 Today, we will discuss one of the most common ligands in organometallic chemistry and that is carbon monoxide. Carbon monoxide is probably most frequently encountered in organometallic chemistry and as significant portion of the literature that is available today in organometallic chemistry deals with at least one carbon monoxide ligand in the coordination sphere of the metal. So, when we look at carbon monoxide and the type of complexes it forms it is interesting to note that almost the whole periodic table can be involved in this chemistry of carbon monoxide and metals. Significantly, we are going to talk about transition metal carbon monoxide complexes and another point of interest is the fact that all the transition metals and when we say transition metals in this color coded periodic table that we see here the transition metals are the ones that range from scandium to zinc. And specifically if you look at the complexes that are formed you can notice that they are all homoleptic complexes that can be generated using carbon monoxide. This is probably unique to carbon monoxide not all the other ligands that we encounter in organometallic chemistry can form homoleptic complexes. By homoleptic complexes we mean those metal complexes in which only one ligand is involved in the coordination sphere of the metal. So, if you look at the first row of the representative row of the transition metals. So, that goes from scandium to zinc and we can we will consider only the elements that are forming good carbon monoxide complexes and those are in the range from vanadium to copper. So, from vanadium to copper you do have frequent occurrence of carbon monoxide complexes and we will see in a moment why this is the case all of them are capable of forming homoleptic or m c o n complexes. So, these are complexes which have m c o n many copies of the same ligand are found in the coordination sphere of the metal. Let us just take a look at the systems that we have here on this table. You will notice that you have for all the metals which have even number of electrons that is chromium, iron and nickel you have no charges on them. In other words since carbon monoxide is a neutral ligand and you have n number of carbon monoxide it does not add any charge to the metal. The metal is found in the 0 oxidation state. So, metal is in the 0 oxidation state and you have n number of carbon monoxide around them. So, if you take the metals which are there with odd number of electrons and so that is vanadium has got 5 d electrons and similarly manganese has got 7 valence electrons in the d shell. So, that might be d 5 s 2 or you might consider it as a total of 7 electrons in the valence shell and cobalt accordingly will have 9 valence electrons. So, the systems which have odd number of electrons tend to form complexes which have a charge. Now, if you go to classical coordination chemistry you will notice that the metals are usually found in positive oxidation states. In these systems where we have these odd number of electrons surprisingly you have a tendency to form negatively charge complexes. In other words the net charge on the complex is minus 1 in many of these cases and if it is minus 1 then it is since carbon monoxide is neutral the metal should be in a negative oxidation state. This is a rather strange instance because metal is usually considered as an electro positive element. So, one would not expected to form a negatively charged species. So, this is one of the first factors that we have to take into consideration when we look at the chemistry of transition metal carbon carbonyls. So, except there is one exception here which is the copper and we will try to explain that as a as time goes on. So, metal is in the 0 oxidation state and we also notice that the coordination number and the geometry appear to be restricted. So, let us take a look at the coordination number you will notice that in coordination chemistry Werner coordination Werner chemistry you will notice that the metal is in a positive oxidation state and you will also notice that many of the complexes are either octahedral or tetrahedral not many other geometries are encountered in coordination chemistry. Unlike that in organometallic chemistry you seem to be having a restriction on the coordination number and the geometry. So, the next aspect that we have to consider is the oxidation state and the coordination. So, this is something which we have just mentioned the metal is found in the 0 oxidation state and the coordination number and the geometry are well restricted. And specifically if you look at the compounds that are there in the first row if you look only at the mononuclear complexes they all have a variable geometry. But nevertheless it is distinctly attached to the number of carbon monoxide ligands that are present. In other words if I have 5 carbon monoxides present in the coordination sphere of the metal then the 5 carbon monoxides are distributed in a symmetrical fashion around the metal. If I have 4 carbon monoxides distributed around the metal then the 4 carbon monoxides are in a tetrahedral environment around the metal. Surprisingly when you talk about these negatively charged systems they are also capable of forming dimeric complexes. So, here are 2 instances