 Today, we will continue our discussions on metal carbonyl compounds. As I had mentioned earlier, metal carbonyls are one of the most important class of compounds in organometallic chemistry. That is because carbon monoxide is like the R group in organic chemistry. Whenever you want a group to be placed in a position on the metal where there is some unsaturation, the carbon monoxide can be used just like the R group can be used anywhere in organic chemistry to plug in an unsaturated position. So, today we will talk about some bonding aspects in metal carbonyl complexes. This is extremely important because as a archetypical ligand, carbon monoxide typifies several interactions that are present in organometallic compounds. Once we understand the interactions in organometallic compounds containing carbonyl groups, most of the interactions can be understood very easily. So, let us take a look at the bonding in metal carbonyl complexes. We should look at some of the key features which characterize metal carbonyl compounds. Since the chemistry is extensive, we have to split this into two lectures. In the first lecture, we covered the fact that most metal carbonyl compounds are in fact typically 18 electron systems. So, they tend to be 18 electron complexes and they also are characterized by the fact that they are metals in 0 oxidation states. When they are in 0 oxidation state or in the negative oxidation state, they apparently do not have this electrostatic component. They must be having a strong covalent interaction to stabilize the system. We also noted that there are two other factors which distinguish them from Werner complexes. These two factors, one is the fact that carbon monoxide can bridge. Many metal centers very easily and the bridging form and the terminal form are almost iso energetic, which means that they can shift from one to the other very easily. It is also interesting that metal metal bonds are facile even when there are no bridging ligands. So, organometallic compounds are distinguished especially carbon monoxide based compounds are distinguished by these strange factors. In the last lecture, we covered the fact that both spectroscopy chemistry and structural features in the metal carbonyl chemistry is unique. So, MCO or metal carbonyl compounds are unique because they have this multiple bonding or the hint of a multiple bond between metal and carbon monoxide. Any bonding model that we propose should account for all these above factors. All the factors that we have looked at in the previous lecture namely in terms of the bonding aspects have to explain the fact that they are neutral systems that they are metal. They can contain metal metal bonds. They have spectroscopic features which are unique and they have structural features which indicate the formation of a double bond. So, let us proceed further. Let us look at just revise some of the structural parameters that characterized the metal carbonyl compound. If one looks at the iron carbon bond in an iron carbonyl complex, we noted the fact that the bond distance should have been approximately 2.10 angstroms. This is the expected value. This is the expected value and what we observe? This is the observed value. So, very clearly the observed value is much shorter either it is 1.833 or it is 1.810. These are two observed values in the complex FECO 5 and we have two observed values because we have an unsymmetrical complex. We have a trigonal bipyramidal system where we have equatorial bonds and we have axial bonds. Surprisingly, the axial bond is a one that is shorter. This is 1.81 and this is 1.83 angstroms. So, this is a very interesting observation. The fact that you have shorter bonds and the fact that you have shorter axial bonds compared to equatorial bonds. If you look at the carbon oxygen bond distance, the carbon oxygen bond distance should have been 1.128 if it is unaffected from the starting bond distance observed in neutral carbon monoxide. So, neutral carbon monoxide has a bond distance of 1.13 angstroms. We can ignore this third decimal place. Let us just round it off to 1.13 angstroms. In the complexes, it usually ranges between 1.14 and 1.16. In other words, this bond distance has been elongated and this bond distance has been shortened. This is what we suspect is coming from the fact that you have multiple bonds between iron and carbon and you have a weakening of the carbon oxygen bond distance. So, this is also reflected in the spectroscopic features of metal carbonyl complexes. In metal carbonyl compounds, the stretching frequency ranges from 1850 centimeter minus 1 from 1850 centimeter minus 1 to 2150 centimeter minus 1. That is approximately 100 to almost 300 centimeter minus 1 lower than what you have for the free carbon monoxide. If you have this reduction in stretching frequency, it is indicative of the fact that it is easier to stretch carbon monoxide in the complex compared to stretching carbon monoxide in the free state. So, clearly the bond order between carbon and oxygen has been reduced and this is identical to the result that we observed from the bond distances. So, these two factors point to a single event and that is the reduction in the stretching frequency reduction in the bond order between carbon and oxygen. So, in general, this reduction in the bond order between carbon and oxygen has been explained using what we call as back bonding, native bonding or the DCD model of bonding. DCD comes from Duer, Chat and Duncanson. The three people who extensively talked about the multiple bond between the metal and the