Cobalt tetracarbonyl hydride

Cobalt tetracarbonyl hydride
Cobalt tetracarbonyl hydride
Other names
cobalt hydrocarbonyl
tetracarbonylhydridocobalt, Tetracarbonylhydrocobalt
Identifiers
CAS number 16842-03-8 YesY=
PubChem 61848
Properties
Molecular formula C4HCoO4
Molar mass 171.98 g/mol
Appearance Light yellow liquid
Melting point

-33 °C, 240 K, -27 °F
Boiling point

47 °C, 320 K, 117 °F
Solubility in water Partially
Solubility soluble in hexane, toluene, ethanol
Acidity (pKa) 8.5
 YesY tetracarbonyl hydride (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Cobalt tetracarbonyl hydride is the organometallic compound with the formula HCo(CO)4. It is a yellow liquid that forms a colorless vapor and has an intolerable odor.[1] Its main use is as a catalyst in hydroformylation.

Contents

Structure and properties

HCo(CO)4-3D-balls.png

HCo(CO)4 is a trigonal bipyrimidal molecule. The hydride ligand occupies one of the axial positions, thus the symmetry of the molecule is C3v.[2] The Co-CO and Co-H bond distances were determined by gas-phase electron diffraction to be 1.764 and 1.556 Å, respectively.[3]. The oxidation state of cobalt in this compound is -1.

Like some other metal carbonyl hydrides, HCo(CO)4 is acidic, with a pKa of 8.5.[4] HCo(CO)4 melts at -33 °C and above that temperature decomposes to Co2(CO)8 and H2.[1] It undergoes substitution by tertiary phosphines. For example, triphenylphosphine gives HCo(CO)3PPh3 and HCo(CO)2(PPh3)2. These derivatives are more stable than HCo(CO)4 and are used industrially.[5] These derivatives are generally less acidic than HCo(CO)4.[4]

Preparation

Tetracarbonylhydrocobalt was first described by Hieber in the early 1930s.[6] It was the second transition metal hydride to be discovered, after H2Fe(CO)4. It is prepared by reducing Co2(CO)8 with sodium amalgam or a similar reducing agent followed by acidification.[2]

Co2(CO)8 + 2 Na → 2 NaCo(CO)4
NaCo(CO)4 + H+ → HCo(CO)4 + Na+

Since HCo(CO)4 decomposes so readily, it is usually generated in situ by hydrogenation of Co2(CO)8.[5]

Co2(CO)8 + H2 is in equilibrium with 2 HCo(CO)4

The thermodynamic parameters for the equilibrium reaction were determined by infrared spectroscopy to be ΔH = 4.054 kcal mol−1, ΔS = -3.067 cal mol−1 K−1.[5]

Applications

Tetracarbonylhydridocobalt was the first transition metal hydride to be used in industry.[7] In 1940 it was discovered that it catalyzed the conversion of alkenes, CO, and H2 to aldehydes, a process known as hydroformylation (Oxo Reaction). Although it has since been largely superseded by rhodium-based catalysts, the world output of C3-C18 aldehydes produced by tetracarbonylhydrocobalt catalysis is about 100,000 tons/year, roughly 2% of the total.[7]

References

  1. ^ a b Kerr, W. J. (2001). "Sodium Tetracarbonylcobaltate". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rs105.  edit
  2. ^ a b Donaldson, J. D.; Beyersmann, D. (2005). "Cobalt and Cobalt Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a07_281.pub2.  edit
  3. ^ McNeill, E. A.; Scholer, F. R. (1977). "Molecular structure of the gaseous metal carbonyl hydrides of manganese, iron, and cobalt". Journal of the American Chemical Society 99: 6243. doi:10.1021/ja00461a011.  edit
  4. ^ a b Moore, E. J.; Sullivan, J. M.; Norton, J. R. (1986). "Kinetic and thermodynamic acidity of hydrido transition-metal complexes. 3. Thermodynamic acidity of common mononuclear carbonyl hydrides". Journal of the American Chemical Society 108: 2257. doi:10.1021/ja00269a022.  edit
  5. ^ a b c M. Pfeffer, M. Grellier "Cobalt Organometallics" in Comprehensive Organometallic Chemistry III, 2007, Elsevier.doi:10.1016/B0-08-045047-4/00096-0
  6. ^ Hieber, W.; Mühlbauer, F.; Ehmann, E. A. (1932). "Derivate des Kobalt- und Nickelcarbonyls (XVI. Mitteil. über Metallcarbonyle)". Berichte der deutschen chemischen Gesellschaft (A and B Series) 65: 1090. doi:10.1002/cber.19320650709.  edit
  7. ^ a b Rittmeyer, P.; Wietelmann, U. (2000). "Hydrides". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a13_199.  edit

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