 What we're going to cover today is a lot of what you, if anybody took the 51p class over the summer, it's a lot of what we're talking about for that class. And so we're going to talk about how to waste and draw organic molecules. We'll be talking about fluid structures, how to draw fluid structures for organic molecules. And then that's what we're going to be working on in this discussion. Later in the week they talk about how to draw reaction specters and then honestly we'll be done with probably by the end of Wednesday's lecture all of the material that we covered in 51p. So it's really fast and so I'm glad that a lot of you were able to take that time to learn that, Joe, because it's super important. All right, so where we left off last time we were talking about ionic and covalent bonding. And we said that when you have an ionic bond you actually transfer electrons completely from one atom to another and then you'll get a positively charged atom and a negatively charged atom. Those two are going to be attracted to each other by a very, very super powerful ionic bond. And generally speaking, if the electronegativity difference is greater than about 1.8 the electrons are going to be transferred completely and you will have an ionic bond. There is one exception to this rule and that is the bond between H-hydrogen and fluorine, so hydrochloric acid HF. That is 1.8 difference in electronegativity. That is definitely a covalent compound. So that's kind of on the edge. And so that's why I wrote about, because it's not a super strict rule, we have a lot of variation in that. It's about 1.8. All right, we don't do too much of that here again because most of the things that you were talking about are covalent and bonded. And so covalently bonded compound add to take a filled valence shell by sharing electrons. So let's see what that looks like. Do that right at the very beginning. All right, so simple example here. A hydrogen plus a great orthoperiodic table combining with another hydrogen plus straight orthoperiodic table. And those two are going to, those electrons are going to overlap. And so we're going to have a hydrogen and hydrogen-hydrogen covalent bond. There's two shared electrons. Each hydrogen's happy because this hydrogen's happy because it's test two electrons in its valence shell for helium configuration. This hydrogen's happy because everyone's happy. And we can write that in a more simplified way. This is equivalent to HH. So that bond signifies a pair of electrons. So each hydrogen shares two electrons and again helium configuration. All right, so that would be between two like atoms. We can also have something between two unlike atoms. Hydrogen and chlorine. Let me make that panel a little bit bigger here. Those can overlap. Chlorine has seven electrons. So let's draw all seven electrons here. The hydrogen's going to overlap. The chlorine. Chlorine has eight electrons in its valence shell. Super happy about that. So there's eight electrons. That's the octet rule. Hydrogen has two electrons. That's the duet rule. And so we get hydrochloric acid. And again, this line between the hydrogen and chlorine represents a pair of electrons. So this is the type of bonding that we see with atoms of similar or the same electronegativities. As those electronegativities get closer. So what we're really talking about is thinking where the electronegativity difference between those two atoms is less than about 1.8. Let's look at how carbon hydrogen and oxygen and the halogens satisfy the octet rule. Four valence electrons. If you look at a periodic table, that's four valence electrons. To complete a shell, it needs four electrons. So it's going to have, in neutral carbon, there's going to be four bonds. So those bonds can be... We can have four single bonds. It doesn't really matter here. We can have four single bonds. We can have a double bond. That counts as two bonds and two single bonds. We can have a triple E bond into something else. Nitrogen has five valence electrons. So it needs three more. So it's going to make three bonds. So we're going to have three bonds in neutral nitrogen and a lone pair. Three single bonds. It could be a double bond and a single bond. It can be a triple bond. Oxygen. Six electrons in its valence shell. Satisfy the octet rule. And so it's going to make two bonds. We're going to have two bonds in neutral oxygen and two pairs of non-bonding electrons. And that could be two single bonds or a double bond. Hydrogen only needs one electron to have two in its valence shell. And so it's going to make one bond. And halogens also only need one electron. So they're going to make one bond. And really, 95% of what we deal with are these atoms. Very small amount of periodic table. Okay? So an easy way to remember this is the Haak 1, 2, 3, 4 rule. Hydrogen and halogens make one bond. Oxygen makes two bonds. Nitrogen makes three and carbon makes four. Okay? And that super important point I want to make here is that this is for neutral compounds. For neutral atoms. Okay? Uncharged carbon makes four bonds. Uncharged nitrogen makes three. Uncharged oxygen makes two. Hydrogen makes one. Halogens make one. Questions so far, anybody? Yes? Excuse me? Helium is the is the atom in the periodic table that is in the first row. And it has a filled valence shell, so two. It has two electrons. So hydrogen, when it gains an electron, it's going to have a helium configuration. All it is is the noble gas that's at the end of that row. Okay? More questions? Anybody? It gets a little bit trickier when we have charged atoms. We do have a lot of charged atoms. I mentioned last time that in GK, if you have a charged