carbon and the reduction in the bond order in multiply bonded systems which are bonded to the metal. In this case, carbon monoxide, it initially has a formal bond order of 3, the bond order between carbon and oxygen. The bond order between carbon and oxygen is actually 3 to start with in free carbon monoxide. So, this bond order is reduced when you have coordination to the metal. As we have seen in many complexes, it is the carbon end which is donating the pair of electrons and it is this pair of electrons on the carbon which is being donated to the metal. In the reverse fashion, in the reverse direction, we have donation of electron density from the metal into the carbon monoxide pi star orbital. So, this is a phenomenon, the second phenomenon. The first phenomenon is very similar to what we observe in Werner complexes. In Wernerian complexes, we would expect donation of a pair of electrons from the ligand to the metal. This is exactly what is happening in the case of carbonyl complexes also. But, in addition to that, we have donation of a pair of electrons from the metal to the carbon monoxide pi star orbitals. In the previous lecture, we just briefly reviewed the orbitals of free carbon monoxide. Let us just take a quick look at that once again before we proceed further. Let us look at the sigma orbitals in carbon monoxide first. We have a very low lying sigma orbital on carbon monoxide and a second and a third sigma orbitals which are in fact the more interesting ones. The carbon monoxide orbital which is the most important one, the sigma orbital which is the most important one is this one which is the highest occupied molecule orbital. This is the highest occupied molecule orbital on carbon monoxide and it is primarily concentrated. The lobes are primarily on the carbon side. So, you have this is the carbon side and this is the oxygen side. You can see that the carbon side has got a very large lobe and this lobe is now pointed towards the metal atom. So, this is the Homo which means the electrons will flow out from the Homo onto the metal. So, this is the primary interaction that we are talking about when you have a lone pair on carbon and an oxygen. But, the lone pair on carbon is more easily donated and that is pumped into the metal vacant orbitals of the metal. So, that is lying at minus 13.02 electron volts that is indicated here. So, the highest occupied molecule orbital is the one which donates a pair of electrons to the metal. Let us now look at the pi orbitals. The pi orbitals are two of them are there which are pictured here. There are totally four and we will come to that in a moment. The pi orbital or the bonding molecule orbital is the one which has got more concentration from oxygen and less from carbon. The pi star orbital which is the Lumo, this is the Lumo lowest unoccupied molecule orbital and the Lumo has got a large contribution on carbon and a smaller contribution on oxygen. So, you can see that very easily in this picture here. This is a large contribution on the carbon and a small contribution on the oxygen. Whereas, the pi or the bonding molecule orbital has got more contribution from the oxygen side. So, this particular combination of the donor orbital being concentrated on carbon and the donor the acceptor orbital on carbon monoxide concentrated on carbon is a major incentive for carbon monoxide to bond through the carbon end to the metal. You will notice, since the pi star orbital is vacant you can now pump electron density from a filled orbital on the metal into this pi star orbital. If you occupy the pi star, if you populate the pi star orbital with electron density then you will reduce the bond order between carbon and oxygen. It would become more easy to pull the carbon monoxide apart and so the stretching frequency is in fact reduced. So, that is how we understand the fact that carbon monoxide has got lower stretching frequency when it is bonded to a metal complex. Now, let us take a brief look on the metal side. Let us understand the metal first. First if you take a molecule like hexacarbonyl chromium you will notice that we need six vacant orbitals where we can accommodate the lone pairs on carbon monoxide. So, if you want to generate six vacant orbitals we have to generate a set of orbitals that can be formed from sp3d2 that is the hybridization scheme that we have learnt from valence bond theory which will lead to six orbitals which are pointed in the octahedral direction. Here I have pictured six orbitals and these six orbitals are coming from three p's, three p orbitals, one s orbital and two p orbitals. So, this is the sp3d2 hybrid and the d orbitals that are contributing to this hybrid are the d x squared minus y squared and the d z squared. So, these two orbitals are the ones which are coming from the d. However, the s and the p are completely involved in the hybridization scheme. So, that you generate six vacant orbitals on the metal and these six vacant orbitals are pictured here. Notice that three orbitals are left without any interaction when you do this hybridization. If I have six ligands which are coming along the x y and the z axis and that is these axis pictured here. If I have all of these axis having one carbon monoxide each then the three orbitals that will be left behind without too much interaction with the ligand would be what we call as a t to g set of the d x z d y z and the d x y orbitals. These are also pictured here. So, this hybridization scheme is very useful in understanding how we will be involving only the d x squared minus y squared and the d