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

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?

• First part of your first statement? Sep 19 '17 at 19:44
• If you relate and compare two properties (here bond lenght and dipole moment) quantitatively, you have to look if other relevant properties (here the effective charge!) are changing. Which they are.
– Karl
Sep 19 '17 at 19:55
• Sep 20 '17 at 0:18

Strictly speaking, you can neither decrease nor increase the bond length a slightest bit.$\dagger$ All molecules of a given compound are equal. It is not like you may grab one with tiny pincers of some sort, stretch it a little and see what happens.

You may, however, consider a molecule of a different compound; indeed, that's what you do in your question. But that's not just a different bond length; it is different atoms with different properties, and hence different everything.

Most importantly for the question at hand, iodine has much lower electronegativity than fluorine, so the effective charges on $\rm H$ and $\rm I$ in $\ce{HI}$ will be much lower than those on $\rm H$ and $\rm F$ in $\ce{HF}$. It just turns out that the increase in bond length can't compensate for this effect, hence the lower dipole moment.

$\dagger$ For those who might want to point out possible vibrational excitation or some other subtle (or not-so-subtle) effect: don't. Let's keep it simple.

• Point taken for no vibrations, but what about going from e.g. $\ce{CH3-NH2}$ to $\ce{CH2=NH}$?
– Jan
Sep 20 '17 at 3:17
• @Jan Just like I said, these are different molecules with different everything. Sep 20 '17 at 4:38

If the same point charges are separated by a larger distance, that's a higher dipole moment.

In the real world, you have to regard the distance and the actual charge difference and distribution.

In the case of HF and HI, fluorine has the highest electronegativity, which means the bond in HF is more polarised than in HI. So, in spite of the lower bond lenght, HF can have, and indeed has, a higher dipole moment.

• but why distance isnt taking into consideration ? Jan 10 '19 at 21:09
• @adosar please explain what you mean by "consideration". I don't get it.
– Karl
Jan 10 '19 at 21:15
• Why hf has greater dipole moment if hcl bond length is bigger ? How can we know that q1d1>q2d2 only from the electronegativity difference? Jan 10 '19 at 23:47
• @adosar I explain exactly that above. It's a fact, we know the distance and the dipole moment from different measurements, so we can calculate the charge difference! And we know about the high electronegativity of fluorine, so the result makes sense. That's nature, I can't change it.
– Karl
Jan 11 '19 at 6:14