molecule, you put brackets and you put the charge on the outside of the brackets. We don't do that here. You want to know which atoms have the charge and you want all formal charges shown. So when carbon, hydrogen, nitrogen, and oxygen deviate from the whole world, they are sitting up with formal charge. I'm going to show you a couple of ways to figure out formal charge. It's something that you're going to need to know. You may just memorize this chart. I'm also going to show you a formula that you can use. You can use that or just kind of a way to think about it. But it's a super important concept and what I see is that there's people who are all the way in 51C and still get formal charges marked wrong and if you add up all the points it ends up being super significant. So it's something that you want to hold up right away. Okay, so carbon if you think about it, when we have neutral carbon it has four electrons. That's more of a periodic table. When it's cationic that means it's lost in electrons so now it only has three. So when it's cationic it's going to have three electrons and so we're going to get three bonds in a pair of non-bonding electrons. So that's what anionic is. We only have three bonds. It only has three electrons so it can make three shared bonds. When it's anionic that means it's going to have five electrons so that's going to make carbanion. We have a special name for it carbanion. And when it has five electrons that means it needs three to fill its valence shell so it's also going to make three bonds. So if you count the electrons here we have one, two, there's two for every bond two, four, six, eight. So notice that with a carbocation it only has six electrons so that's not a very happy compound. All the M's are trying to get an octet that's not a happy compound but that being said we are going to talk a lot about carbocation so you will see them all year. That's something we talk about a lot. Alright so nitrogen number of bonds to the neutral atom here is three number of bonds when managing this cation it's going to make four bonds it's going to need that to make an octet so three bonds here once we remove one of those electrons we're going to have to make an additional bond so positively charged nitrogen has four bonds it can be single, double, triple just as long as all of those add up to four while when nitrogen is anionic it only makes two bonds oxygen on the other hand when it's cationic makes three bonds it's going to look like that on three bonds some combination of making three bonds and oxygen that's negatively charged makes one bond you can memorize that when I see an oxygen and it has one bond and three pairs of non-bonding electrons it has a positive charge or when I have oxygen here that's positive when oxygen makes three bonds it has a pair of non-bonding electrons it's like that it's your choice I'm going to give you some other ways to do that it might be a little easier for you I'm going to do a formula some of you like to use formulas I don't personally like to use formulas too much because it's easy for me to forget them okay so how do we know this what I'm really trying to do here is can I understand this so I don't have to memorize can I understand this so you don't have to memorize it's another way you can do whatever you want so let's look at carbon for example so if we have neutral carbon we have four valence electrons so that would look like this four valence electrons so you can see that when we have carbon with four valence electrons it's going to want to make four shared bonds to make enough time you see the four dots there it's going to make four shared bonds so if we have carbon plus there's only three electrons because we've taken one of those away to get a positive charge and so carbon's going to look like this one, two, three and we put a positive charge it's going to make three shared bonds so this actually has a special name too that's a carbon cation it's so special that we have it has its own name carbon with a negative charge is a carbanion normal carbon would have four if it's a negative charge it has a classical electron so then it would look like this one, two, three, four and then we pair one of those so that means it's going to make it's going to bond with two, three atoms so it's going to bond it's going to it's bonding like a neutral nitrogen it bonds like neutral nitrogen neutral nitrogen makes three bonds this is also going to make three bonds so we would say it's isoelectronic with neutral nitrogen ok? questions on how I did that so far are you ready? let's do one more, let's do oxygen so neutral oxygen has six valence electrons right? so one, two, three four, five six so I took the four electrons I distributed them one on each side and then now we have two pairs of non-bonding electrons and it's going to make two shared bonds if we have O plus I won't do O minus I'm just going to do O plus because I don't have a verb to do O minus so we have O plus that means instead of six electrons we drop one we have five valence electrons like this one two, three, four five that's oxygen with a positive charge I did forget to put the negative charge here for part so we always want to put the charge near where that atom is ok? so oxygen with a positive charge bonds to three atoms so can you see that let's look at that so can you see that positively charged oxygen looks like this it's going to have the same bonding as negatively charged carbon ok? and that's going to have the same bonding as neutral nitrogen let's go back up to neutral nitrogen now it's going fast ok? so see neutral nitrogen three bonds and apparently non-bonding electrons ok? so the charged