<P> The bond length between the carbon atom and the oxygen atom is 112.8 pm . This bond length is consistent with a triple bond, as in molecular nitrogen (N), which has a similar bond length (109.76 pm) and nearly the same molecular mass . Carbon--oxygen double bonds are significantly longer, 120.8 pm in formaldehyde, for example . The boiling point (82 K) and melting point (68 K) are very similar to those of N (77 K and 63 K, respectively). The bond - dissociation energy of 1072 kJ / mol is stronger than that of N (942 kJ / mol) and represents the strongest chemical bond known . </P> <P> The ground electronic state of carbon monoxide is a singlet state since there are no unpaired electrons . </P> <P> Carbon and oxygen together have a total of 10 electrons in the valence shell . Following the octet rule for both carbon and oxygen, 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 . </P> <P> 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 . </P>

Why does carbon monoxide form a triple bond
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