you'll learn, this type of diagram, it's useful for electron configuration and it might be useful to think about it as you're building these atoms electron by electron, but if you really want the precise accurate truth once you get beyond calcium it gets a little bit more complicated. Now one other thing that I want you to appreciate based on what we just learned is patterns in the periodic table of elements. So for which elements are we building out our S subshell? Well, you could see that for all of these elements right over here, these first two columns, we're building out our S subshell. Now it looks like something is missing there. Is there something else that builds out our S subshell? Well, from that point of view we could actually think of helium as being right over here 'cause helium, we're building out that 1s subshell and because of that all of these elements right over here, we say that they are in the S block. Now, which elements are building out their P subshells? Well, all of these elements right over here are building out their P subshells or have it fully built out. And because of that, all of these elements, we call these the P block. And these elements in the middle right over here, scandium is one of them, they are called the D block. Now one reason why folks might have called it the D block is if you really imagine the Aufbau principle as building up atoms, it might be tempting to say, oh well we're building in the fourth row here, we're building the 3d subshell or in the fifth row here we're building the 4d subshell. Now we now know that that actually isn't true, but from electron configuration point of view it can appear that way and so that's why it is called the D block and I will leave you there.