where you have 2 dimeric complexes that are formed. It looks as if you can have some other ways of compensating for the number of electrons around the metal. If indeed the metal was forming this cobalt tetra cobaltate anion to compensate for the number of electrons it was having in the valence shell. Then there seems to be another way of forming the same number of valence electrons in the complex and that seems to be in the form of making a dimeric species. So, interestingly there is one species in this whole series which is iron which forms a dimer. But nevertheless there is a way of making the neutral complex with the right number of carbon monoxides around it. So, even considering the fact that you have a series of complexes right from vanadium to copper which are mononuclear or charged you find that the dimeric species are mostly in the neutral state. But they can also be generated by reacting them with a very good electron source like sodium amalgam or potassium amalgam. You can inject electrons into these dimeric species and also generate negatively charged systems. So, negatively charged metal complexes are formed in organometallic chemistry which is very different from Werner's coordination chemistry where it is very rare. Now what we have noticed in the series of complexes that we have just encountered is that all of them end up with a total of 18 electrons around the metal atom. This was the reason why people started talking about the 18 electron rule. Now the 18 electrons were generated by adding the number of electrons around the metal to the total number of electrons donated by the ligands. In Werner's coordination chemistry what you find is that there do not have this strict following of 18 electrons around the valence shell. Here for example, I have noted that cobalt 3 plus cobalt 3 plus has a valence electron number of valence electrons equal to 6. So, that has got 6 valence electrons it is a d 6 complex. Now if it combines with a water molecule 6 water molecules then you can form a very stable coordination compound which is cobalt a hexa a cobalt 3 complex and that is denoted here. Now you notice that the total number of valence electrons around the cobalt also equal 18. 6 electrons from cobalt and 6 into 12 electrons from the water molecules. So, a total of 18 electrons can be generated a total of 18 electrons can be generated on the cobalt ion. Unlike the cobalt hexa a co complex a very similar complex with chromium 3 plus now has got only d 3 valence electrons. So, it has got 3 valence electrons. So, 3 valence electrons on the chromium plus 6 into 12 electrons from the water molecule. So, you have a total of 15 valence electrons on this complex. So, you can see that Werner complexes even if they form very stable systems they like to be octahedral or tetrahedral, but they do not follow the 18 electron rule or they do not seem to be having an 18 electron rule governing their electronic structure. Unlike this the coordination unlike the coordination compound formed in Werner chemistry. Organovitalic systems tend to form 18 electron complexes. Let us just go back and look at the complexes that we have generated especially the neutral ones. It is very easy to do the electron counting in these cases. Chromium for example, we have just noted it has got a total of 6 valence electrons. Carbon monoxide gives 2 electrons and so 6 into 2 will also give you 12 valence electrons. So, 6 plus 12 gives you 18 valence electrons a total of 18 valence electrons are there. If you take iron, iron will give you 8 valence electrons and carbon monoxide 5 carbon monoxides can give you a total of 10 electrons. So, again a total of 18 electrons are generated around the iron. So, if you count the valence electrons which are available for each of these metal atoms, you will always end up with a total of 18 electrons. This is true of nickel, it is true of iron, chromium and even the vanadium case where you have 5 valence electrons on the vanadium. So, you have to add an extra electron to the vanadium in order to generate the carbonyl complex. In the case of copper, you have a problem positive charge and this is in fact a unique system. It does not seem to follow the 18 valence electron rule, but you can generate tetra coordinate CuCO4 complex also and that is possible and people have generated a range of complexes in this case, but the number of complexes which are positively charged and have only carbon monoxide in the coordination sphere are rather small and copper gold and silver are some of the species which form positively charged complexes and you will notice that if it is CuCO4, then it would have a valence electron count of 18 also. So, let us proceed further. What we have noted is that the organometallic complexes that we are encountering in these systems strictly follow an 18 electron rule and this is very different from the Werner's Coordination Chemistry where you only have octahedral complex being formed or a tetrahedral complex and the driving force seems to be the restriction on the coordination number and geometry rather than on the valence electron count. Now, you will notice that in the complexes the Werner's Coordination Compounds that we have encountered, you have a variety of oxidation states. You might have a