 In part two of this lecture we'll learn how to write electron configurations using Schrodinger's orbitals rather than Bohr's orbits. So with this complicated series of orbitals from Schrodinger's model how do we go about writing electron configurations? It's actually not hard at all. There are a few simple rules that you need to follow and we'll look at the first two of these rules first. The first rule is to start at the bottom and work your way up. Put the electrons in the lowest energy orbital first and work up to higher energy orbitals one by one. This is called the Aufbau principle. The German word Aufbau means to build up. Second, as we've just been discussing, each orbital may only have two electrons in it. So once you've filled an orbital with two electrons, move on to the next one. So let's try a few examples using these two rules. First, let's try hydrogen. This is pretty easy. We've got one electron and we follow the first rule and we put it in the lowest possible orbital, which is the 1s orbital. The electrons depicted as an arrow in the orbital. We don't always want to draw a full diagram to represent the electrons. So to write the electron configuration in shorthand we write 1s1. The first one, the big one, indicates the energy level, the s indicates the kind of orbital and the little one tells you that there is one electron in that orbital. Next let's try neon. Neon is the 10th element on the periodic table, so it has 10 electrons. So we start at the bottom and we fill up the orbitals until we've got 10 electrons. As with hydrogen, the first electron is represented by an arrow. The second electron in the same orbital, two electrons per orbital, is represented by an arrow in the opposite direction, going down. It has to do with the fact that the electrons have what's called opposite spin to each other, but this is not something we need to explore in this course. The electrons in the following orbitals are drawn the same way until we have all 10 electrons accounted for. So you can see that we've completely filled the 1s, the 2s and the 2p orbitals. The shorthand for the electron configuration is then written 1s2 because there are two electrons in the 1s orbital, 2s2, two electrons in the 2s orbital, 2p6 because there are six electrons in the 3 2p orbitals. Okay, example number three, nitrogen. Nitrogen is the seventh element, so seven electrons. As before we fill up the 1s and the 2s orbitals and now we have three electrons left. Perhaps the obvious thing to do at this point would be to completely fill one of the 2p orbitals and half fill another. But this is where the third rule comes in. The third rule for writing electron configurations is that when you have a group of orbitals with the same energy, like three p orbitals together or five d orbitals, the orbitals get filled with one electron each first and only then do the electrons start doubling up. The reason for this is that two negatively charged electrons occupying the same space will repel each other, which is unfavorable. This repulsion is minimized if you spread the electrons out. So let's try nitrogen again, seven electrons to distribute. We fill up the 1s orbital and then the 2s orbital and our three remaining electrons we distribute evenly between the three equivalent p orbitals. Our shorthand doesn't distinguish between the three p orbitals so our configuration looks like one s2, two s2, two p3. So we're just saying that there are three electrons in the three p orbitals. So last example, sulfur. Sulfur is the 16th element so we have 16 electrons to distribute. So we fill the 1s and the 2s, all the 2p orbitals, the 3s and so far that means we've used up 12 of the 16 electrons. The four remaining electrons are going to be put into the three p orbitals. First we put one electron in each and now we have one final electron left over so we put that into one of the three p orbitals. It doesn't matter which one but we put in the left most box by convention. So the shorthand for this will be 1s2, 2s2, 2p6, 3s2, 3p4. You may notice that for larger atoms things are going to start getting a bit complicated. For instance, if you look carefully in this diagram the 3d orbitals are actually slightly higher in energy than the 4s orbital. Which means the 4s orbital actually gets filled before the 3d despite technically being in a higher energy level. However for this course you don't need to be able to do electron configurations up this high.