Cobalt tetracarbonyl hydride

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Cobalt tetracarbonyl hydride
Cobalt tetracarbonyl hydride.svg
Other names
cobalt hydrocarbonyl
Hydrocobalt tetracarbonyl
16842-03-8 YesY=
PubChem 61848
Molar mass 171.98 g/mol
Appearance Light yellow liquid
Odor offensive[1]
Melting point −33 °C (−27 °F; 240 K)
Boiling point 47 °C (117 °F; 320 K)
0.05% (20°C)[1]
Solubility soluble in hexane, toluene, ethanol
Vapor pressure >1 atm (20°C)[1]
Acidity (pKa) 8.5
Main hazards flammable, decomposes in air[1]
US health exposure limits (NIOSH):
TWA 0.1 mg/m3[1]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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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.[2] Its main use is as a catalyst in hydroformylation.

Structure and properties[edit]


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.[3] The Co-CO and Co-H bond distances were determined by gas-phase electron diffraction to be 1.764 and 1.556 Å, respectively.[4] 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.[5] HCo(CO)4 melts at -33 °C and above that temperature decomposes to Co2(CO)8 and H2.[2] 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.[6] These derivatives are generally less acidic than HCo(CO)4.[5]


Tetracarbonylhydrocobalt was first described by Hieber in the early 1930s.[7] 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.[3]

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.[6]

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.[6]


Tetracarbonylhydridocobalt was the first transition metal hydride to be used in industry.[8] In 1953 evidence was disclosed that it is the active catalyst for the conversion of alkenes, CO, and H2 to aldehydes, a process known as hydroformylation (Oxo Reaction).[9] Although the use of cobalt-based hydroformylation 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.[8]


  1. ^ a b c d e f g "NIOSH Pocket Guide to Chemical Hazards #0148". National Institute for Occupational Safety and Health (NIOSH). 
  2. ^ a b Kerr, W. J. (2001). "Sodium Tetracarbonylcobaltate". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rs105. 
  3. ^ 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. 
  4. ^ 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 (19): 6243. doi:10.1021/ja00461a011. 
  5. ^ 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 (9): 2257–2263. doi:10.1021/ja00269a022. PMID 22175569. 
  6. ^ 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
  7. ^ 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 (7): 1090. doi:10.1002/cber.19320650709. 
  8. ^ a b Rittmeyer, P.; Wietelmann, U. (2000). "Hydrides". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a13_199. 
  9. ^ I. Wender , H. W. Sternberg , M. Orchin "Evidence for Cobalt Hydrocarbonyl as the Hydroformylation Catalyst" J. Am. Chem. Soc., 1953, vol. 75, pp 3041–3042. doi:10.1021/ja01108a528