nickel 2 plus H 2 O 6, you can have a cobalt 2 plus H 2 O 6. So, it can be a cobalt 2 plus a co complex or it can be a nickel 2 plus a co complex and these complexes can be nicely colored. Here you can see them with very bright colors blue, green and pink and these complexes tend to be very highly colored. On the other hand if you take a chromium carbonyl complex which we just talked about C R C O 6, this is neutral and it is also interesting to note that it is a colorless species. So, although the oxidation state is one of the differences, it is also interesting that the color of these species tend to be completely colorless or at the most slightly colored at yellow in the case of iron carbonyl like F E C O 5 is light yellow in color. So, these complexes have a different color and they also have different oxidation states compared to classical coordination chemistry. So, nickel tends to form only 2 plus complexes, but cobalt can be cobalt 2 plus, it can also be cobalt 3 plus. So, variable oxidation state is a characteristic of vernis chemistry. This is typical of vernis chemistry whereas, organometallic chemistry tends to stick with simple neutral complexes when it comes to carbonyl homoleptic carbonyl complexes. So, this is also to be noted, the fact that magnetic properties of the Werner complexes are often either diamagnetic or paramagnetic, they tend to vary. It depends on the oxidation state, it tends to depend on the strength of the ligand. If you have very strong ligands, they tend to form low spin complexes or even diamagnetic complexes if the electron count is right. Whereas, in the case of organometallic complexes, almost all of them tend to be diamagnetic. This is a distinct advantage because you can carry out NMR spectroscopy of these complexes. NMR spectroscopy allows you to monitor these compounds in situ in a very easy fashion and this advantage is not present for the coordination chemistry, the Werner coordination chemistry that we encounter. Where you have paramagnetic systems and paramagnetic systems tend to destroy the NMR signals fairly easily. Vanadium carbonyl is an exception because that V C O 6 can be generated and if it is V C O 6, it tends to be paramagnetic, it is not diamagnetic. This is one exception that we note. Let us move on. I mentioned earlier that metal carbonyls tend to form reasonably symmetric structures. They are not necessarily tetrahedral or octahedral, although these tetrahedral and octahedral complexes are indeed observed. Typical examples are nickel tetra carbonyl which is probably the one that chemists study encounter first because that is very easy to synthesize. Chromium hexacarbonyl, that is another molecule and that has got a nice octahedral complex. It is a nice octahedral geometry. You will notice that the complexes are all having carbon monoxide symmetrically distributed around the metal. This seems to be a hallmark of the type of compounds that are formed by metal carbonyls. Apart from these mononuclear complexes, it is also possible to form multi-nuclear complexes with metal metal bonds. Metal metal bonds are reasonably rare in coordination chemistry. In the case of Werner's chemistry, you normally do not encounter metal metal bonds. If you do, you do have the possibility of bridging ligand in addition to the metal metal bond. But in the case of carbonyl chemistry, in the case of organometallic compounds which are supported by carbon monoxide, you see that metal metal bonds are reasonably common. A variety of complexes can be generated which have got both a metal metal bond and also a bridging carbon monoxide. First, let us take a look at those complexes that have a metal metal bond. Here is a set of complexes that I have pictured here. These compounds strictly follow the 18 electron rule. They still follow the 18 electron rule. They have carbon monoxide which are just indicated by a straight line here. For example, if you choose this complex, each one of these lines which I have drawn here represents a carbon monoxide. In order to make the structure readily understandable, all the carbon monoxides are just shown as sticks. So, each one of these lines represents a carbon monoxide. Each metal center has carbon monoxide. So, each metal center has got four carbon monoxides. Each metal center has got four carbon monoxides in this case. So, you form a M3 CO2 complex. This is very common for all the transition metals. In the case of ruthenium and osmium, you can form a system where you have a metal metal bond supporting the formation of a trimeric species. Here you have another instance of a tetrameric species. This is a system where you have iridium with three carbon monoxides. Each of these vertices of this tetrahedron that you have here, each one of these vertices is capped by an IR CO3 moiety. So, each one of these vertices is occupied by an IR CO3 moiety. That is the structure which is shown here. In order to simplify the structure, we have shown it with only simple lines. Another simple complex which is also having a metal metal bond with no bridges. This is unique for organometallic chemistry. As I mentioned before, this is Magnes decacarbonyl, dimanganese decacarbonyl complex which is pictured here. This has got a single metal metal bond which is holding the two