hi welcome to our general chemistry lecture part 2. let's get some screen sharing going okay so for our general chemistry review we're in part two now and what we're going to do is we are going to discuss three different theories that help explain bonding and the first is molecular orbital theory mo theory the second is valence bond theory and the third is vesper theory so molecular orbital theory in organic chemistry we're not going to use that much so i'm just going to show you the basic slide of it we do use valence bond theory quite a bit so we'll go into depth and balance bond theory and then vesper theory is the theory that we used when we did lewis dot structures in general chemistry so let's get started so mo theory okay the idea of mo theory is that so far we've talked about these atomic orbitals we've said there's an orbitals that belong to each atom the 1s the 2s so these these individual orbitals but mo theory goes beyond that and it it creates a a molecular orbital so for a whole entire molecule so for instance a molecule of h2 okay h2 is two hydrogen atoms put together so we'll call that a molecule and instead of using two individual atomic orbitals they come basically create one molecular orbital to encompass the whole entire molecule so this is a more complete way to describe bonding because it includes constructive and destructive interference which has to do with the wave theory um how electrons can act like a wave or like a particle and but mo theory uses the the linear combination of atomic orbitals in a mathematical method to make these new molecular orbitals so it's not easy for us to do at a simple level we can only look at the most basic molecules like hydrogen or helium and so um it's it's more of something we would do with a computer so anyway here is a molecular orbital diagram for hydrogen and basically we have one electron for each hydrogen atom that's in the 1s orbital and when those come together we get a bonding molecular orbital or an anti-bonding molecular orbital and these two electrons come down into the bonding ammo and so that makes for a positive a good outcome a more stable outcome and therefore hydrogen is a molecule that does occur helium however is not diatomic so it doesn't exist as he2 and that's because if we draw let's see if we have one we don't have one but if we draw the mo molecular orbital diagram for helium we would have electrons in the bonding molecular orbital but we'd also have them in the anti-bonding molecular orbital and you can see how that's higher in energy than the original atomic orbitals so that that tells us that the he2 does not want to exist okay you might remember in 1a we calculated bond orders from these molecular orbital theories and we could decipher if a molecule is stable or not based on its bond order so that's all we want to say about mo theory for now we will see it a tiny bit i think it's actually really in 128b maybe where we talk about it maybe we do see it i think a little bit in chapter six seven five maybe five no six six in this chapter anyway okay so now we're going to talk about valence bond theory okay so remember that we say electrons can act like waves and act like particles okay so waves can have positive and negative portions to them okay so when two waves come together they can either be amplified that means get bigger or like you've seen on an ocean they can come together and they can cancel each other out and the waves will go flat so when the wave gets bigger okay the waves reinforce each other we call that constructive interference when the waves get smaller and flatten each other out we call that destructive interference so constructive interference is needed for a bond to occur basically two orbitals come together in a reinforcing way to form a bond and this is our kind of schematic here okay so we've got a 1s orbital and a 1s orbital each of these has one electron in it for each of the hydrogens the orbitals come together in a constructive way and therefore we have a bond that forms so the electrons are being shared between the two atoms okay so the idea of hybridized atomic orbitals in valence bond theory um let's look at um carbon okay so um carbon okay uh let's see here let's just look at the next one here okay so if we have um a carbon well we i'm sorry let's start so we drew the electronic configuration for carbon earlier okay it had two carbons in the 1s i mean 2 electrons in the 1s 2 electrons in the 2s and then one electron in each of the p okay if we were to take this carbon and we were to make bonds with hydrogen okay that means the hydrogen would share an electron with the carbon like the carbon would share an electron with the hydrogen okay and that means this carbon could only make two bonds because it has only two unpaired electrons okay remember the electrons like to be paired up so a hydrogen could pair with one of the electrons in the carbon and the other electrons in the carbon so this carbon could only make two bonds okay but we know that's not what carbon does for instance methane this is natural gas is ch4 it has four bonds to hydrogen so how do we explain that according to our electron configuration carbon would only form two bonds but we know that carbon forms four bonds well the carbon must undergo hybridization of its atomic orbitals to form four bonds so what that means is the s orbitals of the carbon and the p orbitals of