z squared orbital in the d manifold. Let us now proceed to mix what we learnt from carbon monoxide and use what we have learnt from the metal side and try to form a combination that will be stabilizing. So, the first important combination is the donation of electron density from the carbon monoxide. So, carbon monoxide gives a pair of electrons to the metal. As I mentioned before for convenience we have only pictured the d x squared minus y squared orbital, but it is actually a hybrid of the d s and the three p orbitals. So, this hybrid accepts electron density from the carbon monoxide. So, it is like a sigma bond that can be formed because any amount of rotation around this axis which we have between carbon and oxygen will lead to no change in the bonding at all. So, that is a sigma interaction. So, that is very important for us to understand. This is a sigma interaction. In the next case we talk about the pi interaction. The pi interaction is the one which decreases the carbon monoxide stretching frequency. As I mentioned earlier it is the pi star orbital on the carbon monoxide which is capable of accepting electron density from the filled metal orbitals onto the carbon monoxide empty orbitals. This will clearly make the carbon oxygen bond weaker because it is a pi star orbital or an anti-bonding orbital and it will also increase the c o bond distance. So, it will increase the c o bond distance and it will decrease the c o stretching frequency. For the sake of completeness let us take a look at the metal orbitals. In the case of chromium we have a total of 6 electrons and so these 6 electrons will be populating the 3 orbitals which are here in the t 2 g set. And if this is now filled you will notice that they are completely ready to donate a pair of electrons from here into the pi star orbitals of carbon monoxide. So, we have in fact a very nice complementary or synergistic interaction. The carbon monoxide gives a pair of electrons and the metal also gives a pair of electrons from the metal into the carbon monoxide and this leads to a very stable situation. So, as I mentioned we have synergistic bonding the sigma bond and the pi bond are operational here in the same system. And since these two electron density flows are in opposite direction we can have a metal which is neutral or even negatively charged and still form a very nice complex with the 2 electrons species that is carbon monoxide and the metal. Now, how do we show that these factors that we have been talking about are indeed happening in metal carbonyl chemistry. Let us just take a look at the effect of charge since we have mentioned that we can have neutral and negatively charged complexes. Let us just take a look at what happens when you have a negative charge on the metal. So, these complexes are unlike the Werner complexes have got excess electron density on the vanadium. So, vanadium is in the minus 1 oxidation state in this case and titanium is in the minus 2 oxidation state in this case. And we will notice that the stretching frequency the average stretching frequency of carbon monoxide in these octahedral compounds are significantly reduced from what we observe in hexacarbonyl chromium. Free carbon monoxide has got a stretching frequency of 2, 1, 4, 3 centimeter minus 1. Let us just to recollect what we discussed in the previous lecture and now it is decreased significantly in chromium hexacarbonyl and even further when you add a negative charge. So, as the metal has more and more negative charge it will pump in more electron density into the carbon monoxide pi star orbitals. So, it is natural that we have this greater reduction in the stretching frequency. So, that if you want to put it in orbital terms the overlap between the metal d orbital and the carbon monoxide pi star orbitals will be much better when you have a negative charge on the metal. That is because the greater electron density will increase the size of the metal orbital. It will try to expand because of the repulsion between the electrons are much greater and the compare to the nuclear charge which is holding these electrons. So, negative charge tends to increase the size of this cloud electron cloud and better overlap means more electron density can flow from the metal into the carbon monoxide pi star orbitals. So, that explains what happens in these negatively charged systems and if you look at a few complexes which have been synthesized recently by straws and co workers. We find that the carbon monoxide stretching frequency can be close to the free carbon monoxide stretching frequency as in this MNCO 6 complex where you have a positive charge on the metal or it can even be higher than the stretching frequency that you observe in free carbon monoxide. So, here we have an increase in the stretching frequency nearly 57 centimeter minus 1 increase is observed in the case of FeCO 6 2 plus. So, how do we explain this? One factor that readily comes to mind is the fact that because the orbitals have now contracted with the increased positive charge on the metal the orbitals contract in size. So, when they contract in size the overlap between the carbon monoxide pi star orbital and the filled orbitals on the metal they will decrease. This decrease in this size results in poorer electron density transfer from the metal to the carbon monoxide. So, if there is no pi star population the frequency is obviously in it cannot go down below the value of 2 1 4 3 centimeter minus 1, but we still have a problem. We have to explain why the frequency has increased in this complex to 2200 centimeter