Principles of Chemical Science/n * Email this page/nVideo Lectures - Lecture 16/nTopics covered: /nHybridization and Chemical Bonding/nInstructor: /nProf. Sylvia Ceyer/nTranscript - Lecture 16/nLast time we were talking about this valence bond model for describing the bonding./nIn particular, the bonding between atoms in a polyatomic molecule./nAnd we saw that the key to this description was allowing the wave functions, the individual wave functions on an atom to constructively and destructively interfere themselves to form some hybrid wave functions, to form some hybrid orbitals./nAnd then it those hybrid wave functions that overlap with the wave functions of another atom to form a chemical bond./nAnd, in particular, what we treated last time were these sp3 hybrid wave functions that were centered on carbon, that were centered on nitrogen, that were centered on oxygen./nAnd we saw that these sp3 wave functions were this linear combination, or this destructive and constructive interference between the s wave function and the three 2p wave functions./nTo form these sp3 wave functions and correspondingly these sp3 states./nAnd, as I said, we saw that on carbon, we saw that on nitrogen, we saw that on oxygen. And those wave functions end up being pointed to the corner or to the vertex of a tetrahedron where the atom was at the center of that tetrahedron./nIn the case of methane then, if you bond a hydrogen to each one of those sp3 hybrid wave functions, what you get is a tetrahedral geometry around the carbon./nIf you had a nitrogen with those sp3 wave functions around it and you bonded to three hydrogens then you got a trigonal pyramid as the geometry of the ammonia molecule with those two lone pairs sticking out there at another vertex of a tetrahedron./nBut remember we talked about the shape of the molecule as being described by the positions of the atoms and not the positions of the electrons./nAnd we talked about water. Water sp3 hybridized. We add two hydrogens to it./nWe form two bonds that are plane or bent. And then we have two lone pair electrons sticking out of the oxygen. Those are going to play a big role in some of the bonding we're going to look at, at the end of the lecture today./nWe have got to move on. And we have got to talk about another kind of hybridization, which we started to talk about last time, which is this sp2 hybridization. And we talked about it on boron./nWe saw that the sp2 wave functions end up being in a plane which, in the case of boron, if we then bonded hydrogens to those sp2 wave functions we got a planer molecule for boron H3./nBut carbon can also undergo this sp2 hybridization. And so let's take a look at that now. In the case of carbon, we have to do this electron promotion that we talked about also in the case of the sp3 hybridization./nAnd then we let now one of the 2s wave functions and two of the 2p wave functions hybridize constructively and destructively interfere./nAnd the result is three sp2 wave functions or three sp2 states with one electron in each one of them./nAnd then one of those p orbitals or those p wave functions on the carbon is untouched. That is just the atomic wave function, the atomic orbital on that carbon./nAnd if we look at the picture of all of those wave functions here it is. This is the 2s on the carbon, 2pz, 2px, 2py. And what we did is let these three, for example, constructively and destructively interfere to form these three sp2 hybrid orbitals or hybrid wave functions./nAnd the result is that they look like this./nThe result is they all have a large positive lobe on them due to the constructive interference and they all have a small negative lobe here. And then all of the rest of the pictures that we are going to draw, we are not going to draw that small negative lobe because it just makes it hard to draw./nBut the bottom line is that these three lobes lie in a plane./nAnd then, of course, we still had that 2py atomic wave function centered on the carbon. And that is unperturbed. Now what I am going to do is I am going to put all of these three wave functions on the same plot because they all have the same origin./nSo that is what I am doing right there./nThere it is for the carbon with the three positive lobes of that sp2 wave function all in the same plane. Again, I didn't draw the little part of the negative wave function in this picture./nBut now what we also have is that 2py wave function, the 2p wave function on the carbon that hasn't hybridized. Oh, I'm sorry./nThat hasn't hybridized, right? And it is sticking out of the plane of the board, right? That's the 2p orbital that you are used to seeing on the carbon./nAnd