hello General organic chemistry students in this video we're going to talk about the concept of electron configurations bonding and formal charge now before we get into that what I'm showing you right here is a beast of a molecule this is pretty big now some of you might not remember line angle formula or been exposed to it in your general Kim class so this looks totally foreign to you but some of you have covered line angle formula now we are going to get to line angle formula what it means and how to read it but just to tell you the very brief parts of it every time you see an angle like this that represents a carbon if it's just a straight line at each end of that line is a carbon at the end of it and that's the basics of line angle formula more details to come on that later on for sure now looking at this really big molecule the first question I'd like to ask is does anyone know what this molecule is and I know you can't respond to this because it's a recorded lecture but I want you to think about it hm does it look familiar and it might not what this molecule is is an adenine dinucleotide nicotinamide now that seems kind of weird the abbreviation for this is n a D+ now nad+ is an electron carrier in the body what nad+ does is that that H collects a hydride from an enzyme doing an oxidation reaction and it takes NAD and adds a hydride to this top position because as a whole will have resonance down to this lower nitrogen that's why we get that Regio selectivity of the hydride addition so n becomes nadh for a hydride that goes into the Matrix of the mitochondria donates it to complex one and through a Cascade of reaction forms 3 atps so what we're looking at right here is the electron carrier in our bodies that is responsible for the formation of ATP a Denine triphosphate and we'll see that structure later on in this class now looking at this we're going to go through and figure out why is this nitrogen positive why does this carbon have four bonds attached to it and why should it have four and not three or two bonds attached and that's all going to come down to electron configuration and bond types so what I want to do first of all is look at our trusty friend the periodic table of the elements right here now in the periodic table the elements and go ahead and please excuse me you will hear my cat from time to time she likes to chime in when I'm recording videos so you might hear a meow here and there when we look here we have seven rows on the periodic table Row one all the way to row seven now what we want to focus on is what type of orbitals are present in these shells now we deriv this in general chemistry we can go back to quantum chemistry and do that we don't need to do that here I want to give you the basics of what we need to continue on with Organic Chemistry so if you notice in this first shell we have two boxes in the first shell so what that means is is that we have one s orbital in the first shell in the second shell we see 1 2 3 4 5 6 7 eight boxes now each one of these boxes represents an electron so if we take that divide it by two because every orbital can hold two electrons we have to come up with four orbitals of some type so in the second shell we have another 1s orbital and we have three p orbitals in this second shell in the next one we have 1 s and 3 PS as well in the third shell but we do have access to the D orbitals per quantum chemistry so we can start accessing some of that D orbital system now in this fourth shell right here we still have 1 s orbital 3 PS and 5 D orbitals now I'm just telling you the basics behind this again if you want to go back and review how these orbitals come about you can review your quantum chemistry from General chemistry I will never ask you to derive the orbital types now let's not forget in this third one we can have access to 5D orbitals as well it is possible to access those now we're not going to go much further than that because as a whole the D orbitals for organic chemistry we don't care about them whatsoever we've focus on the S and P orbitals so the big take-home message right here is that the first row has one s orbital the second row has one s orbital and 3 p orbitals so what we're noticing here is that every row has one s orbital starting at row two and Below every one of those has three p orbitals so if we look look at Row one again there's two boxes each one represents an electron oh those are two electrons in One S orbital oh in the second row we have eight boxes right here divide that by two four orbitals needed oh lucky for us we have 1 s and 3 p orbitals right here now what do these S and P orbitals look like when we compare an S orbital to a p orbital I'm going to try to draw these relative to one another an S orbital is an entire sphere a p orbital looks like the figure eight that's kind of been elongated on both sides now the length of them should be the same if I was a better artist they would be now right off the bat when you look at these which one has more volume in it the s or the p orbital and it's clearly the S orbital so we have more volume present in the S orbital so if we had to place two electrons in each one of these orbitals so here's one electron here's another one and there's two right there which orbital allows more room for the