minus 1. This is a strange factor that has been observed only recently because these complexes have been characterized fairly recently by straws and co-workers through some very careful chemistry because these complexes are extremely moisture sensitive and are not very stable, but still they have managed to characterize them crystallographically and show and spectroscopically and show that indeed it is possible to have a positive metal ion interacting with the carbon monoxide. So, let us take a look at this a little more closely. Before we proceed to the positively charged compounds let us just take a look at a fairly recent paper that has appeared in the journal of organometallics and this paper deals with the generality of back bonding or this pi bonding that we have synergistic bonding that we just referred to the ligand gives electron density to the metal and the metal in turn populates the ligands anti bonding or empty orbitals. So, that we can have synergistic interaction. So, they have looked at a large number of transition metal compounds which have got carbon monoxide as a ligand and when they examined all the bond distances that are available in this crystal structure database that is a crystal structure database analysis CSD crystal structure database and when they did that they found some very interesting trends. I would like to briefly mention that because this tells you how important back bonding is in organometallic chemistry in metal carbonyl chemistry in particular. So, they what they carried out is a database analysis, but let me explain to you what this figure means. If you take a metal carbonyl complex in the same complex if I measure the metal carbon bond length and that is plotted on the x axis. So, the x axis is actually the metal carbon bond length and if this metal carbon bond length is combined with the carbon oxygen bond length for the same system. So, if you have a metal C O and this bond length goes in the y axis. So, this goes on the y axis and the metal carbon bond length goes on the x axis. So, for each metal carbonyl complex you will have 2 points of reference and if I take this this indicates that for the particular system that I am talking about I have a bond distance approximately 1.825 angstroms and a carbon oxygen bond length of 1.06 angstroms. So, if you do this particular graph if you plot this graph with all the metal carbonyl bond distances you observe for the sets that you have for molybdenum complexes you have a variation in this particular fashion. All the bond length seem to fall in an approximately clustered around a straight line, but you have a negative correlation. In other words as you shorten the molybdenum carbon bond length as you shorten it the carbon oxygen bond length seems to increase. So, the carbon oxygen bond length increases as you shorten the metal carbon bond length. This is in fact what you would observe for a synergistic interaction because if the metal pumps in more electron density into the carbonyl group then the metal carbon bond length should decrease because pi bonding is better and if this happens then the pi star orbital is occupied and the carbon oxygen bond length should increase. So, these two are negatively correlated and so you would expect a straight line with a negative slope and that is what we observe. So, this is very comforting because it is supported not just in a few complexes that we have looked at in the previous slide, but in a wide range of complexes in the database that is available now. What is even more interesting is a fact that if you take a series of compounds the 3D transition metal carbonyl complexes you have a trend which is indicated in this graph right here and you can plot the same type of analysis, same type of a plot as a scatter gram as we would call it. This scatter gram can be done for 4D and 5D transition metals and then you find that the graph is very similar. It has got a sigmoidal nature it has got a sigmoidal nature and you can notice this by the fact that there is an S shaped graph here and the graph is similar for both 3D and 4D metal ions. The only thing that has happened is that the 4D graph is shifted along the x axis and this shifting along the x axis happens because 4D and 5D transition metals are much larger than 3D transition metals and this larger size of 4D and 5D is responsible for this horizontal shift of this graph. But otherwise all the other factors appear to be the same so let us take a look at what would be the factors which are responsible for making a sigmoidal graph when you look at all the complexes put together. So here is the same graph but now I have redrawn it in such a way that we have marked the three regions which deviate from the linear portion of the graph. This is the linear portion of the graph which had a negative correlation between carbon oxygen and the metal carbon distances. So this linear portion of the graph was what we looked at initially and then we looked at all the distances together and we realized that in several systems there is a slight deviation from linearity at the two extremes when you have very large metal carbonyl distances and that is pictured here towards the end. These are the systems which have got very weak interactions between the carbon monoxide and the metal. That is a very small fraction of the total number approximately 6 percent of the metal carbonyl complexes have got this weak interaction between carbon monoxide and the metal. You will find that in these cases the sigma bonding is probably more important and there is very little