Structural Biochemistry/Organometallic Chemistry

Introduction
Organometallic chemistry is the chemistry of compounds that contain metal and carbon bonds. They can form various reactions, similar to organic chemistry. For example, the compounds Cr(CO)6 and [Ni(H2O6)]2+ are both octahedral. The CO and H2O ligands are sigma donors. CO is also a pi bond acceptor. Other ligands like CN-, PPh3, and SCN- can act in a similar way as they can function as both sigma donors and pi acceptors. Sandwich compounds are formed by cyclic organic ligands with a delocalized pi system that bonds to metal atoms. Cluster compounds are two or more metal atoms bonded to organic ligands like CO. Cluster compounds can come in all shapes and sizes. Carbon centered clusters are called carbide clusters, which are carbon atoms that are surrounded by five or more metals.



Background
The first organometallic compound to be reported was synthesized by Zeise in 1827. He obtained yellow needle-like crystals after mixing PtCl4 and PtCl2 in ethanol. KCl was then added as well. However, the structure of this compound was not determined until more than 100 years later. The actual ionic formula of the compound was K[Pt(C2H4)Cl3]*H2O. This structure has three chloro ligands occupying the corner of the square with ethylene occupying the fourth. In 1898, Barbier and Grignard performed reactions between magnesium and akly halides, which resulted in the synthesis of alkyl magnesium complexes known as Grignard reagents. These lead to great advances in organic chemistry, but wasn't as significant in organometallic chemistry. However, this was soon changed with the synthesis of sandwich compound ferrocene, which lead to the current modern era of organometallics.

Nomenclature
The number of atoms that a ligand bonds is indicated by (eta) followed by a superscript indicating the number of ligand atoms attached to the metal. For example, the cylopentadienyl ligands in ferrocene bond through all ligands, so it is named as eta 5 - C5H5. This is then named as Pentahaptocyclopentadienyl due to five bonding positions. Some common ligands are named as carbonyl, carbene, carbyne, cyclopropenyl, cyclobutadiene, cyclopentadienyl, benzene, etc.

The 18-electron Rule
In regular chemistry, the octet rule functions as a way for main group compounds to have a full valence shell of 8 electrons. However, in organometallic chemistry, the metal ion can have up to 18 electrons. For example, the compound Cr(CO)6 has 6 electrons since it's in group 6. Each of the CO then donates 2 electrons, for a total of 18 electrons. This configuration is incredibly stable. Other compounds like Cr(CO)5 with 16 electrons and Cr(CO)7 with 20 electrons are far less stable. Two general ways can be used to count the number of electrons in the complex. The first method is the donor pair method, which considers ligands to donate electron pairs to the metal. The charge on each ligand and the formal oxidation state of the metal must be taken into account to determine the charge. For example, Cl- can donate two electrons and CO can also donate 2 electrons whereas C5H5 can donate 6 electrons. The second method is called the neutral-ligand method. This takes into account of the number of electrons that would be donated if the ligands were neutral. For example, Cl has a -1 charge so it can donate 1 electron. O has a charge of 2 electrons, so it can donate two electrons. However, unlike the first method, C5H5 now has a donation of 5 electrons. A metal bonded to another metal counts as one additional electron per metal. The reason for having 18 electrons is due to the complete filling of the electron shell of the transition metal (s2p6d10). Square planar complexes, have both sigma donor and pi acceptor characters, this becomes an exception to the 18 electron rule where a 16 electron configuration is more stable.

Carbonyl
Hundreds of ligands are known to bond to metal atoms through carbon. Carbon monoxide is one of the most common ligand in organometallic chemistry. It is the only ligand in binary carbonyls. CO has some interesting features. The molecular orbital picture of CO is similar to that of N2. First, the highest energy occupied orbital has its largest lobe on the carbon atom. Because this orbital is occupied, the CO gains a sigma donor ability. CO also has two empty pi orbitals, which gives it its ability for pi bonding. This can be confirmed by x-ray crystallography and infrared spectroscopy. First, any change in the bonding between oxygen and carbon can be seen by the vibration by IR. Second, x-ray crystallography measured the distance between C-O bonds at 112.8pm. Weaking of the CO bond by the factors described above would be expected to cause this distance to increase. CO can also act as a bridging ligand. It bridges two metal atoms, both metals can contribute electron density into pi orbitals of the CO to weaken CO bond and lower the energy. Some common reactions that occur with the carbonyl ligand is:

1) Direct reaction of a transition metal with CO. Most of these involve the metal nickel, which interacts in Ni + 4CO -> Ni(CO)4

2) Reductive carbonylation is the reudction of a metal compound in the presence of CO and an appropriate reducing agent. Ex: CrCl3 + 6CO + Al -> Cr(CO)6 + ALCl3.

3) Thermal/photochemical reaction of other binary carbonyls. The most common reaction of carbonyl is CO dissociation. This must be started by a thermal or photochemical reagent. Ex: Cr(CO)6 + PPh3 -> Cr(CO)5(PPh3) + CO

Carbene
Carbene complexes contain metal-carbon double bonds. It was first created in 1964 by Fischer. They typically contain one or two highly electronegative heteroatoms like O, N and S that are directly attached to the carbene. These are called Fischer type carbene complexes. These O, N, S that are attached to the carbene complexes are more stable than the complexes that lack this atom. The stability of this complex is enchanced by the highly electronegative atom that results in pi bonding. When using NMR, it can be seen that this complex is highly dependent on temperature. When the temperature is lowered, it splits into several peaks.