units together. Once again, you will notice that because manganese has got seven valence electrons, manganese has seven valence electrons, you will need one. It is one electron short. If I form a complex like MnCO5, it is one electron short. It is a 17 valence electron species. Now, there are 17 valence electrons. There is one electron. If it can be shared between the two manganese atoms, then each manganese can end up with a valence electron count of 18. That sharing of that one electron between the two manganese atoms results in the formation of a bond between the two manganese atoms. That is what you have in each one of these cases. Either it forms two bonds as in the case of ruthenium. It has to form two bonds in order to achieve valence electron count of 18. In the case of iridium, it has to form three bonds in order to achieve a valence electron count of 18. The formation of a polynuclear species with a metal metal bond is primarily determined by the electron count. This again brings us back to what we discussed earlier. That is a fact that these metal complexes are significantly driven by the tendency to form 18 electron species. A total valence electron has to be 18. So, another aspect which I mentioned when I talked about polynuclear systems is a fact that you can have bridging carbon monoxides. The fact that you can have bridging carbon monoxides is again a unique factor, which is not found in coordination chemistry in the classical coordination chemistry. Carbon monoxide is a nice terminal ligand or in other words, it can form, it can donate two electrons and form a metal carbonyl bond. That is a terminal ligand. Surprisingly, carbon monoxide is found to be bridging in some of the dimers and the trimmers that can be generated, especially in the case of the first row of transition metals. This is very different from water and ammonia. Water and ammonia also two electron donors. NH3 can donate two electrons. Water can also donate two electrons, but they do not form dimers like the metal carbonyls. Here, I have pictured a series of trimeric and tetrameric species. In each one of these systems, you will encounter a bridging carbon monoxide. In some cases, it is bridging two metal atoms. In this case, for example, it is bridging two metal atoms. We have pictured it like this, where the two lines meet. That represents a bridging carbon monoxide. In other systems, where three lines meet, you have a triply bridging carbon monoxide. You have a triply bridging carbon monoxide, which is capping one of the phases of this octahedral species. You have three metal atoms. You have three metal atoms. A cap, a trigonal phase is capped by a carbon monoxide unit. This is a triply bridging carbon monoxide, which you find in this particular unit right here. You have two metal atoms being bridged by two carbonyl species. You will notice that in this particular instance, we encountered another D-8 system forming a trimeric ruthenium, for example. R-U-3-C-O-12 was also there, but it did not have this bridging carbon monoxide. This has raised a lot of debate as to the facility with which carbon monoxide can bridge and why it bridges 3D transition metals and not necessarily 4D and 5D elements. Both ruthenium and osmium do not form the bridge species, but instead notice that they form only the trimeric species, where you have no bridges, but terminal carbon monoxides. Here, you have all terminal carbon monoxides and in the case of iron, you tend to form two bridges. Iron is the same electron count, but you have two bridges. If you notice, the number of carbon monoxides around the iron atom, however, is identical. The 18 electron count is always maintained. 18 electron count around the metal is maintained, but the number of carbon monoxides can either be bridging or terminal and you can still achieve the same electron count, because the bridging carbon monoxide donates one electron to the iron on one side. This counts for one electron and this counts for one electron. This bond counts for one electron as well. So, the two electrons donated from carbon monoxide are now split between the two iron atoms. In the case of triply bridging carbon monoxides, electron counting is a little more difficult. However, we say that the two electrons are shared between three metal centers. So, we have to do a total valence electron count and then take care of the number of electrons available for the metal atom. So, aggregation at the formation of multinuclear complexes is very commonly encountered in organometallic chemistry. In all these cases, you can have bridging and terminal carbon monoxides in the carbon monoxide chemistry that we have just seen. In one unique instance at least, at least in the case of cobalt di cobalt octacarbonyl, which I have pictured here, it has been noticed that both terminal and bridging modes of carbon monoxide are encountered in the complex. In the case of the solid state structure of CO2CO8, which I have mentioned here, CO2CO8 in the solid state has got two bridges. In the solution state, it has been noticed that there are no bridges in the dimeric species that is formed. So, it appears as if that there is a very delicate balance between carbon monoxide being bridging ligand or carbon monoxide being a terminal ligand. Both