very little very little pi bonding very little pi interaction. Whereas in the other extreme you have a small group of complexes which are pictured here and these small group of complexes have got very short metal carbon distances. So they have a strong double bond character between the metal and the carbon monoxide. So one can almost draw this valence bond structure. So you have a short distance and this is led to an increase in the carbon oxygen bond length. So the carbon oxygen bond length is significantly elongated and this is also a small fraction of the metal complexes that we have. This is approximately 4 percent, but the 90 percent of the complexes are the ones which have a negative correlation negative linear correlation between the metal carbon and the carbon oxygen bond length. So when we have multiple effects like this you have a sigma bond and you have a pi bond. Then you have two different effects operating. When only one is operating obviously the slope of this graph would change and that is indicated in the two extremes of this scatter graph. So this is a very indirect but very clear evidence of the fact that you have two effects operating and in chemistry whenever you have multiple effects you have multiple effects on a particular relationship then you will have non-linear behavior. So this is a clear case where you have two different effects pi bonding which is leading to population of the anti bonding orbitals of carbon monoxide and reduction in the carbon oxygen bond order. You have another effect which is a sigma bonding which results in reduction in the electron density on carbon monoxide. It is coming from this 3 sigma orbital on carbon monoxide which we mentioned earlier mostly populated on carbon and this electron density is something that we will have to discuss later in greater detail during the course of this lecture. So we provided very important proof that the pi bonding is important. You have a strong metal carbon double bond that is formed as we use the electron density on the metal to populate the pi star orbitals and this will reduce the CO stretching frequency. This CO stretching frequency that is reduced on carbon monoxide is significantly affected by the trans ligand. I want to tell you that the trans ligand plays a more important role than the other ligands which are present on the metal complex. So let us take a look at the reason why the trans ligand is so important. Let us consider an octahedral complex again or a square planar complex and in these complexes the T2G set or the DXY, the DXZ orbitals are the ones which are involved in pi bonding and this is pictured here. You will notice that the DXY or the DXZ orbital that we are talking about is pointed towards the carbon monoxide but half of that orbital is in fact pointed in the other direction also. So if you have another ligand in the trans position which is also capable of interacting with the pi or this pi orbital on the transition metal. So here is the metal and here is the ligand which can interact in a pi fashion. The trans ligand in fact shares the same orbital for pi interactions. So you had a carbon monoxide in one side and another ligand in the trans position which is capable of interacting in a pi fashion. So what will happen is the following. Either you will have competition for the electron density that is there on the metal. The electron density that is there on the metal has now to be shared now between the carbon monoxide on one side and the other pi ligand on the opposite side on the trans side. So the same because the same orbital is shared between two different ligands you tend to have very significant effect due to the trans ligand. So let us take a look at some of the effects of this. It not only changes the CO stretching frequency it will also change the metal bond order and hence the bond distance and we will be able to see such effects very clearly in the crystal structure. So I have given you two I am going to show you two examples now where you have carbon monoxide attached to the metal. Here is an iridium complex in this iridium complex you have two carbon monoxide. This is a carbon which is in black and oxygen is pictured in red. So there are two carbon monoxide ligands and you will notice that these two carbon monoxide ligands have got two different trans ligands. When you have a chloride as a trans ligand the metal carbon bond distance is 1.87 angstroms. So this bond distance is 1.87 angstroms and this carbon oxygen bond distance which is shown in green is not very clear. So I am writing it again is approximately 1.08 angstroms. So you will notice that in this particular system you have an iridium complex. When the chlorine is in the transposition what happens is that chloride ligand can have a filled has a filled pi orbital and so the filled pi orbital will not compete with the carbon monoxide for pi bonding and so what happens is that this distance becomes very short and this distance. So this distance is in fact if you compare the two carbon metal carbon bonds this distance is short and this carbon metal distance which is pictured here is close to 1.89 angstroms. So the ligand on the other side is in fact a ligand which can have partial double bond character or this is also a pi acceptor ligand. So this is a pi acceptor ligand in the transposition and so this metal carbon bond is elongated and this metal carbon bond is shortened compared to what you would expect for carbon monoxide in a simple system without pi interacting ligands. So this is a pi donor the chlorine is a pi donor and