the carbon mix okay so we've got px py and pz they hybridize they mix to form four degenerate sp3 orbitals so the s and the p okay there's one s and there's one two three piece one s three piece see that the 1s and the three ps come together to form sp3 orbitals four orbitals combine to form four sp3 orbitals these are degenerate sp3 orbitals they're all of the same energy so what's happening then is the electron from the 2s orbital is mixing with the 2b orbitals to form an intermediate in energy sp3 orbitals okay so now each of these sp3 orbitals can go on to form a bond with a hydrogen right our carbon forms four bonds with hydrogen so the one of the s electron mixes with the p electrons to form four degenerate sp3 orbitals the sp3 orbitals look like this okay they're kind of a combination between the p and the s orbitals they have one large lobe and one small lobe and a node in the middle so here's how we draw this okay we do a lot of schematic drawing drawing in valence bond theory we take our carbon and we draw four sp3 orbitals okay on the back of each of these is the little lobe as well they're not showing those little lobes okay so you've got big lobe little big lobe little lobe big lobe little lobe big lobe little okay but they're not showing the little lobes each of those sp3 orbitals can then bond with a 1s hydrogen orbital and the electrons are shared so if we did not hybridize the orbitals we would have no way to explain how methane forms this compound that it does okay so our answer is to mix the s and the p orbitals into four degenerate sp3 orbitals so what i want you to notice okay here's the let's let's get a pen and color magenta what i want you to notice is sp3 carbon when you see carbon and it has all single bonds attached to it you see that all single bonds attached to it all single bonds attached to it you call this carbon sp3 hybridized okay we're going to use that as a common term later on what's the hybridization of that carbon it's sp3 okay that means it's in all single bonds it means that it has these sp3 orbitals now let's look at a carbon in a double bond okay a carbon in a double bond is what we call sp2 hybridized this carbon is s sp2 hybridized that's because the s orbital is mixed with two of the p orbitals to form three degenerate sp2 orbitals okay so we've got three things mixing it forms three orbitals let's see if we have another picture here okay so here's the carbon it was this carbon right here i'll show you what it looks like in the drawing so let's go back up here we've mixed an s with two p orbitals to form three sp3 orbitals okay there's one p orbital that's just hanging out still okay that p orbital is what's used to form the pi bond the pi bond is the double bond the rest of these orbitals i mean the rest the other bond here so you can consider let's say this top bond the pi bond the other three orbitals are used to form sigma bonds okay there's a sigma bond here a sigma bond here and a sigma bond between the carbons so these carbons are forming each of them are forming three sigma bonds and one pi bond the sp2 orbitals are forming the three sigma bonds and the p orbital is forming the pi bond so it looks like this in our balloon model here here's the carbon okay so this carbon here is this carbon down here that carbon has a p orbital okay the blue and the red or the p orbital and it has one two three the gray sp3 orbitals the sp3 orbitals are overlapping with the hydrogens and with the sp did i say sp3 i'm sorry the sp2 orbitals are overlapping with the hydrogens and the sp2 orbital over on this side is overlapping with the other carbons sp2 orbital these are all called sigma bonds sigma sigma sigma here's the full drawing so i'll try it out for you okay here's the two carbons okay the carbons are making a sigma bond with each other and a sigma bond with the hydrogens each of the carbons is doing that it's using its sp2 orbitals sp2 each of the carbon has three of them remember so the gray is all the sp2 orbitals each carbon has three of them sp2 orbitals are overlapping with the hydrogens in sigma bonds the sp2 orbitals are overlapping with each other carbon to carbon in a sigma bond okay sigma bond means overhead i mean over um sorry end to end overlap okay so they're actually overlapping each other to form a bond sigma bonds are the strongest kind of bonds okay these p orbitals remember there was p orbitals left over each carbon had a p orbital left over the p orbitals form through space bonds that we call pi bonds probably on the next slide okay end to end overlap through space this is a much weaker bond than the sigma bond okay because they're not right on top of each other so each of these carbons has three sigma bonds and one pi bond we're going to when we talk later we're going to see that sigma bonds don't break okay they they're high in energy they don't they don't or they're they're they're um less reactive so they're not going to break the only reactions we're going to see sigma bonds do is when we get to one chapter on radical reactions so all the different reactions that we see are going to be between charged items or these double bonds breaking these pi bonds breaking okay something going on with a pi bond acetylene has one sigma bond and two pi bonds the sigma bonds in acetylene come from sp hybridized orbitals the s mixes with one of the peaks and then there's two p's left over the