electrons to move more freely around them and that's the S orbital so s orbitals are more stable than p orbital so when there's a choice electrons will go to the S orbital before the P orbitals and that's all because of the volume of the orbital space so s is more stable than P now what's the difference between the S orbital in the first shell and the S orbital in the second shell size so let's go ahead and move this down and talk about that a little bit so here if we have our nucleus right here here is the first shell so here's the one s orbital notice I'm putting that comma in or that little dashed line here is the second shell larger than the first shell and there is the 2 s orbital so here the two is referring to the second row row where here the one is referring to the first row there's still only one s orbital per row so this is not saying two s orbitals there is only one now if we lay on top of this the P orbitals here's one p orbital here here's another p orbital here and there is a third p orbital and once again this part should not be going outside the circle it's all within that same Realm of the second shell we can then put the third shell on top of this the fourth shell and the fifth shell it gets bigger and bigger and bigger each one allowing electrons to reside within it but keeping in mind each orbital okay so every time you see an orbital can hold two electrons at Max it doesn't have to hold two it could hold one or have no electrons in it but it cannot exceed two electrons that's impossible so what I'd like to do right now is talk about the basic electron configurations of the atoms as drawn or present on the periodic table which means a net neutral charge which means the number of protons equals the number of electrons giving us a formal charge of zero how do we calculate formal charge we'll get to that in a little bit so let's go ahead and look at hydrogen first now hydrogen is in that first shell which means there is one s orbital present now notice here I'm not putting the hyphen in between it I'm just saying 1 s I'm writing the electron config configuration for this hydrogen atom in that hydrogen atom there's one box which means there's one electron and we write 1 S1 do not leave it as 1 s blank because the blank implies no electron so we have to put the one down if we look at Helium now we're still in the first shell the S orbital and now we have two electrons so what's the big difference between the hydrogen and the helium atoms helium has a filled veence shell now what's a veent shell a veent shell is the outermost filled shell on an atom it doesn't have to be completely filled but so I should say the veence shell is the outermost shell that has electrons on the atom here we happen to be in the first shell and both let's go ahead and expand our electron configuration to lithium now lithium is on the second row which means the first row has to be filled and then we access the second row and we have 2 s and we have 32ps now there's one electron to place in this second shell do we dump it into the S orbital or the P orbitals and it goes into the S orbitals so we don't even have to show the two PS because they're not filled with electrons so the veence shell here is the second shell because it is the outermost shell with electrons the first shell is not the veent shell here but it is for hydrogen and helium it's a case by case basis now let's go ahead and make this a little bit harder let's go ahead and go straight towards oxygen now oxygen right here has eight protons which means we need eight electrons to fill up the first shell has to be filled let's look at the second shell and that should be a two so we have 1 s and we have 3 p orbitals we have six more electrons to place we're going to place two electrons in the s because it's more stable which leaves us four electrons we can place two in this first one and go one and one we would not put two electrons in this p orbital and zero electrons here because electrons repulse one another so they don't want to fill the same shell or the orbital unless it needs to we fill the 2s before the 2 PS because of the pure volume of the orbital itself and the stabilization that's present there and that's what we have in oxygen now let's look at the veent Shell of oxygen in more detail this veent shell right here I'm go ahead and highlight it in red how many more electrons do we need to fill the veence shell and that's two we need one more electron in this p orbital and one more electron in that p orbital so what does that mean it can form two bonds that's why oxygen has two bonds we can sometimes see three bonds we can sometimes see one Bond but two is the magic number for it so if you look at oxygen once again on this periodic table it is two boxes away from the noble gases the noble gases get their name because their veent shells are filled and they're chemically inert they don't do anything they're called Noble because they're supposed to reflect the kings and queens of Europe that did absolutely nothing as their request because they were supposed to be perfect atoms so two boxes away which means we need two electrons so with that idea in mind if you look at Florine which is one box away how many bonds does it need you got it just one but let's look at its electron