of the systems appear to be equally fissile. Again, we have already encountered the fact that you can have iron tricarbonyl, which has got two bridges. Iron tricarbonyl has got two bridges and ruthenium tricarbonyl, which has got no bridges. So, the fact that you can form bridging carbon monoxide complexes and complexes where there are no bridges appears to indicate that there is a very small energy difference between the two states, between the bridge state and the non bridge state. So, here I have pictured the two structures of the di cobalt octacarbonyl complex. In the di cobalt octacarbonyl complex, in solution you find there are no bridges. In the case of the solid state structure, you have a very clear formation of a, you have a very clear indication of a bridge structure. In the case of di cobalt nonacarbonyl that is Fe 2 C 1 9, you find that there are three triply bridging structures. That invokes the necessity of a metal metal bond in order to satisfy the electron count, because if you have six electrons donated by the carbon monoxides and you have, let us do this electron count. So, each iron has got eight electrons in its valence shell. Three carbonyls, the terminal carbonyls which I will mark as T CO, they will give you six electrons as well. If I have this bridging carbon monoxide, if that gives you three electrons to each iron. So, you have a total of 17 valence electrons, you have a total of 17 valence electrons. So, as I mentioned earlier, since we are one electron short, we are one electron short of the magic electron count of 18, you tend to form an iron, iron bond which will involve sharing of one electron from each of these iron atoms. So, that a total valence electron count of 18 can be achieved. So, the iron, iron bond will contribute to one electron. So, that gives you a total of 18 valence electrons for the Fe 2 C 1 9. So, this is chemistry that is completely governed by this tendency to form 18 valence electrons around the metal atom. Now that we have considered the electronic structure around the metal atom, let us take a brief look at some of the properties of these molecules. Wernher coordination chemistry or Wernher chemistry typically involved many ionic compounds and these ionic compounds are solids. In most instances, they are solid complexes very difficult to evaporate. On the other hand, the physical properties of organometallic carbon monoxide complexes are completely different. Many of them or most of them are liquids or gases and even if they are solids, they can be readily sublime. So, their vapor pressure is quite high and that is a matter of concern when you talk about handling them. You need to be extremely careful because they are quite volatile. Now interestingly, many of them are also soluble in organic solvents and that led to the rapid development of organometallic chemistry because it was possible to use organic ligands or organic solvent in order to do the chemistry of organometallic compounds. It was possible to separate them from salts and coordination compounds fairly readily because one could just wash it away. The negatively charged the other important fact that we have already noted is the fact that negatively charged metal ions can be stabilized by carbon monoxide very readily. Those compounds are indeed ionic. They are charged and they are usually solids and they might be colored as well because of charge transfer phenomenon. Let us now move on to the structure of these compounds. We find that the structure of these compounds are those that are governed by this 18 electron rule. But more than that, they are not single iron metal carbon bonds. If you take the typical example of Fe CO 5, you find that the iron carbon bond that you would expect on the basis of the covalent radius of iron and the covalent single bond radius of carbon. If you add these two together, you would get a bond distance of 2.2 angstroms. So, 2.2 angstroms is what you would expect on the basis of adding the two single bond radii of the two elements iron and carbon. But the Fe CO 5 molecule turns out to be having an iron carbon bond distance of 1.83 angstroms in the equatorial position. So, here is a molecule which has a pentagonal 5 ligands. It has got a trigonal bipyramidal geometry. Each one of these lines represents a carbon monoxide. So, you have a carbon monoxide in each one of these points which ends at the end of this line. So, this bond distance is what we are talking about. This bond distance turns out to be smaller than what you expect on the basis of the single bond radius. The axial bond surprisingly is 1.81 angstroms and the equatorial bond is the one which is 1.83 angstroms. So, these distances are longer than the free carbon, the carbon monoxide distances are the ones that we are talking about here CO distance. This distance is longer than what is observed in free carbon monoxide by 0.01 angstroms. So, you have two things that we need to explain. First of all, the fact that all of them have got this 18 electron structure. Secondly, you have to explain the fact that you have shorter bonds than what you expect for a single bond, for a single metal carbon bond. The third point that we note is the fact that you have a carbon oxygen bond distance which is slightly longer. So, you have