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Current time:0:00Total duration:12:33

AP.Chem:

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we've already looked at the electron configurations for potassium and for calcium but let's do it again really quickly because it's going to affect how we think about the D orbitals and so we find potassium which is in the fourth period on the periodic table and if we do noble gas notation to save some time we work backwards and the first noble gas we hit is argon so we write argon in brackets and we know argon has 18 electrons and potassium has 19 electrons so potassium has one more electron than argon and so we put that extra electron into a 4s orbital because for potassium the 4s orbital is lower energy than the 3d orbitals here so we have increasing energy and that electron goes into a 4s orbital so the complete electron configuration using noble gas notation for potassium is argon in brackets for s1 4 but for calcium I should say alright we have one more electron than potassium and so that electron is going to go into the 4s orbital as well and so we pair our spins and we write the electron configuration for calcium as argon and brackets for s2 for the calcium two plus ions so if you think about forming an ion here we're talking about the electron configuration for a neutral atom meaning equal numbers of protons and electrons and so with an atomic number of 2020 protons and twenty electrons if we lose two electrons right we have a net positive two charge so we form the calcium two ion the two electrons that we would lose to form the calcium 2 plus ion are are these right these two electrons right here in the 4s orbital and so the electron configuration for calcium 2 plus would be the same as the electron configuration for the noble gas argon here alright so for potassium once we accounted for argon right we had one electron to think about for calcium once we accounted for argon we had two electrons to think about and so if we go to the next element on the periodic table that's scandium right that's one more electron than calcium so we have three electrons to worry about once we put argon in here like that and this is where things get weird alright so now now we have to think about the D orbitals and once again things are very complicated once you hit scandium because the energies change so when you hit scandium even though these are very small energy differences now the energy of the 4s orbital is actually higher than the energy of the 3d orbitals and so we're talking about once again increasing energy and so that's pretty weird all right so if you think about these three electrons right where are we going to put them well your first guess if you understand these energy differences might be okay well I'm going to follow hoons rule I'm going to put those electrons in the lowest energy level possible here and I'm going to not pair my spins and so I'm going to write my electron configuration like that for scandium and so you might think it would be argon 3 D 3 but that's not what we observe for the electron configuration for scandium so actually two of these electrons actually move up to the higher energy orbital so two of those electrons move up to the 4s orbital here like that and so the electron configuration turns out to be 4s 2 3 D 1 so it's actually 4s 2 3 D 1 or if you prefer 3 D 1 4s 2 once again with argon in front of it so either one of these is acceptable and so this is weird like why do those electrons why do those two electrons go to an orbital of higher energy there's no simple explanation for this right so even though it might be higher in energy for those two electrons it must not be higher energy overall for the entire scandium atom and so there are many other factors to consider so things like increasing nuclear charge right so that scandium has an extra proton compared to calcium and and then there are once again many more factors and far too much to get into in this video and unfortunately there is no easy explanation for this but this is the observed electron configuration for scandium so how do we know this is true how do we know that the 4s orbital is actually higher energy than the 3d orbitals we know this from ionization experiments for example if you form the scandium plus one on the electron configuration for the scandium plus one ion so we're losing an electron from a neutral scandium atom this turns out to be argon 4s 1 3 D 1 or once again you could write argon 3 D 1 for s1 so where did we lose that electron to form our ion we lost that electron from the 4s orbital right we had 4s 2 here and here we have 4s 1 so we lost we lost this electron and that only makes sense if the 4s orbital is the highest and energy because when you lose an electron for ionization you lose the electron that's highest in energy that's the one that's easiest to remove to form the ion and so the 4s orbital is actually higher in energy than the 3d orbitals you don't see this a lot in in textbooks and I think the main reason for that is because of the fact that if if you're trying to think about just writing electron configurations and so your goal is to write let's say you're taking a test and your goal is to write the electron configuration for scandium and so the easiest way to do that let me go ahead and use red here so the easiest way to do that if you want to write the electron configuration for scandium alright you look at the periodic table and if you're doing noble gas notation write the noble gas that precedes it is of course argon right here and so that takes care of the argon portion and then looking at the periodic table you would say oh this could be 4s1 4s2 3d one so that gives you the correct electron configuration argon 4s 2 3 D 1 but it's it's implying that the D orbitals the 3 D orbitals fill after the 4 s orbital and is therefore and is therefore a higher energy and that's not that's not completely that's not true actually and so it does help you to to just assume that's the case if you're writing an electron configuration