it is perpendicular here to the plane in which the sp2 wave functions are sitting. Now what we are going to do is I am going to take this sp2 hybridized carbon where the positive lobes of these sp2 wave functions are all 120 degrees away from each other./nAnd where you would say these wave functions had a trigonal plane or configuration./nI am going to take this sp2 wave function, and I am just going to rotate it so that this lobe here is parallel to the floor. I am going to do that for convenience./nAnd you are going to see why in a moment./nThat is what happens here on the next slide. I just rotated this by 90 degrees, in particular so I could put the z axis here along a bond I am going to form. That is the top view. Now, let's look at what this looks like from the side./nFrom the side what it would like is something like that./nThe carbon is right here at the center. Then the positive lobe, this positive part of the wave function, that is what this is here. This 2py wave function, well, here it is./nIt is perpendicular to this plane that you are looking down on up here. It is perpendicular to this plane. That is that 2py wave function. Positive part of the p wave function, negative part of the p wave function./nAnd then this part of the wave function, well, this is just a projection of these two big positive parts of the wave function./nThis just projects this way and out that way. That is what it looks like from the side./nNow what I am going to do is I am going to bring in another sp2 hybridized carbon up here. I am going to bring it in. It is going to look just like that. Here it comes./nAnd when I do that what is going to happen right here is that I am going to allow the 2sp2 wave function from this carbon to overlap with the 2sp2 wave function from this carbon and I am going to form a bond./nAnd that is going to be a sigma bond, sigma because it is going to be cylindrically symmetric around this carbon-carbon bond access./nAnd it is a sigma bond formed by the carbon 2sp2 hybrid wave function and the other carbon 2sp2 hybrid wave function./nI have formed a bond here. And now what I am going to do is bring this carbon sp2 hybridized wave function or carbon in again./nBut I am going to bring it in and we are going to watch this view. We are going to watch it from the side./nThis time what I want you to see is what is going to happen here to the 2py wave function. Here it goes. We are going to bring it in. And then this 2sp2 is going to form a bond here between this carbon, but now what happens is that those 2py wave functions overlap./nThey constructively or destructively interfere./nIf you want to watch it again, here it comes in, we are going to make this sigma bond, but now these two lobes are going to interfere and these two lobes are going to interfere. And this is going to be constructive interference there because they both have a positive sign or negative sign./nAnd so we are going to have wave function up here and wave function down there./nWe formed another bond right here. The bond that we formed here is going to be called the pi bond. It is pi because it is not cylindrically symmetric around this carbon-carbon axis. There is wave function or electron density if you square it up here and wave function or electron density if you square it down here./nBut right here along the plane there is a node in that pi wave function./nSo this is a pi wave function made out of the atomic wave functions on carbon. The atomic wave functions, the ones that haven't participated in this hybridization. It is made out of carbon 2py and carbon 2py./nWhat I have got here is a double bond./nI have two bonds between the two carbons. I have got a sigma bond and I have got a pi bond. That's great. Now let's bring in some hydrogen. I just did. And when I brought in those hydrogens, well, you can recognize this as ethylene./nThe hydrogens form a sigma bond between the carbon 2sp2 hybrid wave function and the hydrogen 1s wave function./nThat is what forms this sigma bond. It is sigma because it is symmetric around that carbon-hydrogen bond access./nAnd the molecule that we have here is ethylene. Now, let's look at that again because this is really very important. Here I show you just those two 2p2 hybridized carbons forming that sigma bond./nAnd this just represents here the energy levels as I bring them in./nIn this case we have two electrons in this sp2 state and we have one electron in each one of the 2py states. And so this represents this sigma bond here. And then we made a pi bond. And we will watch it now from the top./nWe made a pi bond./nWe let these two 2py wave functions overlap, constructively