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Bridging Carbonyl is defined as a carbonyl (Carbon monoxide) that is bound to two or more metals via the carbon, or both the carbon and the oxygen.

Introduction

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In organometallic chemistry the Carbonyl ligand (CO, otherwise know as carbon monoxide) is the one of the more commonly used ligands, as well as one of the more heavily studied ligands. The carbonyl ligand is typically seen as a terminal ligand as in Iron pentacarbonyl. But more interestingly the carbonyl can also be seen as a bridging ligand between two metals[1]. The typical notation for a CO bridging ligand is notated by the Greek letter µ (mu) followed by a subscript that indicates the number of atoms the bridging ligand is bound to. For example µ3-CO represents a carbonyl ligand that is bridging between three metals.


Interaction between a CO ligand and a metal can be either through Sigma Donation or Pi acceptance[1], as shown below. In Sigma donation the highest occupied orbital (HOMO) has its largest lobe on carbon. It is this orbital that CO can exert its sigma donation. The electron pair that sits in that orbital can be donated toward the corresponding metal orbital, like an unfilled p or d hybrid orbital. CO also contains two empty π* orbital’s (LUMO). These orbitals also have larger lobes than that of the oxygen. Due to the vacancy of the two π* orbital’s, the CO can have Pi acceptance from the metal (considering proper symmetry of the metals d orbital’s). CO being both a Sigma donor and Pi acceptor is called synergistic, where the greater the electron density on the metal, the more effectively it is able to back donate to the π* of the CO ligand. This in effect is called backbonding, where this effect can be strong between the metal and the CO forming a rather strong bond between the two[1]. However in most bridging situations the more interactions the carbonyl ligand has with metals, the weaker the C-O bond becomes, where this lower bond strength can be see under Infrared Spectroscopy as explained later.

Purpose of a bridging Carbonyl

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Bridging carbonyls are seen in a multitude of different fields and studies. From industrial applications as seen with the Ru3(CO)12 which is widely used in photo activated synthesis[2], to biological applications like cyclic enkephalin and dermorphin analogues[3] The binding of the carbonyl can play a vital role in reactivity and functionality of many different applications.

Synthesis

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The most common reaction of carbonyl ligands is dissociation[1], . From this reaction, typically under UV light or heat, a carbonyl ligand dissociates off the metal leaving an open coordination site. Depending on the environment around the open coordination site, further reactions can then occur. Iron pentacarbonyl is an example of this and shown below, where after dissociation of a CO with the use of ultraviolet light, the complex can then form the Fe2 complex. If heat is used instead of light, you form the Fe3 binary complex[1], .


Some other synthesis examples of these bridging CO molecules are shown below. The first example produces a thermally unstable µ2-CO adduct [(Cp*Ru)2(µ2-NHPh)2(µ2-CO)]. Despite the instability it is still able to be found using spectroscopic techniques. This structure then spontaneously eliminates 1 equivalent of aniline (PhNH2) to produce the imido complex[4]. After elimination you end with the amido phosphide complex[5].

The other example demonstrates the formation of a tris-carbonyl-bridged binuclear Rh complex after treatment with 1 bar CO[6]. It was seen that without excess CO this binuclear complex could then degrade to the final product after another dissociation of CO, where the stability of the 2+ cationic species was only made possible with an excess of CO.

Characterization

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The carbonyl bridging interaction can be monitored through Infrared Spectroscopy, where depending on the number of metals in which the CO is bound to, can overall affect the frequency of the CO stretching frequency in the IR spectrum. Terminal M-CO have a v(CO) of approximately 1850-2120 (cm-1). While symmetrical µ2-CO are seen between 1700-1860 (cm-1). µ3-CO and µ4-CO were seen between 1600-1700 (cm-1)[7]. A flow chart is shown below to better express this idea. Since there is an overall shift of about 200-300 (cm-1[8]. , it is expected that bridging ligands tend to have a longer bond length than compared to its terminal counterparts[1].

Computer Calculations

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In order to understand the stability of some of the overall carbonyls, B3LYP calculations were run on various types of Fe3COx type isomers[9]. Some examples are shown below.

