So remember electron configurations are representation of how electrons are distributed into orbital's. And furthermore how those orbital's fit into different energy levels. We can do electron configuration of different elements more particularly the transition metals by either looking at an off about diagram or starting with one S double back to two S, Double back to two p. And so on. Or we could just simply look at the periodic table and realize that it's made up of S P D N F blocks. Remember here are blue ones are s these are P blocks and blocks and blocks. And with our transition metals we can simply look at the periodic table and determine what our electron configuration will be. So we've done this before in the past now we're looking at it more towards the transition metals. Now remember that there are exceptions that exist with the electron configuration of transition metals And that those are the ones that end with D four or D nine. Now, remember these exceptions are normally observed with only the first row of transition metals. Of course, there are exceptions to this and we'll talk about them as we approach them. Remember this? This type of exception is normally reserved for the first world transition metals and those are the ones that end with D four D nine. Now here it says right, the condensed electron configuration for the following element Here we have chromium which has an atomic number of 24, which is why it has 24 electrons. So here we're gonna write out its electron configuration, It's condensed electron configuration in the wrong way and then we're gonna talk about the reasons why we have to change it to something that's more suitable. So we're gonna say that chromium, if we look at the periodic table, I'm up here, we're gonna say that are gone is the last double gas before we get to chromium Romy um is right here. So if we were to do the condensed electron configuration of chromium, we'd say argon for us two Because it's 1, 2, three D. Remember these are three D. So 23 S four, S five S six S. These are the Ds. And then over here these are your apps. These are your peas. So at this point we were acknowledging that we've gone over these types of concepts. You've seen videos of myself for this custom in class on how we have these different types of blocks for the periodic table. If you're still not familiar with them, make sure you go back and take a look at my videos on electron configuration. So chromium, it's condensed electron configuration will be argon for us to three D 1234. So four s 2, three D four. Remember we said that these exceptions happened with D four and D nine. What happens here is that the more correct way of writing the electron configuration of chromium would be argon four S one, three D five. one of the electrons from for us Move to the three D orbital's. Now, why would this happen? Well here we draw the way it would look if we kept the first electron configuration. So argon and for us, remember S has one orbital there's two electron, one spins up, one spins down And there are five orbital's for three D. So up up up up now if we draw it the more correct way, the correct way would be argon for us. Remember there's one here And then three D up up up up up. Alright, so why exactly is the second way this way better than the initial way? There are three reasons. The first reason is that four S and three D are very, very close in terms of energy because they were very close in terms of energy, we would say that they are degenerated so degenerate usually has a negative connotation to it. And chemistry just means same energy at least in terms of orbital's. And if you have the same energy, then we follow hans rule which says that we have Phil orbital's before we attempt to totally fill in an orbital. So that's why we're going up up up all the way through following Hunt's rule. The second reason that this is allowed to happen is that your D orbital's our most stable when they're either half filled all five of them or they're completely filled in in the first way we'd have one spot that's left unfilled. So it's not half filled. So here we want to half fill or totally felt your D orbital's. And then finally, the third reason is something that's new, which is electron electron repulsion. So we know that if an orbital has two electrons in it, following the poly exclusion principle, they have to have opposite spins where one points up on one points down because they have opposite spins are gonna repel each other. This is the electron electron repulsion that we're talking about where two electrons in the same orbital have opposite spins. This actually causes an increase in energy. So this is less stable because it has electron electron pair repulsion here. If we were to take one of these electrons and move into this empty slot here to give us this arrangement here, there is no longer any electron electron repulsion occurring. This helps to lower the energy over all of my orbital diagrams. Okay, so these are the three reasons why this version for the electron configuration of chromium is more stable and suitable than this version here. So keep that in mind. These three reasons why these types of exceptions exist for D four and D9 elements within the periodic table.