interfere to form a pi bond, which is the carbon 2py, carbon 2py. Clue here, every double bond has two bonds. That's why it is double. But every double bond is made up of one sigma bond./nAnd the sigma bond is always the overlap between two 2sp2 hybrid wave functions./nAnd the second bond that always makes up a double bond is a pi bond. The pi bond is always the overlap of two atomic wave functions. In the case here of carbon it is two 2p wave functions that have overlapped./nThat is what a double bond always is./nIt always is so whether it's a carbon double bond or a nitrogen double bond or an oxygen double bond. It is always one sigma, one pi. And there are the hydrogens. And, as we said, those hydrogens there, that carbon-hydrogen bond is this overlap between a carbon 2sp2 wave function and that hydrogen 1s wave function./nAnd now what is the geometry of ethylene? Well, the geometry here of ethylene is that the hydrogen-carbon-hydrogen bonds are 120 degrees./nThat is 120 degrees because that is the result of the hybridization, that 2sp2 hybridization of the wave functions around the carbon./nThis bond here is 120 degrees again because of the 2sp2 hybridization around that carbon./nSo the geometry of this molecule is planer. All of these atoms are in a plane. Those six atoms are in a plane and the bond angles here are all 120 degrees. Now, if you see a carbon that has a double bond./nIf you see a nitrogen that has a double bond./nIf you see an oxygen that has a double bond. That is always a clue that that carbon, that nitrogen, that oxygen has a hybridization of 2sp2. Double bonds have this hybridization always. If you are asked to tell what the hybridization is around some atom, it has a double bond to it, it is a 2sp2 hybrid./nNow what I want to do is I want to take this 2sp2 hybrid and build a larger molecule./nI am going to take the two carbons here that are bonded via this 2sp2 hydrogen./nSo here are two of them. And here are another two. And they are kind of strategically placed for what it is I want to do./nAnd, of course, this is the sigma bond between the two carbons. Here is the carbon 2sp2 sigma bond. Now what I am going to do is bring in another carbon that is sp2 hybridized./nHere it comes. And, again, it is strategically oriented so that right here I have now formed another sigma bond between the carbon 2sp2 wave function and the other carbon 2sp2 wave function./nThe same thing right here./nI just form two other sigma bonds. And now I am going to bring in a sixth carbon sp2 hybridized. And it is going to come in from this side. Again, it is strategically oriented./nI just formed here another sigma bond between these two carbons and another sigma bond between these two carbons./nHey, look at this. It is beginning to look a lot like benzene. [LAUGHTER]/nThat is my favorite part./n[APPLAUSE]/nHere come the hydrogens. This is going to be benzene. So what did I do here? I made some sigma bonds. I made six sigma bonds between the carbon 2sp2 wave functions and the hydrogen 1s wave functions./nNow I've got all of my atoms there and almost all of my bonds./nHowever, we have got to do something here about these 2p atomic wave functions that are centered on the carbon. We are not done yet. What is going to happen to them? Well, they are going to overlap constructively and destructively interfere./nAnd to see that, let's take this benzene molecule and rotate it 90 degrees so that we are going to look at it from the side on view instead of the top view./nSo there it is. Here are my carbon-carbon sigma bonds, the 2sp2-2sp2 bonds./nHere is my sigma bond between the hydrogen and the carbon 2sp2 wave function. And now I am going to let the 2p, the atomic wave functions, not the hybrid wave function on the carbon, I am going to let them constructively and destructively interfere./nAnd here it goes./nWell, they are going to overlap. And I've got something that looks like a pi bond here. I have wave function above the plane of these atoms and I have wave function below the plane of these atoms./nIt is a pi bond that is formed by the overlap of the 2p wave functions, the atomic wave functions on carbon./nIt is pi because it is not symmetric around now the bond axis. There is density up here, density up there, but not in the plane. Let's look at it again from the top view./nWhat did I do? Well, I let these 2p wave functions here overlap./nLet's look at that a little more carefully. What exactly did I do here? Well, what I did is, for example, let the 2p wave functions on these two carbons overlap to form a pi bond. I let the 2p wave functions on these two carbons overlap to form a pi