atoms are a really big deal we talk about them all in the time which is super important let's show you how you use a formula instead the formula is number of valence electrons minus half the number of bonding electrons minus the number of non-bonding electrons ok? so let's do an example of that on the next page so this could be a question like you might see on cypro charges to the atoms that are charging the following molecules so super important things if an atom has a zero charge don't put a zero next to it if an atom has a positive a plus one charge only put plus don't put plus one but these are some inventions we have in drawing so we're not going to put zero on anything we're going to put if we have a plus one charge we put a plus well so we're just talking about we're going to talk about that coming up so you're looking a little ahead we will talk about that coming up yes, if atoms have a lone pair they can make bonds to other species we're talking about just being able to draw a Lewis structure for a molecule but you're absolutely right we will get to that, it's going to take us a while alright, so the Honk rule is super helpful when we're doing a problem like this we do not have to calculate the formal charge for every single atom so for example by the Honk rule hydrogen makes three bonds so if it makes four bonds then that means that it is charged okay and so here's the formula but carbon here makes four bonds so that doesn't have a charge hydrogen makes one bond one bond, so we're not going to waste our time trying to calculate formal charge for the carbon this carbon also has four bonds by the Honk rule that means it's new comb oxygen has one bond so we're going to have two so it means it's charged after a while you're not going to need to use a formula you're just going to know that O- that's O- when it has only one bond it's just going to take a while as you work problems but let's calculate using the formula so we want the number of valence electrons for oxygen six minus one half of two a shared pair of electrons between oxygen and carbon so for the purposes of formal charge carbon owns one of those electrons that's all it owns, one electron for the purposes of formal charge and then we have six non-bond electrons so that's minus one so we'll put a minus there we want low pairs on this oxygen we'll talk about the mistructures coming up when you have a low structure you want to show low pairs and we'll also talk more about this coming up but in G-chem you had a little short cut for writing low pairs you could write a line for a pair of electrons we don't do that in O-chem so you have to take the time to draw two dots rather than a line the reason we have that rule is because when we start getting into mechanisms you're not going to know what's the head of an arrow what's the line, what's the charge it's not readable so that's why we have that alright so I'm not going to always go to a pair of electrons we have non-charging here we have nitrogen which is group five so that's five valence electrons nitrogen owns one electron for every program of bonds that it owns is four for one half of eight and then there's no non-bonding electrons so that's zero that's equal to plus one so nitrogen is a positive charge alright so a couple of things at point I want to make about this is first check 1, 2, 3, 4 rule first first check 1, 2, 3, 4 rule so you don't waste your time up relating zero formal charges don't write plus, don't write one if it's understood write zero if an atom has a zero formal charge anybody? questions to the octet rule we do have some exceptions to the octet rule most of the time we're going to be dealing with the second row of the periodic table carbon, nitrogen, oxygen hydrogen is in the first row carbon, nitrogen, oxygen and the halogens so we're most of the time going to be doing the octet rule 95% of the time but we do know that third row elements have D orbitals available from bottom so they can accept more than 8 electrons the two compounds are similar to this if you're going to see sulfuric acid and phosphoric acid okay so how many valence electrons is sulfur half here counting all the ones in its valence row 2, 4, 12 electrons all together 12 shared electrons it's a third row element carbon will never have 12 electrons around it carbon will only get a maximum of 8 that being said some of you are going to draw carbon with 5 bonds we have a name for that it's called Texas carbon um Texas for everything's very good but you don't want to draw that because carbon cannot exceed 8 electrons in its valence row easy mistake to make phosphorus has how many? all together in its valence row okay 4, 3, 4 elements there's also we also have molecules with open shells sometimes they're not enough electrons to provide an octet in GK we might remember this boron was one of the ones so boron like carbon can make 3 shared bonds and so one of the things that boron likes to do is make EF3 may or may not remember this from GK so what we have is boron here and you would have fluorines 3 shared bonds with fluorine and then we have fluorine here we have a shared bond there fluorines all have octets boron does not okay there's fortunately we don't have to draw it this way all the time we have a simplified way of drawing that we have boron with 3 fluorines each of those lines needs a pair of electrons but we do need to draw low pairs sometimes we leave low pairs off but when we're drawing a Lewis structure we always have to have them on so only 6 electrons so we call that open shell so you may say to yourself boron as an open shell can't you just make another bond with something that's going to go into a pair