two p's that are left over form the pi bonds so this is for each carbon right each carbon has two p's left over that can form the pi bonds and these sp orbitals then go to form the sigma bond between the carbons and the sigma bond between the carbons and the hydrogens so four carbon okay four carbon when you see carbon in a single bond only based on all the stuff we just looked at you're going to say that carbon is sp3 hybridized excuse me when you see carbon in a double bond then you're going to say that that carbon is sp2 hybridized when you see carbon in a triple bond you're going to say that carbon is sp hybridized okay so let's practice that here's ibuprofen this is a line drawing of ibuprofen we're going to learn about line drawings once we start organic chemistry here's the actual structure of ibuprofen all drawn out with all its hydrogens let's just start at the top here's a carbon it has one two three four bonds to it so what's its hybridization all single bonds we would call that sp3 sp3 hybridized carbon this carbon oops i didn't draw it out here there's a carbon right here this carbon is in a double bond and single bonds if it's in a double bond we're going to call it sp2 hybridized carbon in a double bond carbon in a double bond carbon and a double bound all these carbons in a double bond we're going to call them sp2 hybridized so they can also have a single bond but if they're in a double bond we call them sp2 hybridized this one one two three four bonds to that carbon so it's sp3 hybridized all of these carbons here are sp3 hybridized okay how about each of these carbons we would call them sp hybridized because they're in a triple bond this carbon is also the second carbon sp oops erase that scratch that up this carbon here is also sp hybridized the last carbon has four bonds to it i didn't draw it out for you so i'll do that let's see if i can eraser oh good there we go i'll draw it out for you so this last carbon has four single bonds to it so we would say that that carbon is sp3 hybridized okay so this is your cheat sheet of all of valence pawn theory that is your cheat sheet right there even if you didn't get any of the drawing things that's your cheat sheet for valence bond theory okay the other thing i just want to tell you is that when we look at bond lengths the single bond is the longest bond and as we add a double bond to a triple bond the triple bond is the shortest okay so the triple bond is the shortest and because those carbons are so close to each other it's a strong bond but it's also reactive because the pi bonds break easily okay the sigma bond holds tight but the pi bonds break easily so longest is a single bond shortest is the triple bond okay the last thing we want to look at is vesper theory okay and you did this in 1a the book does this a little bit different vesper theory is valence shell electron pair repulsion theory right it's just the idea that we draw that lewis dot structure and then the electrons try to get as far away from each other as possible and so therefore we can detect predict the electron and molecular geometry of the of the molecule based on the electrons getting as far apart from each other as possible now in this kind book it does things a little bit different by calculating what it calls a steric number okay a steric number is just a number another term for number of electron groups so it's the electron groups around the central atom do you remember doing that in 1a we would count the electron groups around the central atom so for instance this carbon has one two three four electron groups around the central atom so we say its steric number is four this nitrogen has one two three four electron groups around it that means the steric number is four and again the oxygen has one two three four electron groups around it so we say its steric number is four okay so we use the steric number to determine the electron geometry it's the number of electron groups around the central atom here are the different electron geometries that we learned okay the different shapes linear an example was carbon dioxide trigonal planar bent tetrahedral it looks like a tetrahedron we've got barking dog all right so sp3 geometry a molecule that we call sp3 when we say a carbon is sp3 we say it has a tetrahedral electron geometry okay so methane has a nice tetrahedral electron geometry there is a equal distance between each of the bonds okay ammonia however has a tetrahedral electron geometry as well but it's a little bit different because it now in place of one of the actual sigma bonds it has a lone pair and then for water it's also called tetrahedral but it's even a little more different because in place of two of the bonds it has lone pairs so what that does is these lone pairs push these their electrons right so remember the electrons are repelling each other so the electrons are pushing the other bonds closer together so these hydrogens here have about a 105 angle while these hydrogens on the ammonia have a 107 angle and the hydrogens on the methane have a perfect 109.5 degrees so the methane is ideally tetrahedral so the ammonia and the water we now use molecular geometry to describe them okay in molecular geometry remember we need to look at the bonding versus the lone pairs so methane has all bonding pairs ammonia has three bonding pairs and one lone pair water has two bonding pairs and two lone pairs so that