Use formal charge arguments to explain why co has a much smaller dipole moment than would be expectedon the basis of electronegativity.

Explanation:

Analizing only the difference of electronegativity between C and O the dipole moment is important. But why is it smaller than expected?

In the figure we can see the Lewis structure of the CO applying Formal charge, in which the oxygen shares one pair of lone electrons with the carbon in addition to the double bond they have, resulting on a triple bond (following the formal charge arguments).

This acummulation of electrons in the triple decreases the effect of electronegativity anf therefore the magnitude of the dipole