While googling some numbers, I found that the dipole moment of acetonitrile is, according to wikipedia 3.92D. Compare this to the dipole moment of ethylamine, according to NIST of 1.220D. Now, I've read elsewhere (and perhaps this explanation is wrong) that polar molecules with pi bonds tend to be more polar than counterpart molecules with only sigma bonds, because there is an overall greater displacement of electron density towards the more electronegative element. In other words, more electrons = more unequal sharing = more polarity.
While this is a quick explanation for this difference, it leaves me scratching my head a little. The difference in electronegativities of carbon and nitrogen is ~0.49 (weakly polar), which in my imagination isn't nearly enough to justify the huge differences between dipole moments of ethylamine and acetonitrile. Do two pi bonds really result in a nearly 3 fold increase in dipole moment?
Compare this with the behavior of carbon with a much more electronegative element, oxygen. The dipole moment of methanol is 1.69D while the dipole moment of carbon monoxide is 0.122D. The more electronegative element produces a smaller dipole in a molecule with greater electron density between the carbon and oxygen, relative to a similar molecule with less electron density. In other words, our previous explanation of more electrons = more unequal sharing = more polarity breaks down in this case.
This is the crux of my dilemma. How can a more electronegative element produce a smaller dipole than a less electronegative element? To ask this question another way, how can a less electronegative element produce such a large dipole? Is carbon monoxide a strange case, or have I confused myself (which isn't so uncommon)?
Thanks!