a very small elongation of the carbon monoxide bond and you have a shortening of the iron carbon bond. Before we proceed further and explain the electronic structure of these complexes, I want to briefly mention the spectroscopy that can be done with these compounds. These compounds have got very significant carbon monoxide stretches which change when they are complex. That is free carbon monoxide stretches that change when they are complex to the metal. The free carbon monoxide stretch which is something that every organometallic student or every student of organometallic chemistry memorizes is a fact that it is 2143 centimeter minus 1. So, this magic number that you have with carbon monoxide is because of the triple bond that you have between carbon and oxygen. So, it is quite difficult to stretch this carbon oxygen bond. So, this carbon oxygen bond which is difficult to stretch appears at 2143 centimeter minus 1, but when it is complex to the metal, it becomes easier to stretch this carbon and oxygen. So, this carbon oxygen stretching frequency drops down by 100 to almost 200 centimeter minus 1 when it is complex to the metal. So, this is another interesting phenomenon that we have to notice and that we have to explain when we look at the electronic structure. But this gives us a very nice way of studying the structure of these molecules because carbon monoxide stretch has got a very strong dipole moment change when it and so it can be readily observed in the infrared spectrum of these molecules. We will look at some of the ways by which it can be used effectively. In a few compounds, this carbon monoxide stretching frequency is not decreased, but it actually is increased. This is very rare. This is quite rare. This is centimeter minus 1 and those systems we find that the metal is positively charged. So, let us take a look at the bond dissociation energy of metal carbon carbonyl complexes. You will see here a series of numbers which are the energy required to break the metal carbonyl bond. Chromium carbonyl for example requires an average of 25 kilocalories per mole to break this C-R-C-O bond. So, this bond is worth about 25 kilocalories per mole. That is what we are saying in this table. Similarly, if you look at all the metals, you find that the average bond dissociation energy is not varying much as you go from left to right in the periodic table. All of them have an average of approximately 25 to 26 kilocalories per mole. On the other hand, as you go down the periodic table, in the chromium series for example, if you go from chromium hexagon carbonyl to molybdenum hexagonal, the bond energy seems to increase significantly. It is almost increased by 10 kilocalories per mole and that is about 33 percent increase and by another 16 percent increase when you go from molybdenum to tungsten. So, it is possible to have very strong bonds in the case of 4D and 5D elements. This happens to be a general phenomenon in organometallic chemistry. The metal carbon bond strength that you observe in most cases increases as you go down the group. For the same group, if you have a bond strength of 100 kilocalories per mole as you go down the group, you will increase it to by 150 or to 200 as you go to the 5D series. So, the metal carbon bond becomes stronger as you go down the group. Let us come to the synthesis of metal carbonyls, a little bit of chemistry before we go back to the understanding of metal carbonyl chemistry in terms of electronic structure. If you take iron powder and treat it with carbon monoxide at 200 atmospheres pressure and a high temperature or relatively high temperature to 200 degrees Celsius, you tend to form the volatile metal carbonyl complex which is FeCO5. This turns out to be the simplest way to make iron carbonyl and it is fairly inexpensive because carbon monoxide is cheap and iron powder is cheap and you can make iron carbonyl in significant amount. Iron carbonyl however is not readily available these days because it turns out that you can make very interesting metal coatings which can make for example, airplane invisible to the radar by making a coating of iron which is generated from iron carbonyl. So, it becomes difficult to make or purchase iron carbonyl in spite of the fact that it is very easy to make. So, let us move on here we have iridium complex being generated with carbon monoxide COD which is pictured here is which is indicated here as COD is in fact, cyclo octadiene which is normally 1 5 cyclo octadiene and 1 5 cyclo octadiene and L is the ligand trimethyl phosphine which is easy to understand. So, if iridium has got these two ligands cyclo octadiene is a very weak ligand we will encounter it later in this lecture series. It can be replaced by carbon monoxide to form this familiar complex which is a square planar complex with CL and CO in the trans geometry and 2 L ligands which are PMA 3 coordinated in like fashion. So, this is the complex which you have here this is the complex that we are talking about and this is very readily generated using carbon monoxide and carbon monoxide and this is a COD ligand which is a cyclo octadiene ligand. So, in the previous two examples we have generated the nickel generated the metal carbonyl complex starting directly from the metal atom in the 0 