will not compete with carbon monoxide for pi accepting. So this bond becomes very strong this becomes very strong and this becomes very weak the carbon oxygen becomes very weak. So this double bond is weakened and so you have a bond distance of 1.08 angstroms. Whereas in the other case where you have a carbene ligand which is on the transposition this bond distance is increased the metal carbon bond distance is increased to 1.89 angstroms. So that is elongated and this is in fact this is the distance which is in fact weak because you have a competing this metal carbon bond distance is longer and it is weak because you have a competing pi acceptor ligand. So let us take a look at another system where you have a different type of a pi interaction. Here is an amino benzoic acid this is an amino benzoic acid which is coordinated to a rhodium compound and in this particular instance you have again a chlorine which is a pi donor and this is a sigma donation sigma donor ligand. Once again you find that when you have a sigma donor ligand this bond length is 1.841 angstroms and it is shorter when you have a pi donor ligand in the transposition. So this is shorter this is longer and that is because this is only a sigma donor ligand this is a pi donor ligand. So this pi donation has a significant effect in the transposition and you have a shorter metal carbon bond and you have elongation of this carbon oxygen bond this becomes longer. You will see that this distance which is again in green I will write it out for you once more 1.13 angstroms is a bond distance in the transposition. So you can see that the trans ligand has a significant influence and the carbon monoxide bonding because they have to share the same orbital they have to share the same orbital when you have a metal carbonyl bond and there is sigma and the pi interactions are complementary they are in fact synergistic. So you have significant changes when you either have a sigma donor or a pi donor or a pi acceptor. So far we have been looking at structural evidence and we have also been looking at some spectroscopic evidence for the fact that there is significant pi bonding in the metal carbonyl complexes. In fact there is also some chemical evidence for pi bonding and by chemical evidence what I mean to say is that if you react chromium hexacarbonyl in principle if I have two ligands coming in I should be able to form both the cis complex and the trans complex. So if I treat C R C O 6 with two equivalents of trimethylamine I should get both cis and trans and if I have more than two ligands that is if I treated with three ligands then I should have both the facisomer and the merisomer. But surprisingly from the reactivity of chromium hexacarbonyl we can see that the cis complex is the only one which is formed when there is a dye substitution and it is only the facisomer which is formed when you have a trisubstitution. This is because you would rather have a ligand in the transposition which is not a pi acceptor ligand and if you have a sigma donation ligand sigma donor ligand as trimethylamine then it would be better to have it in the transposition where it will push electron density into the metal and lead to better carbon monoxide metal binding. So this trans ligand in the case of the tris complexes where you have substitution of three ligands has got very nice relationship with the carbon monoxide stretching frequency and so I am going to show you this relationship once more in a graphical form. What I have pictured here is the set of complexes where you have P R 3 or three R groups which are attached to the phosphorous and they because of the electron withdrawing nature of fluorine or chlorine here I have fluorine on the phosphorous here I have chlorine on the phosphorous. So P C L 3 and P F 3 will have difference in the ability to donate or remove electrons from the molybdenum and so if you have P F 3 then you have less electron donation into the molybdenum. If you have less electron donation to the molybdenum then the carbon monoxide will have higher stretching frequency and that is exactly what we see here and if you substitute each one of these chlorins with R groups here I have shown you an example where two of the chlorins have been substituted by phenyl groups then I tend to have smaller stretching frequency and that is because chlorine withdraws electron density more than a phenyl group and so by virtue of the electron density withdrawn from the molybdenum or the poor electron density in molybdenum you have changes in the CO stretching frequency of the trans ligand. P M E 3 you have a very strong electron donor and so you have smaller or weaker stretching frequency for the three carbon monoxide in the transposition. So you can in fact plot this and again you will see a linear relationship in a small region where you have only one of the effects dominating and that is in this particular case it is mostly the pie effect which is being observed. So you have the effect of the trans ligand you have P F 3 this is P F 3 and the right mode side and P C L 3 this is P C L 3 ligand and P O E D 3 if you look at the stretching frequency of all these compounds then you tend to have a linear relationship in the stretching frequency of the carbon monoxide which is present in the transposition. So let us now discuss the importance of the sigma bonding orbital earlier we have discussed extensively what pie donation does in terms of spectroscopy what it does in terms of the bond distances and now I want to briefly mention what exactly happens when you donate electron