Does the dipole moment increase or decrease by increasing the bond length?

Question

It is established the dipole moment is a result of multiplication of the magnitude of charges (Q) and the distance between them (r). What I understand is that when an electron and a proton get closer (bond length decreases), polarity decreases hence the dipole decreases. When they move further away from each other(bond length increases), the polarity increases and hence the dipole moment increases. However, in case of halides (H-X), let's take for example HF and HI, HF has a bigger dipole moment than HI, it is stated that the bond length of HF is smaller than HI. Why do I find this contradictory? Can you help me with What I am missing here?

Answer 1

The piece you are missing is the fact that these aren't ions, the bond is covalent shared electrons. The much larger electronegativity difference in H-F causes the charges to be separated more by the more uneven distribution of the electrons. H electronegativity is 2.1 vs F 4.0, large difference = highly polar bond (remember what bonds can participate in Hydrogen Bonding as an intermolecular force??) So there is a significant charge separation in the H-F bond, it is a very highly polar bond, almost approaching ionic character!

But with H-I, the electronegativity difference is 2.1 vs 2.5 = 0.4, barely polar at all. So it's not just about bond length, you have to also look at how much the charge is actually "separated" in that bonding situation.

Heck, think about an iodine molecule - pretty high bond length, right? But it is NONPOLAR - there is NO charge separation, no net dipole, because the atoms have the same electronegativity! Hope that helps...