configuration to justify that so Florine its first shell is completely filled and we're looking at the second shell so we write the three p orbitals now in the second shell we have 1 2 3 4 5 6 seven electrons to place two go in the S two go in the first P two go in the second P for a grand total of six and there's one and here is the orbital that can do bonding on Florine to give it its one Bond needed to look like a neon atom so when we're looking at this all the atoms on the periodic table desire to look like noble gases but the electron configurations that we ask for when we just say write the electron configuration for oxygen we mean per its periodic State as shown on the periodic table if we want to write these in its true through real life State we're going to talk about that in a moment but I want to continue on our discussion of electron configurations now over here on the side and I apologize for kind of doing this like this CU once I go to the next slide I can't go back with these recordings in this program I want to look at carbon carbon has six electrons so we put two into the first shell 2 into the 2s and we go 2 P1 2 P1 2 p 0 and this is what you were taught in general chemistry for the electron configuration for carbon by itself and if it's truly by itself this is correct but if it's actually in play in real life this is incorrect the actual electron configuration for carbon is this 2s1 2p1 2p1 2 P1 wait a second how on Earth did one of these electrons flip to the p orbital P orbitals are less stable than S orbitals as electrons add into the orbitals it increases the energy of the orbital so once the two electrons are present and we fill these two this orbital is almost equal in energy to this p and the electron flips and it moves to the p orbital and that's how we get this electron configuration for Elemental carbon in real life now this is what we taught you in gen Kim but this is the truth behind carbon in real life all right so with this information right here we've talked about veent shells and how to figure out the number of bonds but the one thing that we haven't talked about yet and I'm going to move this up here oops maybe I'll move this up let's see this is always the tricky part on this program there it is I want to talk about Lan pairs what is a lone pair of electrons a lone pair of electrons Rons are two electrons in an orbital so one orbital that are nonbonding meaning they are not bonded to anything else it's just by themselves so a lone pair of electrons are non-bonding electrons not bonded to anything else we just talked about Oxygen's electron configuration up above let me write it down again 1 S2 2 S2 2 P2 2 P1 2 P1 now regards to its veent shell and that's all that we care about we don't care about this first shell because we can't get to it until the second shell is devoid of electrons and that's not going to happen we're not going to get rid of all of Oxygen's electrons when we look at this veence shell what we're noticing here is that this s orbital and this p orbital have two electrons in them these are lone pairs of electrons both of them right here there are two lone pairs of electrons on oxygen these two over here these P orbitals are what we call bonding electrons meaning that they want to interact with other atoms so that they can actually form a Cove valent Bond and we're going to get back to how this bonding occurs in a little bit but I really just want to hyperfocus in on what is a non-bonding versus a bonding electron system so the non-bonding is not participating in a bond to another atom where bonding electrons form a bond to another atom so with that let's go ahead and look at floring up above how many lone pairs are on Florine now we don't count the first shell because we only care about the veent shell so we have one 2 three lone pairs on Florine with one open orbital so now if I was to scroll back up oops this is where the once again this app gets a little tricky here oh dear nope bear with me for a second as I try to figure this out nope ah nope when we look at the periodic table of the elements what we're looking at here is a lot of information and if we look at this second row right here we have to have eight electrons to fill that shell now keep in mind only neon has eight electrons in the s and p orbital system so when you look at nitrogen for example it is three boxes away from neon isn't it which means it has five electrons in its l in its veence shell so five electrons in the veence Shell so when we try to write the electron config configuration in the veent Shell we're looking at a 2s2 2p1 2p1 2p1 one lone pair on nitrogen three bonding orbital systems right here this is going to be the same for phosphorus arsenic antimony bismuth and so on and that's the periodic trend down the table we just said that Florine has three lone pairs chlorine has three so does bromine iodine and animon as well oxygen right here has two guess what so do sulfur and selenium selenium and polonium and that's our periodic trends coming back into play in a wonderful way now let's go ahead and look at the electron configuration for Boron see if we can move this up nicely okay so if we look at Boron we have three boxes in so the first shell is filled the second shell third shell not I'm