oxidation state or in the plus 1 oxidation state as in the case of iridium. Here we have started with nickel in the plus 2 oxidation state and we are going to reduce it. We can also reduce it using carbon monoxide itself. Carbon monoxide is a reducing agent because we can oxidize carbon monoxide using the oxygen atoms which are available in this oxide material and converted to carbon dioxide. So, carbon monoxide is in fact a reducing agent. So, you can use carbon monoxide as a reducing agent or you can use an inorganic reducing agent as S 2 O 4 2 minus. So, either way you can generate a metal complex fairly readily by taking a slightly oxidized system a 2 plus species in that particular case nickel or even in the case of iridium you have a plus 7 oxidation state and it is you can still reduce it using carbon monoxide to generate a metal carbonyl complex. Here is one example where we have a carbonyl which is a 0 oxidation state which is in the 0 oxidation state. This particular complex needs a little bit of explanation. Here you have 2 carbon monoxides trans to each other and you have 2 nitrogens which are coordinated in this fashion with 2 methyl groups on the nitrogen. That is this ligand T M E D A and you have 4 carbon monoxides, 2 carbon monoxides trans to each other and 2 carbon monoxides cis to each other. Now, if you treat this metal complex with sodium, sodium as I told you earlier can pump in electrons especially if it is used as an amalgam, sodium and mercury turns out to be a very good reducing agent. It can generate what we can call here as a chromium 2 minus complex because the negative oxidation state of chromium here is in this particular case it is 4 minus 4 electrons are being pumped in and you have chromium in 4 minus oxidation state. So, these are unique complexes where you have high negative charge on the metal atom and this is usually generated with either sodium in mercury or sodium as molten sodium that you can generate by slightly warming the sodium metal. Nickel is the only transition metal that reacts with carbon monoxide to form a carbon complex at room temperature and pressure. This is popularly known as the Mons process because he was the person who first discovered the reaction of nickel and carbon monoxide leading to the formation of nickel tetra carbonyl. Now, let us get back to the fact that the carbon monoxide stretching frequency has reduced from 214 3 centimeter minus 1 to 1850 to 2100 approximately 100 to 150 centimeter minus 1 decrease or 100 to 200 centimeter minus 1 decrease in the stretching frequency. So, this decrease has to be explained when we look at the bonding. So, let us go and look at the bonding that is involved and we also note the fact that this carbon monoxide stretching can vary depending on the type of carbon monoxide that you have. If it is a terminal carbon monoxide, then the stretching frequency decreases significantly and if it is bridging carbon monoxide, then it is more like a ketonic carbon monoxide. The carbon monoxide stretching frequency can be as low as 1600 centimeter minus 1 when it is triply bridging as in the case of the compounds that we encountered the iridium polynucleic systems. We had triply bridging carbon monoxide and in those cases the carbon monoxide is in fact more like ketone or C double bond O or a C plus O minus species O. You have very low carbon oxygen stretching frequencies. It is very useful to use carbon monoxide to identify the type of compounds that they form. For example, I have pictured here a MCO2 system. One of them is trans. Here is a trans compound and here is a cis compound and the two carbon monoxides can be readily distinguished because in the infrared spectrum of this molecule you will observe for the trans isomer only a very strong signal, a single strong stretching frequency which is coming from the asymmetric stretching of the two carbon monoxide units which will couple to each other. Here you have a symmetric stretch where the two carbon monoxide units are pulled in a symmetric fashion. Both of them are either pulled apart or in the other case where you have an asymmetric stretch. This is the new asymmetric stretch. Then you have pulling apart of the two carbon and oxygen bonds, pulling apart of the carbon oxygen bonds and in the other carbon oxygen you have a compression of the C and O. This leads to an asymmetric stretch which is observed as a very strong band in the infrared spectrum. But the symmetric stretch is hardly is a very weak signal or if it is not observed in some cases. On the other hand if you have a cis complex, then you tend to have a strong band for the symmetric stretch and a strong band for the asymmetric stretch. So, carbon monoxide infrared spectra are very useful. You can also use it for triply bridging systems. Here I have shown the two possibilities when you have carbon monoxide. You can in an octahedral complex if you have triply ligated systems. Then you have triply carbon monoxide ligated systems. You can have the meridional isomer or the facial isomer. The meridional isomer will have three bands as pictured here. Three infrared stretching frequencies usually the two asymmetric stretches which you can have for this complex are strong bands and the symmetric stretch is weak. Whereas, in the case