[schafer]

Hey,
I had two questions regarding CO (which probably will lead to further questions).

1. Since carbon usually have 4 bonds (like in the case of CO2), I don't understand how can it bond with only 1 oxygen and still be stable.

2. If carbon is burnt in an environment with limited oxygen, will we receive CO2 and have 'leftover' carbon, or burn all carbon with the product of CO?

Thanks a lot !

[Corribus]

Stable is a relative term, of course, but the C≡O bond is reasonably strong and the molecule is not a free radical, so there's no huge, glaring reason for it to be unstable.

If there is not enough fuel for complete combustion, usually you will get formation of CO as a product. There's a substantial entropic driving force to form gasses.

[critzz]

CO is far less stable then lets say CO₂. This is why CO is harmfull (it binds to hemoglobin in blood faster than O₂) and CO₂ is not.
But I guess in air it's relative stable.

[Corribus]

You cannot generalize in this way. CO is more stable than CO₂ at high temperatures, for example, because the entropy driving force to form CO becomes more favorable than CO₂.

http://en.wikipedia.org/wiki/Boudouard_reaction

In addition, thermodynamic stability does not necessarily translate into kinetic stability.

[billnotgatez]

https://en.wikipedia.org/wiki/Carbon_monoxide

Have we really answered the question why the carbon in this case has 3 bonds rather than 4.

Even with the following excerpt from the above link, I think many would still be confused.

Carbon and oxygen together have a total of 10 valence electrons in carbon monoxide. To satisfy the octet rule for the carbon, the two atoms form a triple bond, with six shared electrons in three bonding molecular orbitals, rather than the usual double bond found in organic carbonyl compounds. Since four of the shared electrons come from the oxygen atom and only two from carbon, one bonding orbital is occupied by two electrons from oxygen, forming a dative or dipolar bond. This causes a C ← O polarization of the molecule, with a small negative charge on carbon and a small positive charge on oxygen. The other two bonding orbitals are each occupied by one electron from carbon and one from oxygen, forming (polar) covalent bonds with a reverse C → O polarization, since oxygen is more electronegative than carbon. In the free carbon monoxide, a net negative charge δ- remains at the carbon end and the molecule has a small dipole moment of 0.122 D.[18]

The molecule is therefore asymmetric: oxygen has more electron density than carbon, and is also slightly positively charged compared to carbon being negative. By contrast, the isoelectronic dinitrogen molecule has no dipole moment.

If carbon monoxide acts as a ligand, the polarity of the dipole may reverse with a net negative charge on the oxygen end, depending on the structure of the coordination complex.[19] See also the section "Coordination chemistry" below.