sorry not the third shell but the second shell with the three p orbitals we we have three electrons that we have to place and what happens here is that we get one electron in the S one in the P one in this other p and this p is devoid of electrons boron has a great tendency of violating the octet rule it only needs six electrons in its veent shell to be satisfied because these three orbitals are bonding it does not use this p orbital whatsoever and this is kind of laying the seed for a reaction called hydroboration that we'll talk about in a little bit having an open p orbital making it an electrophile a lot of big terms right there but we're going to focus in and we're going to learn all these terms don't you worry so now that's the basics of electron configurations in real life not real life I'm sorry based upon the periodic table we now want to focus on these in real life so if we were to see oxygen in real life oxygen could have a min-2 charge meaning if picked up two electrons from somewhere to fill its veence shell to look like neon it desires to look like neon because it's a full veence shell and therefore it's stable the electron configuration for this oxygen and ion is 1 S2 2s2 2 p oops Got ahead of myself there 2 P2 2 P2 and 2 P2 a completely filled shell now this is what we call the expanded electron configuration showing every little bit of detail can we condense this down some absolutely and we do have it called the condensed electron configuration so the condensed one is 1 S2 2 S2 2 P6 where I'm not saying What electron is in which p orbital I'm just saying that there are six electrons in the three p orbitals and that's fine to do if I asked you right write the electron configuration for oxygen in its real life State meaning -2 charge if you wrote 1 S2 2 S2 2p6 that is completely fine we can also use the Noel gas configuration it doesn't help us too much right now but it will for bigger elements where we put the noble gas for the shell that's filled such as helium right here and then we put 2 S2 2 P6 now for oxygen with 2 minus what we would actually do is draw it as NE in these brackets meaning it looks just like neon's electron configuration now let's practice this a little bit more what I would like you to practice right now is take nitrogen and put it in a minus three state what's the electron configuration for this nitrogen anine nitrogen normally has five electrons in its veent shell we've just dumped in three more from a variety of sources which means the first shell is completely completely filled and let's write all the orbitals present in the second shell if we dumped in three we had five initially that means we have eight electrons it's a filled shell it looks just like the electron configuration of neon and that's how we write electron configurations of charged species what about lithium plus that's a tricky one lithium is in the second shell so it looks like we have one electron in that veence shell but if it's a positive charge it means that missed an electron so its veent shell is now 1 S2 the S orbital in the first shell lithium by itself though is 1 S2 2s1 so when you have a lithium cat ion it is the veent shell is the first shell when you have elemental lithium meaning in its Elemental form the veent shells in the second shell a big difference right there and those are overall the basic comparing and contrasting of electron configurations inside atoms themselves now what I like to talk about next is the bonding of these types of systems so let me open up a new page all right how does bonding work here and what types of bonds can we form in chemistry so the two big types of bonds that we have are calent bonds and ionic bonds calent bonds are made up of two nonmetals oops two nonmetals an ionic bond is made up of a metal and a nonmetal atom and that's the big difference remember metals like to give away electrons non-metals like to hold them so if you were to look at sodium attached to O right here and I want to look at that Bond sodium is a metal oxygen is a non metal this is an ionic bond so how does this species actually look like when it's in solution those electrons between that sodium and oxygen go to the oxygen so we have the negative charge here's the hydrogen and here's our sodium with a positive charge so oxygen has taken the electron from sodium onto itself and that's how we see the electron transfer forming these charged atom States so that's what an ionic bond does what type of bond is holding this oxygen and hydrogen together and right now you might be saying wait for the periodic table hydrogen is a metal but alas it's not hydrogen is placed right here for the electron configuration only it has the same electron configuration for all these group one elements in terms of its type of atom hydrogen should be over here on top of carbon hydrogen is a non metal so if you're not looking at electron configurations always assume hydrogen is right on top of carbon right here so these are two non-metals bonded together making a Cove valent Bond now in an ionic bond we saw that the more electr negative atom basically the non-metal steals all the electrons here we have an equ not really an equal sharing but we have a sharing of electrons between the hydrogen on the oxygen it's not equal by any means