density from the carbon monoxide in the sigma orbital. Let me just briefly change over to remind you of what we had discussed earlier the effect of charge and effect of sigma and pie bonding the two fold one is the fact that you have electron density donated into an empty orbital on the d orbital on the metal from the carbon monoxide which is in a linear which is shown here along say the x axis and it is donating electron density into the d x squared minus y squared orbital. So what we have is in fact an effect due to mostly the sigma bond when you have charge on the positively charged metal atom I told you that the positively charged metal atom will have a smaller d orbital because of contraction of d orbital and so the pie bonding becomes insignificant. So the pie interactions becomes in become insignificant when you have a positive charge on the metal. So the primary interaction when you have a positively charged metal atom is due to sigma bonding. Now how does this affect the c o stretching frequency and how does it affect the bond length first of all because you the pie effect is an additional effect you tend to have weaker carbon monoxide metal interactions and you have longer metal carbon bonds. But it seems to be having a strange effect on the carbon monoxide stretching frequency. So the charge effect that is recently been discussed about is the fact that when you have a positive charge next to the carbon on the carbon monoxide then it tends to have an influence which increases the carbon oxygen bond order. So we will discuss this in the subsequent slides. There is also another rationale that has been given and that is the fact that the highest occupied molecule orbital on carbon monoxide is in fact slightly anti bonding in nature. And because it is anti bonding you tend to have the reduction in the electron density on this homo which is slightly anti bonding you tend to have stronger c o interaction once you remove electron density. But compelling evidences come to indicate that it is primarily a charge effect and that is if you place a positive charge next to the carbon monoxide in the carbon end then you tend to have a stronger stretching frequency for carbon monoxide primarily because the electron withdrawing nature of carbon is enhanced by the positive charge close to the carbon monoxide. So let us first look at some of the positively charged complexes it is not easy to make positively charged metal carbonyl complexes. Here I have shown you nickel tetra carbonyl zinc 2 plus is an iso electronic system but it does not form a similar complex at least not easily. You will remember that nickel carbonyl was very easily generated by passing carbon monoxide at one atmospheric pressure but zinc 2 plus on the other hand does not form a similar complex at all and it will form very nice burner complexes with compounds like water. As I told you straws generated or synthesized a wide range of complexes with positively charged metal ions and interestingly the stretching frequency in all these cases ranges is in the higher side compared to 2 1 4 3 centimeter minus 1 which is the free carbon monoxide stretching frequency which is indicated here the free carbon monoxide stretching frequency is indicated here as 2 1 4 3 centimeter minus 1 and in the positively charged metal complexes the frequency is significantly enhanced to 2248. As I mentioned to you there are two factors that we have to note one is a fact that if you remove an electron from carbon monoxide and generate C O plus the stretching frequency goes up. This appears to indicate that removal of electrons from the highest talk about molecule orbital tends to have a strengthening effect for the carbon monoxide bond. So, the bond now becomes stronger. Now, two explanations were given one is a charge effect and that seemed to indicate that if you place a positive charge next to the carbon the electron withdrawing capacity of carbon is now increased. So, it will have greater pulling power of the electron density towards itself and will behave more like nitrogen. If it does that then the bond order between carbon and oxygen increases significantly and that is why the stretching frequency goes up. So, C O distances are in fact in the case of positively charged complexes it is in fact 1.115 angstrom this is in fact shorter than what you would expect. We had earlier mentioned that the pi effect has the pi effect increases this carbon oxygen bond length whereas the sigma effect now we see it shortens the carbon oxygen bond length. But as I told you in the graph that there are hardly 4 percent of the complexes where there is a delta plus or a slight positive charge on the metal atom. So, in most complexes the range is on the lower end it the frequency of carbon monoxide is in the lower end it is reduced from 2, 1, 4, 3 centimeter minus 1. But in a few complexes it is in fact enhanced that is because of the positive charge on the metal. When you have a very strong positive charge as in these two cases which are shown here then the frequency is about 100 centimeter minus 1 more than what you expect for free carbon monoxide. So, back donation is not important or significant when you have a metal which is positively charged when the metal is positively charged back donation is not important. If back donation is not important then the significant interaction is donation of electron density from the 5 sigma or the 3 sigma. In some books it will be referred to as 3 sigma because they ignore the core orbitals in other books it is referred to as a 5 sigma orbital. This is