Hi guys welcome back to another video where we're revisiting ionization energy because this is the part I skipped because we didn't Know orbitals, but now we know orbitals we can do this and What we're doing is we're looking at why? Well, let's go back to here. We saw that it increases as we go along the period of the periodic table So as we go along a row so like I'm from Lithuania to Berlin to Voron to carbon nitrogen and fluorine It increases the ionization energy with some exceptions and We're gonna look at those exceptions now. So these exceptions are here of circlet them with giant red circles And as you can see these buck the trend so you'd think that Boron had a high ionization higher ionization energy compared to Brillium and oxygen had a higher ionization energy compared to nitrogen, but they don't and By looking at their orbitals we can find out why so Let's have a look at this Brillium. It's element number four and Boron is element number five So I can go to my trusted paint. I know that Boron has got four electrons and I can pencil this in And I can draw my four electrons onto Boron one two three four That's nice and stable. It's got four electrons both of its two S shells are completely filled When we go to no that's Borillium when we get a boron with five electrons we have to put it there This electron is in a much higher energy orbitals in the 2p orbitals all by itself Hence it's easier to remove Compared to this stable configuration of Borillium Hence Borillium with the four electrons here has a nice stable configuration It's low in energy meaning it's hard to remove one electron It's hard to remove this one electron from Brillium as it is to remove this higher energy electron that is in Boron So the meaning when we go back to here we can see that it takes less energy This is ionization energy here. It takes less energy to remove this high energy electron from Boron Then it is to remove the stable electron from Brillium Let's go up to nitrogen nitrogen's element number seven so we can just add a couple more things to this one two three four five six seven So that's nitrogen nitrogen has singly occupied all of its 2p orbitals and That's when it has them all filled singly occupied. That's nice and stable It minimizes repulsion and so it Minimizes energy it maximizes the repulsion of the electrons so it minimizes the energy of the system So it doesn't want to lose this one electron because it's it's nice and stable like this When we go to oxygen you have to double up on our electron So that is suddenly not as stable Hence this electron is easier to remove to remove and it takes less energy to remove this electron of oxygen Because remember with ionization energy we're just removing one electron So with nitrogen we've got all our p orbitals with one electron in them Oxygen with added one electron to our p orbitals which has got a pair Meaning it's now easier to remove the hence its lower energy The same thing happens over here with magnesium aluminium with magnesium we've got This so that's 10 electrons this is 12 electrons so this is magnesium. Oh, sorry. This is yeah, this is magnesium with 12 electrons aluminium has got 13 electrons and this is happening there is one electron in its 3p shell Meaning it's easier to remove so aluminium has got a lower ionization energy as we can see here aluminium is got a lower energy compared to magnesium in terms of how much energy it takes to remove one electron from it Next one phosphorus and sulfur Same thing happens here. This is phosphorus Which is nice and stable because it was all singly occupied and we double them up here This is suddenly Easier to remove to get back to this stable configuration So let's look at the notes for that In group two and 13 elements there's a small decrease in ionization energy is on with 15 and 16 So by group two I mean between this and this Since this is group two in this week group 13 or group three depending on if you account the transition metals in between So this is between group two and three this between group five and six there's a small decrease Not between here. This is five and six These are some notes for your book to put in as an element moves from having electrons in an s orbital to a p Subshell or s subshell to a p subshell for example brilliant to born and magnesium to aluminium the ionization energy decreases Electrons in the p subshell have a higher energy than those in the s subshell and an easier to remove Cool as an element moves from having an unpaired electrons Having unpaired electrons the paired electrons in their p subshell so I'm going from nitrogen to oxygen and phosphorus to sulfur The ionization energy decreases There is extra electron electron repulsion from the paired electrons Making the paired electrons easier to remove the the new paired electron. It's easier to remove and hence has a lower ionization energy Fantastic So there'll be a small task in the Google classroom to help explain this and get these notes down and have a great day. Thanks for watching