but a coent bond is defined as the sharing of electrons an ionic bond is basically the stabilization of charge and why is it the stabilization of charge this sodium being positive wants to be close to that oxygen being negative to neutralize their charges calent bonds share the electrons if it's equal or unequal that's another question so if you have an unequal sharing of electrons we have a polar calent bond if we have an equal sharing of electrons we have a non-polar coal Bond so this is equal sharing and this is Nal sharing of the electrons so if we go back and look at this coal bond between the oxygen and hydrogen is it calent yes is it polar or nonpolar this is a polar bond because the oxygen is more electronegative than hydrogen pulling the electrons closer to it further from the hydrogen and that's an important concept right there so let's go ahead and look at this molecule right here so in this molecule how many ionic bonds are present a being 1 B being 2 C being 3 D being four e being five and F being being none so in order to answer this we have to ask ourselves which Bond or which or or atoms is a metal bonded to a non-metal here we have hydrogen it's nonmetal non-metal so that's a calent bond coent coent calent calent coent Cove valent oxygen to potassium potassium is a metal so this Bond right here is an ionic bond and these are important distinctions to make in chemistry very important now we're going to utilize this as we're building up more knowledge in this class in terms of electron densities and basic chemical reactions what I'm hoping to do right here is just to go over the basic concepts of bonding types coent versus ionic bonds so with this basic concept of bonding done the last thing we want to talk about is formal charge when I showed you NAD there was a nitrogen with a positive charge on it let's go ahead and go to the next page and talk about that in a little bit more detail so n a I'm not going to draw the whole structure out but we had this part of the structure right here and it was connected to a carbon down below this nitrogen has a positive charge on it how on Earth do we know it has a positive charge so if the positive charge wasn't there and I was asking you to calculate the formal charge for all the atoms in this molecule this is how we do it to calculate formal charge we're going to use the following equation we're going to take the number of veence electrons present in the free atom meaning not not in this molecule just by itself in nature so there's that and from that we are going to subt subtract excuse me the number of Lone pairs of electrons on that atom in its present State minus one half of all bonding electrons let's go ahead and practice that right here let's look at nitrogen nitrogen in its free state has five electrons in its veent shell remember it's 1 S2 2s2 2p1 2p1 2p0 I'm sorry 2p1 all around all across the board so 2s2 2p1 2p1 2p1 for the five electrons as shown though we're going to subtract lone pairs that nitrogen if we look at it has 1 2 3 four bonds too many bonds on it so if it has four bonds Each Bond has two electrons in it that's eight electrons which means this nitrogen cannot have any lone pairs let me go through that again the nitrogen has four bonds being shown each bond has two electrons which means there's eight electrons on nitrogen as shown there are no lone pairs on this nitrogen so there's zero right here now I just said that there were eight bonds so we're going to take2 times the number of bonds in this molecule and there's two electrons per bond for a grand number of eight where did that eight come from again there's 1 2 3 four bonds each bond holds two electrons and that's where the eight comes from so we now have 5 - 0 - 4 12 * 8 what does this equal 5 - 4 is a +1 oxidation States and that's how we get the form noral charge for this nitrogen what I'd like you to do right now is I'm going to go ahead and show the hydrogen on it because we don't know line angle formula not all of us there's a carbon right there that little angle what's the formal charge of that carbon now carbon in its neutral State how many veence electrons are present in its veence shell there are four veence electrons there's absolutely no lone pairs in this species right there because there's four bonds right four bonds two electrons per bond is eight the veence shell is filled so we take2 minus those four bonds with two electrons in it for eight electrons we get 4 - 0 - 4 for a formal charge of zero and that's why carbon has no formal charge in this structure absolutely no charge whatsoever now with that in mind that ends this video right here talking about electron configurations bonding and formal charge please practice problems you can look at KH Academy look in your textbook even if you want you can go to the internet and type in formal charge of atoms and you will see lots of resources I highly recommend that you only look at academic websites not other ones other ones are paid services and they're also not corrected or updated so there could be wrong information on them academic sites are constantly being updated and corrected if there's any errors in them um I hope each of you are doing well I look forward to talking to you in our uh discussions and have a wonderful evening