Metal Carbonyls

Metal carbonyls are a fascinating class of coordination compounds in which carbon monoxide (CO) acts as a ligand bonded directly to a transition metal. Classic examples include tetracarbonylnickel(0), Ni(CO)₄, and pentacarbonyliron(0), Fe(CO)₅. What makes these compounds extraordinary is that CO — a molecule we normally associate with toxicity and combustion — becomes a remarkably stable, tightly bound ligand when coordinated to metals such as nickel, iron, or chromium. Understanding why requires exploring one of the most elegant bonding concepts in inorganic chemistry: synergic (or synergistic) bonding, which combines two simultaneous but opposite flows of electron density between the metal and the ligand. In most coordination compounds, the ligand donates a lone pair of electrons into an empty orbital on the metal centre — a straightforward Lewis acid–base interaction known as σ (sigma) donation. However, CO possesses low-lying empty π (pi-star) antibonding molecular orbitals that are perfectly suited to accept electron density back from filled d-orbitals on the metal. This second component, called π back-donation (or back-bonding), flows in the reverse direction: from the metal to the ligand. The two effects reinforce each other — the more the metal accepts σ electron density, the more positive charge builds on the metal, encouraging it to donate back via π bonding. This mutual reinforcement is why the bonding is described as synergic. Synergic bonding in metal carbonyls has profound practical consequences. The C–O bond order decreases as electron density from the metal is placed into the CO π orbital, which weakens and lengthens the C–O bond relative to free CO. This weakening can be directly measured using infrared (IR) spectroscopy: free CO absorbs at 2143 cm⁻¹, whereas the CO stretching frequency in metal carbonyls is typically reduced to 1850–2125 cm⁻¹, providing chemists with a sensitive, quantitative probe of the degree of back-bonding. Metal carbonyls are also industrially significant — Ni(CO)₄ is used in the Mond process for purifying nickel, and Fe(CO)₅ serves as a precursor for the synthesis of iron nanoparticles and other organometallic reagents.

Key takeaways
  • The Molecular Orbitals of Free CO
  • σ Donation: CO → Metal
  • π Back-Donation: Metal → CO
  • The Synergic Effect and Bond Order Changes
  • Infrared Evidence and the 18-Electron Rule

The Molecular Orbitals of Free CO

Carbon monoxide is isoelectronic with N₂. Its highest occupied molecular orbital (HOMO) is the 5σ orbital, which is a lone pair located mainly on carbon (slightly antibonding in CO itself). Its lowest unoccupied molecular orbital (LUMO) consists of two degenerate 2π antibonding orbitals oriented perpendicular to the C–O axis. These two frontier orbitals are the key players in synergic bonding. The 5σ HOMO has the right symmetry and energy to overlap with empty metal d-orbitals or s-orbitals of appropriate symmetry (dz² or s), while the 2π LUMOs have the correct symmetry to overlap with filled metal dxz and dyz orbitals.

σ Donation: CO → Metal

In the σ component, the carbon end of CO uses its 5σ lone pair to donate electron density into an empty hybrid orbital (typically of dz² or s character) on the metal. CO always binds through carbon, not oxygen, because the HOMO (5σ) has greater electron density and better directionality at carbon. This donation creates a metal–carbon σ bond and formally increases the electron count at the metal centre. On its own, σ donation would build up negative charge on the metal, making further donation less favourable — this is where back-donation becomes essential.

π Back-Donation: Metal → CO

In the π component, filled metal d-orbitals of t₂g symmetry (dxz, dyz in an octahedral complex) donate electron density into the empty 2π* antibonding orbitals of CO. This flow of electrons is perpendicular to the M–C axis and constitutes a dative π bond from the metal to the ligand. Back-donation relieves the excess negative charge that σ donation placed on the metal, simultaneously strengthening the M–C bond (by adding π character) and weakening the C–O bond (by populating its antibonding orbital). The result is a shorter M–C bond and a longer, weaker C–O bond compared to what pure σ donation alone would predict.

The Synergic Effect and Bond Order Changes

The two effects are self-reinforcing: σ donation makes the metal electron-poor (more Lewis acidic), which enhances its tendency to accept even more σ donation while simultaneously making it more willing to engage in π back-donation to relieve charge. Conversely, π back-donation makes the metal less electron-rich, encouraging more σ donation from CO. This cycle is the essence of synergic bonding. The net result: the M–C bond has partial double-bond character (σ + π), and the C–O bond order drops below 3 (free CO has a formal triple bond, C≡O). In heavily back-donating systems — such as the anionic complex [V(CO)₆]⁻ — the C–O stretching frequency can drop below 1900 cm⁻¹, indicating significant population of the π* orbital and substantial reduction of C–O bond order.

