Orbital Box Diagrams

so in a previous video we took a look at something called electron configurations both expanded electron configurations and condensed electron configurations now really this kind of accounted for two of the quantum numbers one being that principal quantum number which assigned that number to a particular energy level and also our secondary quantum number or angular quantum number or shape quantum number that really gives us the S and the p and the D distributions of those orbitals but we do have some other things that we have to sort of account for and there are other ways to come up with electron configurations that we can use to help us depict say even bonding and that tool that we use is just kind of an extension of electron configurations and we refer to these as orbital box diagrams and so these take our electron configurations one step further because each box now represents a maximum of two electrons that can be assigned for any particular orbital you see an orbital can have a Max maximum of two electrons so if we look in a particular energy level and in a particular Su each of those is going to have a maximum number of orbitals and therefore maximum number of electrons that they can have and we already kind of have an idea of this because of the electron configurations that we drew let's take a look at lithium as an example so we can see that lithium has three electrons associated with it from its atomic number and as we saw with the electron configurations that's going to be 1 S2 and 2 S1 but but what we're going to do for our orbitals now is we're going to use boxes to represent these orbitals so we have a box representing the 1s and we have a box representing the 2s and in that first orbital we have two electrons and what we're going to do to represent that is we're going to draw two arrows now we're going to draw them in opposing directions because this sort of relays to something that we refer to as P's Exclusion Principle where no two electrons can occupy the same space or have the same set of four quantum numbers we're going to apply it here and indicate that they have opposing spins now you might remember with our um opposing spins quantum number plus a half minus a half instead we're going to have one Arrow going up and one Arrow going down typically the first electron that's assigned to an orbital is pointed up then as we move now to the 2s we're going to have another arrow representing that third electron now it you can see doesn't have another electron that's pairing up with it so as a result it's going to be the only electron in there now if we take a look at say something like um oxygen that is going to have some electrons that are in its 2p suev the 2p because it can contain a maximum of six electrons is going to have three boxes to represent the three orbitals that are in the 2p subl so as a result we are now going to draw one box for the 1s one box for the 2s we're going to separate a little bit draw three boxes for the 2p and we're going to continue to fill with electrons so now we're going to put another electron in the 2s we are now going to put another electron in the 2p but before we continue on to put another electron in the 2p what we're going to remember is that we're going to put it in a separate orbital first and this is according to something called Hun's rule where we are going to fill all the available orbitals within a particular energy level first we're not going to put them into the same orbital because being electrons they are potentially going to interact so if they can move to another orbital or occupy another orbital within the same energy that is they don't have to expend any energy they are going to do so so you can see that the orbital box diagram for oxygen is going to look something like this and again if we are going to count up all of those electrons we can see that we should have eight electrons here now once we have each of the orbitals within a particular energy level filled and we still have electrons to assign we can then start pairing them up so if we get into more complex examples like again my friend tungsten we can use condensed orbital box diagrams and so the condensed orbital box diagrams takes the electron configuration of tungsten from this into something that looks like this and we can see that we have the appropriate condensed orbital box diagram for Tungsten and again we can count up all of those electrons to make make sure that we have those 74 electrons accounted for now with orbital box diagrams what they allow us to do is start to explain bonding a little bit better you see if we use orbital box diagrams now we can start to illustrate bonding as the bonding occurs between sus and we can use something simple like sodium and chlorine forming sodium chloride to help us depict this we can see that sodium being a group one Alkali metal is going to show a tendency to have one electron taken from it and the chlorine which is a group seven hogen is going to show a tendency to take an electron and you can see as it takes an electron it is then going to fill that P suble that it has and then the sodium is now going to have one electron taken from it and so it's going to empty the s subl that it has and now these two become isoelectronic that is having the same electron configuration as the nearest noble gas and we can depict these now as ions sodium becoming a cation a positive ion because it has had an electron taken from it and we have our chloride ion an anion a negative ion because it has taken the electron and so now we can see a way that we can use orbital box diagrams to depict what's going on in an ionic bond so very quickly an excellent example of multivalent metals and how we can use orbital box diagrams to explain it is to take a look at iron you see Iron has two very common valences three and two and if we take a look at how the iron would become a 2+ ion or iron 2 it's because two electrons are being taken and if we just took the outermost electrons at least on terms of how we assigned them we would take two from the D now if we took two from the D that would kind of leave some sort of half-filled orbitals and some empty orbitals all within the same sublevel and that's not what's observed what's observed instead is that two electrons are taken from the S and that we have an empty 4 s orbital and that we have the six remaining electrons occupying the D now what happens with the 3+ and perhaps you can start to picture it is that we have those two electrons taken from the S and then one taken from the D and in terms of how this Arrangement looks now look at this we have an empty suev in the S and we have entirely half-filled orbitals within the D and so you can see there's a bit of symmetry here and the belief is that if we have a suev that is uh fully filled or half filled or empty it provides greater stability for that particular ion or element than say something that's partially filled now I don't want you to go moving electrons around and thinking that this is always going to be the case we're limited by what nature does and again this is just our method of illustrating what we observe all we can observe is what happens in nature and then try to fit that to our constructs not the other way around so don't be trying to move around electrons because it looks a little prettier but it does allow us to explain what is observed in nature and it is observe that iron has these electron configurations in order to form 2+ and 3+ respectively it should be noted that when elements bond the electrons are going to empty from the highest principal quantum number first so if we're looking at something say that has an electron in the 6s or two electrons in the 6s and has electrons in the 5D it's more than likely the these electrons are going to be removed from the highest principal quantum number first that is the six before they're going to be removed from the five so some periodic tables will actually show the order of electron emptying or what the configuration of that element is going to look like once all of the electrons have been assigned to it and it should be noted that as we assign a particular electron they're negatively charged so they are going to interact and they're going to change the distribution of those orbitals so again the way that we're going through this is in terms of the AL bow principle and electron filling but you will from time to time come across periodic tables where it's as I would call it more in terms of electron emptying or where electrons would be taken from in terms of bonding so hopefully after watching this video you have a better understanding of not only electron configurations and how they translate into orbital box diagrams but how orbital box diagrams can be used to explain things such as ionic bonding and potentially multivalent elements thanks for watching

Channel: JFR Science

Views: 77,936

Keywords: Chemistry Help, High School, John F Ross, SCH3U, SCH4U, Chemistry Support, Chemistry Tutor, Orbital Box Diagrams, Electron Configurations, The Atom, Quantum Theory, OBDs, E Configs, E Configurations, Atomic Theory, Electrons, Bonding, Multivalent Metals, Multivalency

Id: EUvzNFa3WOg

Length: 9min 5sec (545 seconds)

Published: Thu May 11 2017