bond./nAnd I let these two 2p wave functions overlap to form a pi bond./nWell, that is very nice but I could also have let the 2p wave functions on these two carbons overlap or these two carbons. In other words, I could have made a pi bond between these two carbons or these two carbons or these two carbons./nSo which one do I choose? Well, in recitation the other day you should have looked at the Lewis structure of molecular benzene or benzene./nAnd what you should have seen is that you would be able to write several different Lewis structures all which have the same set of formal charges./nAnd that you wouldn't be able to decide which structure you should have based on the formal charges./nRather, what you had was a resonance structure. And that is exactly what you have. Instead of these six extra electrons, which are centered on the carbon, those 2p electrons, there is one here, one there, one there, one there, one there, one there./nInstead of each one of those electrons being shared between just one of the adjacent carbons, those six extra electrons are actually delocalized around all of the carbons./nAnd so what you are going to form here is not a clear double bond between this carbon and this one and this one./nInstead, what you are going to form is kind of a half of a pi bond. That is you are going to let those six electrons be equally distributed, so to speak, around all of the six carbons./nAnd so this pi bond here isn't quite a full pi bond. It is kind of half of a pi bond./nBecause this carbon is sharing its 2p electron with this carbon and with this carbon and vice versa all the way around./nAnd so we have this resonance structure. I drew this as kind of a fuzzy green line there. We have these six pi electrons that are delocalized amongst the six atoms of this carbon right./nSo these double bonds here are really a bond and a half./nThis bond is more than a single bond in terms of its bond strength and in terms of its distance. It is closer than a single bond, but it isn't as strong as a double bond, nor is it as short as a double bond./nIt is somewhere in between./nThat is the structure here of benzene. So that takes care of sp2 hybridization. Now we've got one other kind of motif and that is called sp hybridization. And we are going to use carbon again as the example of this sp hybridization./nAgain, since we are starting with carbon here, we are going to have to undergo this electron promotion process./nWe are going to take an s electron. Are there some questions here that I can help you with? No. OK. We are going to take that s electron and promote it to the 2p state. And then we are going to do a hybridization, but this time what we are going to do is let the 2s atomic wave function hybridize with one of the 2p wave functions./nAnd the result then is that we are going to have a new wave function that we are going to call sp, and there is only going to be two of them because we only let one of the 2s and one of the 2p constructively or destructively interfere./nSo we are going to get two sp states or two new sp wave functions./nEach one has got an electron in it. Each state has got an electron in it. And then we've got left over two atomic states centered on the carbon, the py and the px. Each one has an electron in it./nYes? Well, by convention here yes./nThe p state we are going to hybridize is going to be 2pz, and that is because that is the wave function I am going to make a bond to. And I want that along the internuclear axis./nIn real life you cannot tell what is x, y or z, but by convention we are always going to put the z axis along the bond axis./nAnd I will say a little more about labeling py and pz and an example that I am going to do in a few minutes, so I hope that will clear things up./nIn the picture form, here are our three atomic wave functions again. And I am going to let these two constructively and destructively interfere./nAnd the result is then two new sp wave functions. Now, given that I am letting only two wave functions constructively and destructively interfere, it is a little bit easier to see how I get these shapes./nHere it goes. This 2s wave function, remember the 2s wave function always has one sign? Say it has a plus sign it never crosses the axis. There are no nodes, right?