Metal Carbonyls: Synergic Bonding and What the Spectrum Reveals

The bonding, in two halves

Free carbon monoxide is a poor Lewis base — its lone pair is not especially available. Yet it binds transition metals extremely strongly. The resolution is that two interactions operate together, each reinforcing the other.

• σ donation: the carbon lone pair donates into an empty metal orbital.
• π back-donation: a filled metal d orbital donates back into the empty π* antibonding orbital of CO.

These are synergic: donation puts negative charge on the metal, which makes back-donation more favourable, which removes charge from the metal, which makes further donation more favourable. Each strengthens the other, which is why the overall bond is far stronger than either interaction alone would suggest.

The infrared consequence

Ordering a set of carbonyl complexes by stretching frequency, given their charges or their other ligands, is one of the most common questions in organometallic chemistry — and the reasoning always runs through electron density at the metal.

The same logic applies to co-ligands. A strong π-acceptor elsewhere on the metal competes for the same d electrons, reducing back-donation to CO and raising its frequency. A strong σ-donor co-ligand does the opposite.

Terminal versus bridging carbonyl

A bridging CO shares its π* orbitals with two metals, receiving back-donation from both. The C–O bond is therefore weakened more, and the stretching frequency drops further — typically well below the terminal range.

This makes infrared spectroscopy a direct structural tool: counting bands and noting their positions distinguishes terminal from bridging carbonyls without any other measurement. Deducing a structure from a list of stretching frequencies is a standard exam task.

Synthesis

• Direct combination of the metal with carbon monoxide, which works for a few metals under pressure.
• Reductive carbonylation, in which a metal salt is reduced in the presence of CO.
• Photolysis or thermolysis of a simpler carbonyl to give clusters with metal–metal bonds.

The historically important application is the Mond process, in which nickel forms a volatile carbonyl that is distilled away from impurities and then decomposed to give very pure nickel. It works because the carbonyl forms and decomposes at conveniently different temperatures.

Structure and the 18-electron rule

Most simple carbonyls obey the 18-electron rule, and it predicts their formulas. A metal with an even electron count forms a mononuclear carbonyl; one with an odd count cannot reach eighteen alone, so it either dimerises through a metal–metal bond or takes an extra electron as an anion.

This is why some metals give simple mononuclear carbonyls while their neighbours give dimers — a pattern that looks arbitrary until the electron count is done, and a very common question.

Frequently asked questions

Why does CO bind through carbon rather than oxygen?

Because its highest occupied molecular orbital is concentrated on carbon, despite oxygen being more electronegative. That orbital is what donates to the metal.

Why does a negative charge lower the stretching frequency?

Extra electron density on the metal increases back-donation into the CO antibonding orbital, weakening the C–O bond further and lowering its frequency.

How do I tell bridging from terminal CO?

By the stretching frequency. Bridging carbonyls absorb substantially lower because two metals back-donate into the same antibonding orbital.

Why are metal carbonyls often toxic?

Several are volatile and readily absorbed, and they release carbon monoxide in the body. Nickel tetracarbonyl is notoriously hazardous, which is a practical point worth knowing alongside the chemistry.