# Advanced Inorganic Chemistry/Dative ligands: CO and phosphines
## Dative Ligands: CO and Phosphines:

Dative ligands represent a class of compounds that form dative covalent bonds, otherwise known as coordinate bonds, in which both electrons come from the same atom. In the case of transition metals, a dative ligand can form a coordinate bond with a transition metal. The dative ligand can form a sigma bond (σ) with the metal center and can donate both of its electrons to the metal, contributing to a transition metal complex. For transition metals, a dative ligand will always first form a sigma bond with the metal center. The pi* (π*) antibonding orbitals of the dative ligand, if energetically available, can also interact with the filled d orbitals of the transition metal. This interaction is known as pi backbonding and results in the transfer of 2 electrons from the d orbital of the metal to the pi* antibonding orbital, thus relieving negative charge from the metal center. This represents a transfer of electrons from an atomic orbital (d orbital of the transition metal) to a molecular orbital (pi* antibonding orbital). The d orbital and the pi* orbital have the requisite symmetry with one another to facilitate this process.

## CO Ligands:
### Backbonding and IR Absorption

CO ligands are carbon monoxide ligands that participate readily in transition metal complexes via coordinate bonds with a transition metal center. CO ligands are neutral 2 electron donors and first form a sigma bond with a metal center. As discussed above, the d orbitals of the transition metal are symmetric about the pi* orbitals of the CO compound and backbonding occurs between the metal and the CO. This process is illustrated in Figure 1. Pi backdonation is supported by IR (infrared spectroscopy) data of various metal-carbonyl complexes. Using group theory, one can consider the consequences of backbonding and the affects this would have on the IR absorption of a carbonyl-containing compound. Pi backbonding donates electrons from the transition metal center to the pi* antibonding orbital of CO, the LUMO of this ligand (Lowest Unoccupied Molecular Orbital). This in turn lowers the bond order between the carbon and oxygen, weakening their interaction. This would cause one to suspect that metal carbonyl complexes would absorb light of a lower frequency upon excitation by infrared radiation. This is consistent with experimental data, with few exceptions, for several metal carbonyl complexes when compared to the IR absorption of free carbonyls. This is depicted in Figure 2.
Figure 2: IR Absorption Frequencies for Metal Cabonyls :
| Compound | ν$_{CO}$ (cm$^{-1}$) |
| CO | 2143 |
| V(CO)$_6$$^{1-}$ | 1859 |
| Cr(CO)$_6$ | 2000 |
| Mn(CO)$_6$$^+$ | 2100 |
| Fe(CO)$_6$$^{2+}$ | 2204 |
| Fe(CO)$_5$ | 2022, 2000 |
| Ru(CO)$_5$ | 2038, 2022 |

### X Ray Crystallography

Homoleptic carbonyl complexes typically form octahedral complexes and the normal modes of vibration transform as a$_{1g}$, e$_g$, t$_{1u}$. The character table of the octahedral complex, however, predicts that only the t$_{1u}$ set will be IR active. In this case, the e$_g$ set will be active in
Raman spectroscopy. For metal-carbonyl complexes with substituted ligands, the IR absorptions become more numerous due to the lower symmetry of the each complex. A schematic of this is illustrated in Figure 3. Although IR characterization is a powerful tool to confirm the properties of metal-carbonyl complexes, X-Ray crystallography is another useful method of confirming the existence of pi backbonding. X-ray crystallography can be utilized to confirm the presence of a backbonded metal carbon complex. If a carbon is backbonded to a metal complex then the covalent bond between the two atoms will be stronger and this should indicate a smaller bond distance. X-ray crystallography is capable of determining this bond distance by sending X-rays through the material and measuring the diffraction pattern that is produced to measure the bond distance within a crystal structure.

### Fluxional Carbonyls

Fluxional carbonyls arise from the ability of carbonyl groups within a transition metal complex to move from one metal to another via bridging. This is easily accomplished due to the fact that the ground state energy for a terminal carbonyl group is similar to that of a bridging carbonyl, allowing carbonyls to interchange between metal centers. This typically occurs within first row transition metal complexes and the presence of fluxional complexes can be confirmed in disubstituted metal carbonyl complexes by way of cis-trans isomers. NMR spectroscopy can then be used to determine the conditions upon which all isomers can be converted to either a cis or trans configuration based on the peaks in the 13C NMR spectrum of the molecule. If the trans configuration produces chemically distinct environments for two carbonyls then this will be demonstrated in the NMR spectrum, and the reverse case is true for the cis configuration in which carbons in chemically equivalent environments will express only one peak.