/nNo radial nodes. When you have a node that is when the wave function changes sign./nI am going to put that 2s wave function in the same place in space as this 2pz wave function. When I do that, since this is a wave, it is going to constructively and destructively interfere. Up here, where I am constructively interfering this positive wave function with the positive part of the 2pz wave function, I am going to get a lot of positive wave function./nThat is where this lobe comes from./nBut down here where this is positive and this is negative, that is going to be destructive interference. And because this negative part of the 2pz lobe is actually larger than the 2s that is positive, well, I am going to have a little bit of leftover of a negative wave function./nHere it is easier to see how you get this shape by the interference of that wave function with that wave function./nAnd then, correspondingly, I am going to let this wave function and that wave function destructively interfere. If this is still positive, positive minus a positive is going to give me a little bit of a wave function and it is going to be negative./nAnd then positive minus a negative, that is going to give me a positive. Here is my big positive part of the wave function./nThat is a little bit easier to see now than in the sp3 case where it is not so easy to see./nThese are now my two new 2p wave functions. And the atomic wave functions here, I haven't done anything to them. Now I am going to put these wave functions all on one plot. They all have the same origin./nIt is the carbon right in the center./nI am going to put them all on one plot. Here they are. And what I did is rotated the z axis up and down here. The z axis in this picture is coming this way./nIt is parallel to the floor. Again, strategically placed for the next bond formation. But here you can see the 2px wave function, the atomic wave function untouched. And perpendicular to it you can see the 2py wave function untouched./nNow I am going to take another sp carbon wave function and bring it in, and I am going to let the sp wave functions on the two carbons overlap so that I form a sigma bond./nSigma formed by the carbon 2sp, carbon 2sp. That's my sigma bond there./nBut now you know what is going to happen? We are going to let the atomic wave functions on the two carbons constructively and destructively interfere./nOh, I am going to bring in hydrogens first. Sorry. We brought in the hydrogen and formed this sigma bond between the carbon 2sp wave function and the hydrogen 1s wave function. And we've got acetylene./nExcept now we've got to let the 2p atomic wave functions on each carbon interfere./nExcept I wanted to point out this geometry. [LAUGHTER] I got this all wrong. This is 180 degrees because the sp wave functions lie in a line. The hydrogen-carbon-carbon-hydrogen here, that bond angle is 180 degrees./nWe have a linear molecule./nAnd now we are going to let the 2p wave functions interfere. When we did that we are going to form a pi bond. The interference of this 2py and this 2py is going to be a pi bond. It is pi because it is not cylindrically symmetric around this carbon-carbon bond axis./nThere is electron density above the plane of the slide and below the plane of the slide./nWe've got a pi bond here. And then, finally, we are going to let these two atomic wave functions constructively and destructively interfere, the 2px./nAgain, we are going to form a pi bond. It is pi because it is not cylindrically symmetric around the bond axis. There is electron density up here and electron density up there./nWhat we have here is a triple bond./nWe have, between the two carbons, one sigma bond, one pi bond and another pi bond. The two pi bonds are perpendicular to each other. That is important. So a triple bond is always composed of one sigma bond and two pi bonds./nThat is the case, whether you're looking at a triple bond on carbon or nitrogen or anything else, well, almost anything else that forms a triple bond./nBut for your intents and purposes carbon-nitrogen are going to form a triple bond./nThat is important here. Those are our hybridization schemes that we are going to look at./nAnd it is going to allow us to describe the bonding in lots and lots of molecules. And one of the molecules that I want to describe the bonding of is this molecule./nIt is methyl nitrate. This is going to help you out on the homework here./nSuppose you are asked to describe the bonding in methyl nitrate and you are given this structure. Well, the first thing you have to do is write down a skeletal structure./nActually, in your mind get down what the Lewis structure is for methyl nitrate. A couple of days ago I told you that when you see a CH3 species like that it is a methyl group./nThat is always terminal./nThat is a carbon with three hydrogens bond to it. And so that is what we are going to do. We are going to make this carbon there with the three hydrogens bound to it. And then, if you have a long molecule like this, one place to start is to then