From the complexes A and B, A was calculated to have a higher energy of about 22.5 kcal/mol (B3LYP)[6]. Showing better stability in the unbridged Fe3(CO)11 complex. However, these complexes are not known experimentally[9] and therefore further calculations with other Fe complexes were done of Fe complexes that are known to exist in the lab.

C, D, and E are structures of Fe3(CO)10 isomers containing a µ3, terminal, and µ2 CO bridging ligand respectively. From the calculations it was seen that structure C or the µ3-CO was found to more energetically favorable[9]. Other test were done with Fe3(CO)9 isomers where structure F (shown below) with 3 µ2- CO was also calculated to have a global minimum compared with G with only one µ2- CO , having a overall lower energy of about 8.1 kcal/mol through use of B3LYP calculations[9].

Other Bridging Modes

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In addition to symmetrical bridging modes of CO, sometimes CO can be found bridging either asymmetrically or through donation from a metal d orbital to the π* orbital of CO[1], as shown below.

Asymmetrically bridging mode

Donation from a metal d orbital to the π* orbital[10]

For the latter case, this type of interaction was seen in this reaction:

[(η⁵−C₅H₅)Mo(CO)₃]₂ ↔ [(η⁵−C₅H₅)Mo(CO)₂]₂ + 2CO
1960, 1915 cm⁻¹ 1889, 1859 cm⁻¹

This reaction is driven by heat, where carbon monoxide is driven off. The product can readily go back to the starting material under reaction with the carbon monoxide. As seen there is an overall shift in infrared bands to lower regions, where the Mo-Mo bond distance was shorten by a distance of 80pm[11][12]. This shift is caused by the bridging mode shown above where donation of the other metal d orbital contributes electron density into the π* of the CO. This donation weakens the CO bond, resulting in this lower energy shift.

Electron Counting

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In terms of electron counting terminal carbonyls are counted as donating two electrons to the metal complex. For bridging carbonyls however the electron count is changed depending on the number of metals the carbonyl is bridged to. An example is shown below for the Fe3(CO)9 complex (f) and (g).

From the electron counting it is seen that for complex F, due to its symmetry, all Fe have the same 16 electron count.Complex G on the other hand contains two Fe centers with a 15 electron count, while only one has a count of 16 electrons. Since most organometallic complexes like to form 16 or 18 electron complexes, this lack of 16 electrons for structure G could be the contributing factor to the higher energy calculation.

See Also

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References

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  1. ^ a b c d e f g G. O. Spessard, G. L. Miessler, Organometallic Chemistry, 2nd ed., Oxford Univeristy Press: New York, 2010, pp. 79-82
  2. ^ Y. Wuu, J. Bentsen, C. Brinkley, and M. Wrighton, Inorg. Chem., 1987, 26, 530.
  3. ^ F. Katarzyna, O. Marta, W. Jacek, N. Chung, P. Schiller, P. Danuta, Z. Agnieszka, P. Agnieszka, W. Ewa, I. Jan,. J. Pept. Sci. 2005, 11(6), 347.
  4. ^ S. Takemoto, T. Kobayashi, H. Matsuzka, J. Am. Chem. Soc., 2004, 126, 10802.
  5. ^ S. Takemoto, Y. Kimura, K. Kamikawa, and H. Matsuzaka, Organometallics., 2008, 27, 1780.
  6. ^ W. I. Dzik, C. Creusen, R. Gelder, T. P. J. Peters, J. M. M Smits, and B. Bruin, Organometallics., 2010, 29, 1629.
  7. ^ P. Li and M. D. Curtis, J. Am. Chem. Soc, 1989, 111, 8279.
  8. ^ J. H. Jang, J. G. Lee, H. Lee, Y. Xie, and H. F. Schaefer III, J. Phys. Chem., 1998, 102, 5298.
  9. ^ a b c d H. Wang, Y. Xie, B. R. King, and H. F. Schaefer III, J. Am. Chem. Soc., 2006, 128, 11376.
  10. ^ A. L. Sargent and M. B. Hall, J. Am. Chem. Soc., 1989, 111, 1563.
  11. ^ D. S. Ginley and M.S. Wrighton, J. Am. Chem., 1975, 97, 3533
  12. ^ R. J. Klingler, W. Butler, and M.D. Curtis, J. Am. Chem. Soc., 1975, 97, 3535.
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