## Phosphine Ligands:
### Phosphine Backbonding

PR3 ligands are phosphine ligands that behave quite similarly to carbon monoxide ligands when interacting with a transition metal center. Phosphine ligands are also neutral and also contribu te to 2 electrons to the transition metal center. Similar to CO ligands, PR3 ligands first donate a sigma bond to the transition metal center, and then the d orbitals of the metal pi backbond with the phosphine and donate 2 electrons to the sigma* antibonding orbital. A schematic of this is shown in figure 4. Although a phosphine ligand has the potential to have large steric groups that impede the formation of transition metal complexes or decrease the bond enthalpy of these coordination bonds, the charged nature of some groups attached to the phosphorus can have significant impacts on the stability of the structure. The electronegativity of the groups attached to the phosphine play a huge role in the stability of the transition metal complex. Strongly electronegative groups on the phosphorus atom can drastically lower the energy of the of the sigma* orbital, thereby making the phosphine a significantly better pi-acceptor (more accessible energy level for donation by the metal). In descending sigma* antibonding orbital energy, the following is obtained:
P-C > P-N > P-O > P-F

### Phosphine Bonding and IR Absorption

While electronegativity of the groups attached to the phosphorus are important in the stability of the transition metal complex, several other factors can influence this stability. Namely, steric hinderance can play enormous role in the ability of a phosphine group to coordinate with a transition metal center. The following is a comparison for several phosphine complexes in descending order from the ligand with the most pi-acceptor character.
PF3~CO > PCL3 > P(OMe)3 > PPh3 > Py
This demonstrates that the more electronegative phosphine groups have higher stability as a result of having greater pi-acceptor character (reduction in energy of sigma* antibonding orbital). They are also smaller allowing them to interact more closely with the metal center as opposed to phosphine with phenyl groups attached. Additionally, these groups, as a consequence of their higher stability, will absorb higher frequencies of IR radiation.

### Cone Angle and Percent Buried Volume

In order to quantify the affects of steric hinderance, cone angle is often used as a metric to define how readily a phosphine ligand or any ligand will dissociate from a transition metal center. For example, a phosphine ligand such as triphenylphosphine will most likely have a larger cone angle than PF3 and will most likely dissociate much more easily than PF3 due to the large size of the three phenyl groups attached to the phosphorus. The cone angle is a rough measure of this property but can be highly inaccurate. This is because a smaller ligand can coordinate more closely to a metal center, thereby exaggerating the size of its relative cone angle. The opposite case is true in that a larger ligand will coordinate further from the metal center, which could underestimate the relative cone angle of this molecule. Therefore, it can be more useful to instead utilize a parameter known as percent buried volume, which is defined as the percent volume occupied by a ligand within a sphere. This is an arbitrary sphere with a defined radius and contains the transition metal at its center. This method is a more useful characterization tool because it does away with the cone angle and considers the spatial volume occupied by the ligand in all dimensions. Inconsistencies between cone angle predictions of dissociation and percent buried volume predictions of dissociation have been documented.

## References

Organometallic HyperTextBook: Ligand Cone Angles, www.ilpi.com/organomet/coneangle.html.
Symmetry, www.chemedx.org/JCEDLib/QBank/collection/ConcepTests/inorsym.html.
Libretexts. “Carbon Monoxide and Backbonding.” Chemistry LibreTexts, Libretexts, 5 June 2019, chem.libretexts.org/Courses/Johns_Hopkins_University/030.356_Advanced_Inorganic_Laboratory/Lab_BCD:_Four_Coordinate_Nickel_Complexes:_Ligand_Effects_and_Organometallic_Catalysis/3_Carbon_Monoxide.
“X-Ray Diffraction.” Rigaku, www.rigaku.com/en/techniques/xrd.
* ↑ Organometallic HyperTextBook: Ligand Cone Angles, www.ilpi.com/organomet/coneangle.html.
* ↑ Symmetry, www.chemedx.org/JCEDLib/QBank/collection/ConcepTests/inorsym.html.
* ↑ Libretexts. “Carbon Monoxide and Backbonding.” Chemistry LibreTexts, Libretexts, 5 June 2019, chem.libretexts.org/Courses/Johns_Hopkins_University/030.356_Advanced_Inorganic_Laboratory/Lab_BCD:_Four_Coordinate_Nickel_Complexes:_Ligand_Effects_and_Organometallic_Catalysis/3_Carbon_Monoxide.
* ↑ “X-Ray Diffraction.” Rigaku, www.rigaku.com/en/techniques/xrd.