just bind this atom to the next atom./nSo I did that. I put the carbon bound to the oxygen./nAnd then following that rule, I took the oxygen and bound it to the nitrogen. And then following that rule, I got two oxygens here and I put the oxygen on this nitrogen and the oxygen on that oxygen./nAnd then I drew the Lewis structure. I counted my electrons, drew it up and there it is. There is a Lewis structure. But now what I am going to do is calculate the formal charge./nAll of these atoms here, if I do it right, have a zero formal charge./nI then find a formal charge of plus one on the oxygen and minus one on the nitrogen. And I see that the sum of the formal charges is zero. And the overall charge on that molecule is zero, so it looks like I did everything right./nHowever, is this the correct Lewis structure for methyl nitrate? No./nIt is not the correct structure for methyl nitrate. And it isn't because we have a negative formal charge on the nitrogen and a positive on the oxygen, and oxygen is more electronegative than nitrogen so this is not an adequate structure for methyl nitrate./nThis does not describe the chemical bonding in methyl nitrate. What does?/nWell, to do that let's look at the board here./nLet's look at the bonding here in methyl nitrate./nHere is another Lewis structure that I could draw./nCH3, oxygen, and then I could have put the double bond here between the oxygen and the nitrogen and then I could have bonded the two oxygens to that nitrogen./nAnd there is a bunch of lone pairs here on the oxygen. That's all OK. But now, if I go and calculate the formal charges on that, I am going to get a plus one on this oxygen./nI am going to get a plus one on that nitrogen and I am going to get a minus one on this oxygen and minus one on that oxygen./nThat is a lot of distribution of charge here away from the distribution of charge in the isolated atom. Let's see if we can do a little better. Again, starting with the methyl group here at the end, let's do this./nLet us bond the oxygen to the carbon and then singly bond that nitrogen to the oxygen and then doubly bond one oxygen and then singly bond the other oxygen./nAnd there is going to be a resonance structure to this because I could have switched where the double bonds were here./nAnd if I go and I calculate the formal charges, I find now everything has got a zero, except the nitrogen has a plus one and this oxygen has a minus one./nAnd that is a much better situation than a lower energy situation than a plus one on the oxygen and a minus one on the nitrogen. There is the resonance structure. I won't draw it. But that is the Lewis structure here for methyl nitrate./nNow, clue./nIf you see an NO2 written like that, that is always going to mean the two oxygens are bound to the nitrogen. And if you ever see the word nitrate like in that expression, that means all of those oxygens are going to be bound to the nitrogen in some way./nOn an exam, if you see an NO2, those two oxygens are going to be bound to the nitrogen./nIf you see the word nitrate, the three oxygens are going to be bound to the nitrogen. Yes?/nIf I do that then I am going to have these five bonds to the nitrogen, and that is going to give a very high formal charge./nIt is not going to be the lowest energy structure./nThat is our structure. But now the question is how to describe the bonding in this molecule. Let's do that. Let's describe that bonding. Let me erase this structure here./nWhat do we have? We have three carbon-hydrogen bonds./nAnd this carbon here is bonded to four different atoms. That carbon has no double bond, has no triple bond. If it is bound to four different atoms and it has just single bonds, the carbon is always sp3 hybridized./nThat is something to know./nWhat that means is that these carbon-hydrogen bonds are sigma bonds. They are sigma bonds formed between the carbon and the 2sp3 wave function and the hydrogen 1s wave function./nThat describes the bonding in the carbon-hydrogen bonds. What about this bond, the carbon-oxygen bond here?