of electrons and then we'll have to take electrons that's exactly what boron does it likes to accept a lone pair of electrons but that's a chemical reaction between two compounds so a very common reaction which you may or may not remember from GK is boron trifluoride and so Lewis acid we'll talk about that in chapter 2 likes to accept a pair of electrons from something that has low pairs and so this is a common thing okay and what's going to happen and we'll talk about curvy arrows curvy arrows coming up but with curvy arrows we will see that nitrogen is going to donate a pair of electrons to boron and that's going to make everybody happy okay so that arrow notice I used a fissure arrow last time this is a double headed arrow okay so a double headed arrow means moving two electrons double headed arrow shows transfer of two electrons and that's a curvy arrow we use those a lot in GK so that's not something you've ever saw in GK okay so that's going to react and once you get a reaction there doesn't this going to look like you get a new covalent bond so let's label the new covalent bond we're also going to have some charges here so this is a new covalent bond we're supposed to have three bonds when it's neutral it means it's charged so we can use one of the three methods nitrogen with covalent you're going to say oh it has a plus charge I've done this so many times I know it has a plus charge you're not even going to think twice so one of the three methods you can use here I'm going to show you another method that you can use instead so nitrogen with covalent for the purpose of formal charge nitrogen owns one electron from each of those bonds right so nitrogen owns one electron from each of those bonds so it owns four electrons nitrogen's group five it only owns four electrons that means it's lost an electron it has a positive charge so that's another way to do it all I'm really doing is the exact same thing you do when you do that calculation using the formula the problem with the formula it's easy to make a mistake I saw one student one time a charge of 256 so I don't know how they got that but you can make a number that doesn't make sense alright let's do the same thing for boron well boron's owns one electron from every single bonding owns four electrons boron's group three it only has three electrons that means it's gained an extra electron so it has a negative charge so again whatever you find fastest it works better for you questions representation of structure we have a lot of different ways to represent structure molecular formula wow we don't use that too much I know you use that a lot in G cam we don't use that too much because it's not very helpful it's only helpful if you're trying to calculate the mass of that of that species that mass of that molecule and you'll use this in the lab so that way you know you say okay I have two carbons so that's 2 times 12.01 is the mass of carbon when you're calculating those calculates that's the only thing that's useful for it doesn't really tell us anything about how those are bonded together so it's not super useful for us then there's condensed structure condensed structure is a way this is the same molecule but it actually tells us a lot more about it that's absolutely the same molecule but now it tells us how things are bonded together and that's super important for us but there is a little bit of a learning curve in figuring out what this means so what this means is that we have a carbon that's bonded to three hydrogens and then it's bonded to another carbon which is bonded to one hydrogen and then that's bonded to an oxygen which is bonded to one hydrogen so that's what it tells us we can actually draw a structure where we would be hard pressed to draw one for the first one and notice this is the charge here the charge is on the atom that it belongs to tells you which atom is charged if I gave you the molecular formula you wouldn't be able to do that you wouldn't know which atom is charged so what this means is carbon's bonded to three hydrogens and then it's bonded to another carbon that carbon's bonded to a hydrogen because the hydrogen follows it and then it's bonded to an oxygen which is bonded to a hydrogen so what that would mean is there's actually more than one way to draw this again we will talk about that coming up but it says the charge is on the oxygen so that means that that means that the oxygen has three bonds the carbon's not charged so it has four bonds so by showing that cosmic charge oxygen it allows to draw the structure of that so oxygen with three bonds and a pair of non-bonding electrons has a positive charge so it gives a lot of information it also saves us a lot of time when we're drawing so there are some guidelines let's go through those guidelines here all atoms are drawn in but bond lines are generally admitted atoms are usually drawn next to the atoms which they are bonded so that's why we have CH3 means this CH parenthesis are used around similar groups bonded to the same atom small pairs are usually admitted formal charges are shown on the atom to which they belong super important let's do some examples when I'm looking at here I see CH3 that means I'm going to draw this okay I have a carbon and then I have CH3 two times bonded to that carbon and then I have a CH2 and a CL so once you can draw that so CH3 bonded to a carbon the parenthesis that shows that this carbon is bonded to two other menthols I'm going to call these bonds out very much longer because we have these shortcut ways of drawing things and then that carbon is bonded to a CH2 and then we have a torn lone pairs are optional in condensed structures so we don't