Infrared Evidence and the 18-Electron Rule

Infrared spectroscopy is the primary diagnostic tool for characterising metal carbonyls. Because the C–O stretch is an intense, easily detected band, its frequency directly reports the extent of back-bonding. More electron-rich metal centres (lower oxidation states, more electron-donating co-ligands, or anionic complexes) show lower ν(CO) values. The 18-electron rule also underpins the stability of many metal carbonyls: Ni(CO)₄ (d¹⁰, 4×2 = 8 electrons from CO → 18 total), Fe(CO)₅ (d⁸, 5×2 = 10 electrons from CO → 18 total), and Cr(CO)₆ (d⁶, 6×2 = 12 electrons from CO → 18 total) all satisfy the 18-electron count, explaining their exceptional stability among homoleptic carbonyls.

The Key Equation

\ceNi(s)+4CO(g)−>Ni(CO)4(g)\ce{Ni(s) + 4CO(g) -> Ni(CO)4(g)}
  • \ce{Ni} — Nickel (—): Transition metal in oxidation state 0; d¹⁰ configuration provides filled dxz and dyz orbitals ideal for π back-donation into CO π* orbitals.
  • \ce{CO} — Carbon monoxide ligand (—): Acts as both σ donor (via 5σ lone pair on C) and π acceptor (via empty 2π* orbital); binds through the carbon end to the metal.
  • \ce{Ni(CO)4} — Tetracarbonylnickel(0) (—): Tetrahedral 18-electron complex; C–O stretching frequency at ~2057 cm⁻¹ (compared to 2143 cm⁻¹ for free CO), confirming back-bonding.

Worked Example

Problem. The C–O stretching frequency is measured in three species: free CO (2143 cm⁻¹), Cr(CO)₆ (2000 cm⁻¹), and [V(CO)₆]⁻ (1859 cm⁻¹). Explain the trend in C–O stretching frequencies in terms of synergic bonding, and predict whether the M–C bond length increases or decreases from Cr(CO)₆ to [V(CO)₆]⁻.

ν(\ceCO):free CO>\ceCr(CO)6>[\ceV(CO)6]−\nu(\ce{CO}): \text{free CO} > \ce{Cr(CO)6} > [\ce{V(CO)6}]^{-}

Answer. The C–O frequency decreases because greater π back-donation in [V(CO)₆]⁻ (driven by the extra negative charge making vanadium more electron-rich) populates the CO π* orbital more extensively, weakening the C–O bond. Simultaneously, the M–C bond in [V(CO)₆]⁻ is shorter than in Cr(CO)₆ due to increased M–C π character.

Figure
Diagram
Synergic Bonding in Metal Carbonyls: σ Donation and π Back-DonationMd-orbitalsCOC≡O triple bondσ donation(5σ lone pair, C→M)π back-donation(filled d → CO π*, M→C)Metal d-orbitaldxz (filled)CO MO Diagram5σ (HOMO)2π* (LUMO)(degenerate)Effect on Bond ParametersFree CO: ν = 2143 cm⁻¹Cr(CO)₆: ν = 2000 cm⁻¹[V(CO)₆]⁻: ν = 1859 cm⁻¹Bond order = 3Back-donation moderateBack-donation strongestincreasing back-donation →
Origins · 1 minLudwig Mond and how this was discovered
LM
Ludwig Mond
1839–1909 · German-British

Ludwig Mond discovered tetracarbonylnickel, Ni(CO)₄, in 1890 while investigating the corrosion of nickel valves in his chemical plant. He observed that CO gas at room temperature reacted with nickel metal to form a volatile compound. Together with his assistant Carl Langer, Mond demonstrated that Ni(CO)₄ could be thermally decomposed at higher temperatures to deposit pure nickel — the basis of the Mond process still used industrially for nickel refining. His discovery was the first metal carbonyl ever identified and opened an entirely new field of organometallic chemistry. Mond reportedly almost poisoned himself during early experiments with Ni(CO)₄, which is one of the most acutely toxic substances known — it is both a potent carcinogen and rapidly fatal at very low concentrations because Ni(CO)₄ penetrates biological membranes and releases CO and reactive nickel inside cells.

  • Ligand Field Theory — Ligand field theory describes how d-orbital energies split in coordination complexes; understanding t₂g orbital occupancy is essential for analysing which d-electrons participate in π back-donation in metal carbonyls.
  • The 18-Electron Rule — Most stable homoleptic metal carbonyls (Ni(CO)₄, Fe(CO)₅, Cr(CO)₆) obey the 18-electron rule; understanding electron counting explains why these specific stoichiometries are observed.
  • IR Spectroscopy of Coordination Compounds — The C–O stretching frequency measured by IR spectroscopy is the primary experimental evidence for back-bonding; learning to interpret ν(CO) values quantifies the extent of synergic bonding.
  • Oxidation States of Transition Metals — Metal carbonyls typically feature metals in zero or negative oxidation states; lower oxidation states mean more electron-rich d-orbitals and stronger π back-donation.
  • Molecular Orbital Theory — Synergic bonding is understood through the frontier molecular orbitals of CO (5σ HOMO and 2π* LUMO); a firm grasp of MO theory is a prerequisite for understanding back-bonding.