/nHow do we describe it? Well, we already said this carbon is sp3 hybridized. And so we can write this carbon as sp3./nAnd then this oxygen here, well, this oxygen has got two bonds to it. It has no double bonds, no triple bonds, so this oxygen is also sp3 hybridized./nThis oxygen looks like the oxygen in water. This is going to be a sigma bond./nIt is going to be symmetric around that axis between carbon 2sp3 and oxygen 2sp3. We've got that bond. Next we have an oxygen-nitrogen bond. How do we describe the oxygen-nitrogen bond? Well, the oxygen we already said was 2sp3 hybridized./nWe know that./nBut this nitrogen has got a double bond to it. If it has a double bond to it, well, that nitrogen is 2sp2 hybridized. But this is still a single bond between the oxygen and the nitrogen there. It is symmetric./nAnd so that bond is a sigma between the nitrogen 2sp2 wave function and the oxygen 2sp3./nNext the nitrogen-oxygen double bond, how are we going to describe that? Well, it is a double bond. It means we have two bonds./nOne of those bonds is a sigma bond, as I said. It is going to be a sigma bond between the nitrogen 2sp2 wave function and the oxygen 2sp2 wave function./nBecause oxygen here is also double bonded. If we use that rule, we will consider the oxygen to be 2sp2./nIt is a sigma bond. But now we have two bonds, so the second one has to be a double bond. What is that double bond going to be? Well, it is going to be between the nitrogen 2p wave function./nBecause pi bonds are always between the unhybridized wave functions./nNotice I am going to leave out the x or the y here because I don't really know. And between the oxygen 2p wave function, I left out the x and the y there. Then, finally, I have the nitrogen-oxygen bond, that second nitrogen-oxygen bond./nAnd what do I have there? Well, it's a single bond./nIt is going to be sigma. It is symmetric around the axis. It is going to be between the nitrogen 2sp2 wave function and the oxygen 2pz wave function. It is going to be 2pz in that case because we are here along the bond axis./nIt is sigma bond. It is not a pi bond./nThe sigma bonds are along the bond axis. So that is methyl nitrate. I am a little bit rushing to do something here, and that is I want to describe to you one other kind of bonding./nBut this is a bonding that is not a bonding within the molecule./nIt is actually bonding between molecules. And this is called hydrogen bonding. In hydrogen bonding, what you have to have is a hydrogen, and you have to have a hydrogen bonded to an electronegative element like oxygen./nA good example is water./nWater is bound to an electronegative element, and we've got two lone pairs. But in hydrogen bonding you've got bonds between molecules. This is intermolecular bonding. And what happens is this hydrogen here will interact very strongly with the lone pair electrons here on its neighboring oxygen forming a bond./nIt does so because this hydrogen is kind of deshielded./nThe oxygen being so electronegative it has kind of pulled all the charge toward it. This hydrogen kind of looks a little bit positively charged, delta plus where this is kind of delta minus. And so this hydrogen being really small, it can get close to the oxygen lone pairs where there is a lot of electron density in form of a bond./nAnd that bond ranges from 20 to 60 kilojoules per mole./nCertainly not as strong as a covalent bond. A C-H bond is 400 kilojoules per mole, about the tenth of that strength. And it turns out to be very important. Hydrogen bonds determine the properties of water./nThey determine the structure of proteins. DNA owes its helical structure to hydrogen bonding./nTrees stand up straight because of hydrogen bonding in the cellulose chains. The strength of nylon is determined by hydrogen bonding./nAnd whether or not you have a bad hair day is determined by hydrogen bonding. For example, the slides on the side, what you see is the protein of hair. This is a polypeptide./nIt is a repeating unit here of C-O, C-H, N-H, C-O, C-H, N-H again and again and again./nHere is another polypeptide or another peptide strain with this repeating unit. Well, what happens, when your hair is actually wet is the hydrogens on the nitrogen, you can see in that chain up there./nWhat happens is that they are hydrogen bonded to water molecules./nAnd you can also see in the second chain over that the oxygen on the carbon here is hydrogen bonded to a water molecule. When your hair is wet the actual polymers in your hair kind of slip by each other./nHowever, if you then take your hair and you put it some very strange configuration and you dry it, you remove these water molecules that are hydrogen bonded to your hair, what happens then is that these two chains now kind of lock into registry./nThey kind of bond together because these water molecules are gone here./nAnd so now they hydrogen bond to each other. The hydrogen bonds to the oxygen on the CO and the two chains are in registry. And then when you let your form go, well, your strands actually stay in that registry for a little bit until the humidity wipes that out./nBad hair day. See on Friday.
Channels: Chemistry (General)
Tags: Hybridization Chemical Bonding
Uploaded by: mitlectures ( Send Message ) on 16-04-2009.
Duration: 51m 30s