have to put the lone pairs on the torn there's a couple of ones that are going to be a little trickier than others this next one let's have a little bit CH3 carbon bonded to hydrogen and also bonded to oxygen and so what that means is that this is what we have I'm going to go ahead and draw in those lone pairs it is not and as commonly a common mistake it is not this you know why we know it's not that and because of the way this is drawn the hydrogen is bonded to the carbon not the oxygen if it was this it would be CH3 COH it's not it's CH3 CHO which means that that hydrogen is bonded to the carbon so there's a couple of ones that are going to throw you off I'm going to show you a couple more that will throw you off let's label this first one wrong because bonding to carbon not oxygen how do we know if a hydrogen is bonded to carbon it follows this when we write thick and dense structure I'm giving you all the hardest ones CH3, CO, CH3 once a tricky one it's going to throw you off CH3 don't move on CH3 first one because I think that's going to throw you off I draw the second one to make it more clear I add more information if I can I draw this way and the reason I draw it that way is because if I don't draw it that way then 80 to 85% of the class will get it wrong on the test and that's not what I'm looking for I'm looking for you to get answers right I want to be able to convey what I'm what I'm trying to convey we believe that everybody in the room so it is not so these are the tricky ones that you just have to remember how do you know it's not this well once we start drawing Lewis structures we're going to come back to some of these ones and you're going to see that you're going to draw something that's really weird that doesn't make sense that has charges that it's not supposed to have it doesn't have the right amount of electrons okay so it's not that this one here is also a really, really common one for students off CH3 so probably everybody has a whole CH3 on so there's two ways to draw this I see this way in some textbooks I see this way in bio books I don't draw it that way I draw it this way and the reason I draw it this way is because those two oxygens are bonded to carbon and so it looks like this that's what it looks like it is not it is not this we know it's not those once we draw Lewis structures so that will be Wednesday's lecture we're going to come back and look at this and we're going to show why these two other structures are not correct we're going to show why this is correct these two are not correct but we need a little more information before we do that so you just have to remember that the CO2 in a condensed structure means carbon bonded to two oxygens one will have a double bond and one will have a single bond okay one will have a double bond one will have a single bond same thing with this last one here CH3 I'm pretty sure we're good on the CH3 here that's not a tricky one see how that double bond is drawn in the in that structure that's not required I always do it that is not required with that double bond here I always do it again to make it more clear so we have a carbon carbon double bond then we have our CO2 so the CO2 means the carbons bonded to two oxygens one with a double bond one with a single bond so I have a double bond here let's put long hairs here and one single bond so that's CO2 and then we have a CH3 so those are the trickiest ones questions on taking condensed structures and turning them into loose structures yeah it does yes and and you get used to kind of leaving long hairs off because you don't have to draw them in all types of structures so I'm just having to go back and work on it so thank you for coming anybody else? Questions? alright let's talk about loose structures that's what you really want to know about, right? structures we get connectivity that means how the atoms are bonded together tells us connectivity and the location of all non-bonding and bonding electrons as well as formal charges they do not give information about three dimensional orientation in space so you can see the way that I drew this here those are not writing on a molecule but it doesn't matter it's a loose structure we will draw on structures that show correct angles but it's not a loose structure so by the way, if I tell you to draw a loose structure on the test, I will remind you that I want to see all non-bonding electrons and formal charges I will tell you that so you won't forget so let's do the loose structure of the sample compound here we already drew on the previous page we are going to come back to the structure to talk a little bit more about it this is required what else is required? formal charge is always required low pairs are sometimes optional formal charge on atoms is always required no matter what type of structure you are drawing required required formal charge always required all types of structures not all pairs do have to have formal charges there is a loose structure then we have 3D structures here is where we attempt to give the 3 dimensional orientation of molecules that is chapter 5 we are not going to have to do this for a while but if we were going to draw the same compound showing correct angles we won't be able to finish this but we will start it how do we know what the correct angles are? we are going to talk about that after we finish the loose structure this here where a straight line is sponsored in the plane of the paper a wedge means out of the plane of the paper that means the hydrogen is coming out of the plane of the paper and a dash means that that atom is going behind the plane of the paper but we are going to stop right there sorry sorry sorry