but that's not what's happening in reality and we need to think about the other elements here so we just did scandium right so next let's move on to titanium thinking about titanium so the next element in the periodic table if your question on test was right the electron configuration for titanium the easiest way to do it is just once again to think of argon so put argon in brackets and then think to yourself this would be for s1 this would be for s 2 this would be 3 d1 and this would be 3 d2 so you could write 4 or s 2 and then 3 d2 or once again you could switch 3 D 2 and 4s 2 and so once again this is implying this is implying the D orbitals fill after the 4s orbital which isn't true but it does get you the right answer and so it's useful to think about it both ways it's useful to think about the energy levels properly but the same time if your goal is to get the answer the fastest way possible looking at the periodic table and running through the electron configuration might be the best way to do it on a test so let's look at some of these other elements here so we've just talked about scandium and titanium all right so let's go down here and let's look at this little setup here alright so we just did we just did scandium and titanium alright so scandium was argon 4s 2 3 D 1 so we talked about two electrons in the 4s orbital one electron in 3d orbital and then we just did titanium 4s 2 3 D 2 or once again you could switch any of these and when you're doing orbital notation right adding that second electron to 2a D orbital here's the electron that we added so we didn't pair up our spins who are following hoons rule here next element is vanadium so we do the same thing one more electron we add that electron to a D orbital but we add it to we don't add it to one of the ones that we've already started to fill here right we add that electron to to another D orbital and so once again following one's rule so things get weird when you get to chromium so let me use a different color here for chromium so if you're just thinking about what might happen for chromium chromium one more electron to think about than vanadium alright so you might think oh let's just add that one electron to to a 3d orbital like that and then be done with it so so if you think about it you might guess 4s 2 3 D 4 so let's go ahead and write that so 4s 2 3 D 4 so question mark but that's not actually what we get we get 4s 1 3 D 5 so that electron so this this electron here and let me go ahead and use red so you can think about this electron expected it to be there we expected before us to 3d for it's like that electron has moved over here to this empty orbital to give you this orbital notation alright so that's that's just an easy way of thinking about it and in reality that's not what's happening if you're building up the atom here because of the different energy levels but just to make things this is make things easier when you're writing electron configurations you can think about moving an electron from the 4s orbital over to the empty the last empty D orbital here so some people say that this this half-filled d subshell let me go in circle here this half-filled d subshell is extra stable and that might be true for the chromium atom but it's not always true so it's not really the best explanation the real explanation is extremely is extremely complicated and actually just way too much to get into for a general chemistry course so next element is manganese right let me go ahead and stick with blue here so manganese one more electron than chromium here so chromium we had we had six electrons here at manganese we need to worry about seven electrons and so this is this is kind of what we expect right just going across the periodic table let me go ahead and do this for for manganese let me use green here so you might say okay that's for s one that's for s 2 and then 3 D 1 3 D 2 3 D 3 D 3 d 4 3 D 5 so you might say to yourself 4s 2 3 D 5 and so this kind of proceeds how we would expect it to all right and the same thing with iron so 4s 2 3 D 6 all right so that takes care of iron and once again now you can start to pair up your spins all right so so while we've seen that in earlier electron configurations next cobalt one more electron to worry about right so 4s 2 3 D 7 makes sense and you can see here would be the electron that we added and we paired up our spins again nickel same trend right so we add one more electron 3d8 so that makes sense here's the electron that we added and once again we get a weird one right so when we get to copper alright so copper you might think let me let me let me use red for copper so we know copper is red so we think about it writing 1 our electron so if we took the electron configuration here for nickel we added one more electron you might guess that would be the orbital notation for copper but that's not what we see we've taken this electron here right and moved it over to here like that and this gives us this gives us say a a filled a filled d subshell here so once again one explanation we'll see for that is that this kiss is extremely stable for copper and that might be true for copper all right and that leaves us only one electron here in our 4 s orbital and so the electron configuration is 4 s 1 3 D 10 but all these general chemistry explanations are just a little bit too simple for reality but you know if you're just starting out there they're pretty good way to think about it and then finally and then finally zinc zinc makes sense right we're adding one more around one more electrons we just took care a copper for zinc we have one more electron and so you could think about this being this being 4s 2 right here and then we have 3 D 10 1 2 3 4 5 6 7 8 9 10 so 4s2 3d10 or 3 D 10 4s 2 with argon in front of it gives you the complete electron configuration and you can see you've now filled your 4s orbital and your and your 3 D orbitals like that so once again pretty complicated topic and hopefully this just gives you an idea about what's going on

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