![]() Okay let's do the orbital diagram for iron, iron we know is on its ground state of 26 electrons, so we know the first electrons are going to go into the 1s orbital and we said 2 electrons can fall into the 1s orbital. Let's put all these stuff into play, how this all come together. So just like electron are both negative they're all negatively charged they're not going want to be really close to each other, so they're in equal energy they're going to occupy all the energy levels of that same energy first before they pair up because they typically don't like being really close to each other. Lastly the Hands Rule states that they must occupy all orbitals of equal in energy before pairing up. It can hold 1 but it cannot hold more than 2. So an orbital can only hold 2 electrons and that's it no more. The second rule that we're going to talk about is the Pauli Exclusion Principle which basically states that there are maximum 2 electrons per orbital. Okay so I'll actually set this into play and in just a second. Then they're going to after the 4s then they're going to fall into the 3d and then the 4p and just keep pushing the arrows down if you have more principle energy levels so on and do forth. Then they're going to fall into the 3p orbital and then the 4s orbital notice we skipped the 3d they did not go arcos like that we're following the actual arrows not across. Then we're going to say "okay electrons are going to fall into the 1s orbital first the lowest in energy as Auf Bau diagram described." Okay falling into the 1s orbital, then they're going to fall into the 2s orbital after that okay great then they're going to fall into the 2p orbital and then the 3s orbital making our diagonal. I'm not going to draw the entire thing but you get how it comes out. So how are we going to remember which one is lower in energy than the others? Instead of having to carry this around and have to refer back to this actual diagram there is an easy way to remember which way, how electrons fall in the orbitals and in the sublevels.Īlright we're going to make this chart which might seem familiar from the classroom, the first energy level is the 1s orbital then we're going to do the 2s and the 2p then the 3s, 3p, 3d okay then the 4s, 4p, 4d, 4f and so on and so forth. Okay as we go up and we jump up to the 3s orbital then the 3p orbitals then up here it gets all funny like all the things, all the sublevels and all the orbitals kind of get jumbbly in terms of like what you predict. Notice that there's 3 orbitals within the 2p sublevel we have learnt that before, that makes sense. You jump up a little bit in energy and we get the 2s orbital that make it the 2p sublevel. So let's look at the Auf Bau diagram which actually show this for us, okay so down here we have the 1s orbital but the 1 dash indicates that there's 1 orbital within the 1s sublevel which makes sense that it is the lowest in energy, it's its first principle energy level. ![]() Then we have to think okay with the sublevels, I mean the orbitals how are they falling in terms of like energy which one is lower in energy, which one is higher in energy. The first one being the Auf Bau Principle, the Auf Bau Principle states that each electron occupies the lowest energy orbital available. In order to figure out where electrons go in an atom we have to follow 3 main rules. Orbital diagrams are a pictorial description of electrons in an atom. Alright let's talk about orbital diagrams. ![]()
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