of the facial isomer you can have depending on the symmetry, you can have a small symmetric stretch and a very strong asymmetric stretch. So, based on the elemental formula, you would be able to distinguish the MCO 2 and the MCO 3 systems. Once you know how many carbon monoxides are there, you can figure out the geometry of the complex, whether it is cis or trans, whether it is facial or meridional depending on the number of carbon monoxide bands that you have. Carbon monoxide complexes can also be studied using carbon 13 NMR spectroscopy and the only difficulty in these cases is the fact that the natural abundance of carbon is only 1 percent. Since, carbon is not a very sensitive nucleus, you tend to have very weak signals for the carbon monoxide in MCO N complexes. So, what one usually does is you add some agents, external agents into the sample so that the relaxation is fairly fast. So, you can have reasonable signals for the carbon that is the 13 C in the carbon material carbon. Now, we have considered a lot of things, a lot of aspects of metal carbonyl chemistry. We have studied that they tend to form 18 electron complexes. They tend to form specific structures which are even negatively charged and they are very different from Werner coordination compounds. They can be studied by infrared spectra, which indicate significant decrease in the carbon monoxide stretching frequencies. Lastly, I also want to ask this question, why do they form at all? Because carbon monoxide is a very stable molecule. It is very difficult to generate isolated metal atoms. So, isolated metal atoms are formed in the gas phase, if you vaporize them, but this requires a very large amount of energy, very often 1000 to 2000 k. If you heat a metal atom to a very high temperature then you can vaporize the metal. So, why is it that you can readily form metal carbonyl complexes? In the case of nickel, you can even pass carbon monoxide at room temperature and atmospheric pressure and you can form nickel tetra carbonyl. So, these two complexes are even more interesting. You can readily form metal carbonyls at room temperature with carbon monoxide. Iron will form FeCO5 very readily and this can be done even at room temperature using carbon monoxide with very pure iron, which can be readily distilled. These are volatile molecules, which are readily distilled. So, carbon monoxide is a stable molecule. It has got a very small dipole moment. The carbon end is negative and the protonation of carbon monoxide leads very little energy, but it does give you can protonate carbon monoxide. That being the case, why is it that they form such stable compounds? In fact, theory estimates that the bond dissociation energy of metal carbonyls can be as high as 70 to 100 kilocalories per mole. The carbon oxygen bond and the metal carbon bond are significantly different from what you would expect. So, all these things have to be explained. Let us take a look again at the structural parameters. We have already mentioned this that based on the iron carbon bond distance, you expect distance of 2.10 and what you get is a much shorter bond distance. The carbon oxygen distance is in fact elongated. It is usually between 1.14 and 1.16. You have a spectroscopic feature, which suggests that this frequency decreases. So, all these things have to be explained. One way to do this is to use molecular orbital theory. We will use molecular orbital theory to explain the bonding in carbon monoxide in detail in the next lecture. We will introduce this structure of carbon monoxide itself right now. We will take a look at some of the valence orbitals. The valence orbitals of carbon monoxide are the ones which have got a very large contribution from the carbon end. So, this is a carbon end and this is the oxygen end. The valence orbital has got these are only the sigma orbitals. You can notice that the sigma orbital has got a very large contribution on the carbon end. These are the pi orbitals now. You have a pi and a pi star. This pi star also has large contribution on the carbon end. This is the introduction to the molecular orbitals of carbon monoxide. Now, we can see how these valence orbitals can interact with the metal in the next lecture. Let us just summarize some of the aspects that we have learnt in metal carbonyl chemistry. These are the key features that we have encountered in today's lecture. First of all, 18 electron complexes appear to dominate. There are very few complexes in which you have, very few complexes in which you have more than or less than 18 electrons. You also encountered the fact that neutral metal atoms, neutral metal atoms and not ions as in Werner complexes are encountered. In Werner complexes, mostly they are charge systems. Thirdly, carbon monoxide can bridge and form both terminal and bridging carbon monoxide units. Also, we encountered several systems where there are metal metal bonds even when there are no bridging ligands. This is again uncommon in Werner's chemistry. We will study in the next lecture that the best explanation for all these factors is to use what is called the DCD model, where there is a give and take of electrons. This appears to be a general phenomenon not just in metal carbonyl chemistry, but in the whole of organometallic chemistry.