makes the molecular geometry of these molecules different the molecular geometry of methane is tetrahedral the molecular geometry of ammonia is trigonal pyramidal and the molecular geometry of water is bent okay so we just said if the hybridization is sp3 then the molecular geometry is i'm sorry the um electron geometry is tetrahedral okay we can also use the steric number so vesper theory to predict the hybridization okay so again if the steric number of methane is four the number of electron groups around the central atom then we say the molecule is sp3 so i mean the atom is sp3 so this carbon is sp3 hybridized then nitrogen is sp3 hybridized and the oxygen is sp3 hybridized okay in this molecule which was ethene those are h's on the end i know they're not drawn very well okay this carbon has one two three electron groups around it remember a double bond counts as one electron group a triple bond counts as one electron group so this carbon only has three electron groups around it its steric number is three so what did we say this carbon was sp2 hybridized okay they're just kind of showing you there's different ways to come to the same answer okay i told you to cheat i just look at the carbon and like it's in a double bond so it's sp2 hybridized okay they're counting steric numbers to do it so okay so um here's a problem it says analyze the steric number hybridization electron group geometry and molecular geometry for this mean okay this is a a functional group we'll learn functional groups when we start organic chemistry this is called an amine and this nitrogen has how many electron groups around it one two three a lone pair a single bond and a double bond they all just count as one electron group okay each one is one electron group so i have three electron groups around the central atom if i have three electron groups that means the steric number is three so that means that this nitrogen is sp2 hybridized right the nitrogen's in a double bond it's sp2 hybridized right i don't even need to do stair groups to figure that out as long as it's in a double bond it's sp2 hybridized okay the electron geometry then with three electron groups around it we can go back to our old chart for this one okay three electron groups that means that the electron geometry um three electron groups over here sorry three electron groups the electron geometry is trigonal planar now depending on if they're all bonding or if there's in our molecule there's two bonding and one lone pair therefore we're going to call that molecular geometry bent gosh okay so the electron geometry is trigonal planar and then we said because there's two bonding pairs and one non-demanding pair we'll call the molecular geometry bit okay so this is the summary that's in your client book again this is the same stuff you learned in 1a they're just changing it up a little by introducing this thing called the steric number okay so if there's four electron groups around the central atom we call it sp its steric number is four we call it sp3 and we say that the molecule is tetrahedral okay in electron geometry if it has one lone pair we'll call it trigonal pyramidal it has no lone pairs its molecular geometry is tetrahedral as well okay two lone pairs we call it bent if the steric number is three it's sp2 hybridized that central atom is sp2 hybridized we'll call it trigonal planar again one lone pair versus no lone pairs and steric number two we call it sp hybridized we call that molecule linear and that's all it can be is linear okay i think we should stop here actually let's just finish this off okay so we're going to move you can always take a break if you want so we're going to move now to look at molecular polarity okay again i said that these tiny little pluses and minuses in organic chemistry are what are going to cause these molecules to react with each other remember that opposite charges attract to each other it's like if you've ever heard that that term opposites attract right so if you have a positive charge and a negative charge they'll attract to one another it's like taking two magnets and they come together okay so um and that's what's going to happen with our molecules in organic chemistry a negative charge is going to come attack a positive charge and everything's going to be caused by these tiny little charges so um electronegativity differences cause induction induction is simply a moving of electrons in their orbitals okay this repo results in a dipole moment a dipole moment as we said is simply a partial positive charge and a partial negative charge on a molecule so electrons are more pulled one way than the other so different bonds have different dipole moments they are so for instance and this has to do with the electronegativity of the atoms right an oxygen is more electronegative than a carbon it has a dipole moment of 0.7 okay again you don't need to look at these calculations you don't need to memorize any of these numbers it's just knowing that there is polarity associated with these bonds okay and h is a polar bond a c o is a polar bond these have dipole moments this carbon is partial positive charge this oxygen has a partial negative charge the atom that's more electronegative will have more of the electron density okay and this molec in this molecule here the o is more electronegative than the hydrogen so it's going to have the partial negative charge associated with it