The Electron Configuration: Study with Video Lessons, Practice Problems & Examples

Now, before we can talk about electron configurations, realize first that Mendeleev organized the elements in the periodic table through periodic law. In periodic law says that when arranged by increasing atomic weight, elements in the same group share similar chemical properties.

These properties would influence their electron arrangements and by extension, their configurations. It would affect their trends and reactivity and their atomic structures.

So remember, in order to 1st understand electronic configurations, we need to go into periodic law 1st and it helps us to understand how the periodic table itself is organized.

As we stated earlier, the periodic law influences the electron arrangements of the elements and the electron orbital diagrams are the visual representation of electrons within orbitals. Now we're going to say here we have what are called degenerate orbitals. These are electrons in the same set of orbitals having same energy and they're filled using Hans rule. Now Hans rule says that these degenerate orbitals are first half filled before being totally filled.

So if we take a look here, we have our sublevel or subshell can hold a maximum of two electrons. It has one orbital. Within that orbital we have two electrons. One spins up, one spins down. So that would mean that the sublevel has a maximum of two electrons. For P sublevel we have 3 orbitals. Following Hun's rule we would half fill them first. So we go up, up, up. Each orbital we know can hold a maximum of two electrons, so we come back around, down, down, down. So the P sublevel holds a maximum of 6 electrons.

For D we have 5 orbitals. Here. Hans Rule says we have filled them first. Since they're all D set of orbitals, they all have similar energy. So then we come back around down, down, down, down, down for a total of 10 electrons. And then finally the F sublevel has seven of these orbitals. Half fill them again according to Hun's rule. So only half of them, according to Han's rule, come back around to totally fill them in. When we do that, we get a total of 14 electrons.

So just remember, periodic law influences the electron arrangement of elements, and it's these orbital diagrams that depict the visual representation of electron within any given orbital based on sub shell level or subshell letter. So just keep that in mind. S can hold a maximum of two electrons, P can hold up to 6D, can hold up to 10, and F can hold up to 14.

Here it says we need to properly fill in the orbitals of an atom that possesses 8 electrons within its D set of orbitals. Now we know that these are D set of orbitals. Because there's five of them. We have to fill in eight electrons.

Remember, since they're all the same set of orbital types, being D, they all have the same energy and therefore are degenerate. We would half fill them based on Hun's rule. So we go up, up, up, up, up. For our first five we need to fill in eight, so come back, around, down, down, down.

So we stop there because again we only have 8 electrons to fill in. So here we don't need to completely fill in the last two, they will be left just one electron in each. So this would be the way to properly fill in these D set of orbitals that have 8 total electrons.

When writing the electron configuration of an element, it's important that it represents the distribution of electrons from 1S to 2S to 2P and so on within orbitals using the aufbau principle. Now what is the aufbau principle? It says starting with 1S we're going to have electrons filling lower energy orbitals before moving on to higher energy orbitals.

We begin with 1S when we're doing the full ground state electron configuration of an element or an ion. To help us with this, you could take the aufbau diagram approach. We start out with 1S, then we move on, loop back around to 2S then we loop back around to 2P, then to 3S. Then we loop back around to 3P and then 4S. Then continuously loop back around to 3D to 4P to 5S.

Now this is our aufbau diagram. In the aufbau diagram we have here 1S all the way down to 8S. Then we have 2P to three on to seven P then we have 3D to 6D and then we have 4F and 5F. Another way we can look at determining the electron configuration has to do more with the periodic table.

So if you click on the next video, let's reimagine what the periodic table will look like when dealing with the electron configuration of elements and ions.

So here, if we reimagine the periodic table, we'll see it in this depiction. Now realize here we have blue, we have yellow, we have purple, and we have red sectors. Now each of these is called a block. The S block is what's in blue. This is the P block, the D block, and the F block.

Remember, the S sublevel has one orbital and that one orbital can hold a maximum of two electrons, which is why the S block is these two columns to represent the two maximum electrons that the S sublevel can hold. The P sublevel has three orbitals, right? And each one again can hold a maximum of two electrons. So P theoretically can hold a maximum of 6 electrons. That's why the P block has 12345 and six slots for it.

D can have up to 10 electrons because it has five orbitals. And if you count you'll see in here there's ten spots. And then for the F sublevels they can hold U to 14 electrons. And if you were to count these rows in red, you'd see they come out to 14. Now realize here that in this periodic table, the first slide here, which represents hydrogen, starts off our electron configuration as 1S1. Then as we move to the next one we add another electrons. So Helium is 1S2.

Then when we get to the second row, since we're in the 2nd row, we now have two. This is still the block, so this is 1S22S1 and it continues onward and onward. Over here we'd go 1S22S22P1. So following this pattern this would be 3S44S55S66S7 and then here this would be 3P44P55P66P7. When we go to the D block, realize that it's going to drop down by 1, so here this is 4S3. But then when we go to D block it drops down by 1 number so now it's 3D4. This would be 4D55D6.

Now notice how this number goes from 57 to 72. That's because 58 to 71 are here. Remember this? This red line here says that this entire red row exists in between 57 and 72. And then we go between 89 and 104. Because 90 to 103 is down here. These are your F blocks. They also drop down by 1 number, so this is 4F5 and 5F6. So basically, if you can reimagine the periodic table in this fashion, you can use it to figure out the electron configuration of any element or ion given to you. So we're going to put this periodic table to use in order to do future electron configurations.

To help us with this example question, I decided to leave this periodic table up so we can see how to best use it. Here it says write the ground state electron configuration for the following element. Here we're dealing with fluorine and they tell us that it has an atomic number of nine, which means it has nine electrons. On this periodic table, we find fluorine right here.

Ground state means that we're going to start out with 1S orbital and work our way up to fluorine. So we're going to count to fluorine. So we'd say 1S2 1S2 2S2 because of 1/2 slots and then we have to count to F so that would be 2P5. We're in the 2P row. And how many slots do we have to count to get to fluorine? 1 2 3 4 5 2P5.

So here this would be the ground state electron configuration of the fluorine atom. And realize here that these are the number of electrons. So when you add them up 2 + 2 + 5 , that gives me 9 electrons , which is related to the atomic number here of fluorine . Remember, when an element is neutral, its atomic number tells us both the number of protons and the number of electrons.

So again, rely on this depiction of the periodic table to help guide you to the right electron configuration of any element or ion given to you.