the highest occupied molecule on carbon monoxide and the carbon monoxide stretching frequency will go up if it is anti bonding in nature and that is one explanation that is given. The other explanation is that because you have a positive charge this positive charge makes the carbon delta plus and if carbon is delta plus it attracts electron density towards itself. It is behaving more like nitrogen and the carbon oxygen bond strength is temporarily increased and the frequency is enhanced. That is why the CO stretching frequency goes up in the case of metal carbonyls with a positive charge. So, here I have just reviewed for you all the orbitals which are of importance this is the sigma orbital which donates electron density. You will notice by looking at the 3 sigma orbitals this is the first sigma orbital this is the second sigma orbital and this is the third sigma orbital which is a highest occupied molecule orbital. The number of nodes between the carbon and oxygen keeps increasing and in this case it is the maximum and usually the anti bonding character is indicated by a node between the two atoms. Here although it is not extremely obvious there was a hint that there is a node between oxygen which is present here and carbon which is present here. That would lead to a slight weakly anti bonding nature for this orbital. However, it is now confirmed or it is assumed that the positive charge on the carbon is primarily responsible for the frequency shift and the anti bonding nature of this highest occupied molecule orbital is minimal if not non existent. So, C O let me summarize by saying that C O is rarely found bonded to positively charged metal atoms. In complexes where they are bound to positively charged centers then the frequency goes up beyond 2 and 4 3 centimeter minus 1. There are two suggestions which are given by two different groups of people. One group suggested that it was the anti bonding nature of carbon monoxide highest occupied molecule orbital. The alternate suggestion that was given by Franking and Strauss is now shown to be a better explanation and that suggests that the charge is in fact due to polarization of the pi cloud stronger pi bond because of the polarization of the pi cloud. So, that carbon monoxide now resembles nitrogen because the carbon has been ripped off an electron or some electron density has been removed from carbon and that positive charge makes carbon more electronegative and the polarization makes the pi cloud the pi bond between the two atoms stronger and you have stronger triple bond between carbon and oxygen. So, one other way to look at this phenomenon is to look at the stretching frequency in B H 3 C O because B H 3 does not have a pi filled pi orbital now to donate electron density. You have the stretching frequency going up beyond 2 1 4 3. A mild increase is observed about 20 centimeter minus 1 increase in the B H 3 C O complex. C O stretch goes up and that is indicative of the fact that you have a slight positive charge on the boron and this positive charge on the boron is responsible for increasing the C O stretching frequency. You will notice that this is not an electrostatic bond. It is quite strong. The B H 3 C O interaction is almost 23 kilo calories per mole in strength and so it is not just an electrostatic interaction. It is a covalent interaction between the B H 3 and the carbon monoxide. There are several reasons why carbon is bound to the metal in carbon monoxide. One reason is that the donor orbital is more localized on carbon. The second reason is that the acceptor orbital which is pi star on carbon monoxide has greater contribution from carbon leading to more stabilization when the carbon is bound to the metal. There is yet another reason. There is a pi orbital on carbon monoxide which has a similar p orbital on carbon which is also suitable for interacting with the metal filled orbital. If this filled orbital interacts with the pi orbital, there will be destabilization but because carbon contributes less to the pi orbital on carbon monoxide, it is better for the carbon to be pointed towards the metal. This will lead to a greater stability of the metal complex. One has maximized the bonding interaction by pointing the carbon towards the metal and reduced the repulsion due to the filled pi orbital on carbon monoxide interacting with the filled T 2 G orbitals. So, there are several factors that we have seen. Key features of metal carbonyl bonding. We have better bonding by having the carbon end pointed towards the metal. Better pi acidity of the carbon monoxide can be realized only if we point the carbon end because carbon monoxide has got pi star orbitals more concentrated on carbon. We also have reduced repulsion between the filled orbitals on carbon monoxide and the filled metal orbitals. So, this leads to reduced repulsion, better pi bonding and better sigma bonding when you have carbon monoxide interacting through the carbon end of carbon monoxide. Lastly, there is a charge effect which polarizes the carbon making it more electronegative when there is a positive charge on the metal. So, if you have a positive charge on the metal, then the charge effect makes the CO bond stronger. So, that leads to strong or higher stretching frequencies. So, we have covered several aspects of pi acidity and these are listed here. Based on what we have studied, one should be able to draw the structures of complexes, predict the structural features and also predict stability of the metal complexes. This should lead to a better understanding of the spectral features.