What is the order of dipole moments for $\ce{SO3, SiO2, P2O5,}$ and $\ce{Cl2O7}$? This question is from a book and the answer given is $\ce{Cl2O7 < SO3 < P2O5 < SiO2}$. Can someone explain why? When I drew the Lewis structures and applied VSEPR theory, it seemed that all the given compounds have zero dipole moment.

  • $\begingroup$ Can you double check that it says SiO2 has the largest dipole moment? SiO2 should have no dipole moment, it doesn't say SO2 does it? That would have a dipole moment, and...maybe the largest of those listed. $\endgroup$ – J M Aug 8 '12 at 14:14
  • $\begingroup$ Make sure you distinguish between ions and molecules first. Also, don't be shy to look on the internet for the structure of these compounds. My hunch is that you had P2O5 wrong, but I could be mistaken. $\endgroup$ – CHM Aug 8 '12 at 15:50
  • $\begingroup$ @CHM Regardless, something seems off, as SiO2 certainly should not have the largest dipole moment. And if they really did mean the SO3, and not sulfite ion, then that also should have no dipole moment. So, the answer key still needs explaining if that is really what was provided as the correct answer. $\endgroup$ – J M Aug 8 '12 at 18:07
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    $\begingroup$ The answer is really strange. $SO_3$ is plane triangle and has dipole moment zero, $P_2O_5$ is symmetric tetrahedron as dimer and has dipole moment zero, $SiO_2$ is insoluble solid with covalent grid and the only one that can have dipole from this molecules is $Cl_2O_7$, which is asymmetric in O atom as it has structure $ClO_3-O-ClO_3$. The polarisation of bond Element-Oxigen, however, grows in this order, as electronegativity of the element in question falls. $\endgroup$ – permeakra Aug 9 '12 at 5:09
  • $\begingroup$ Something is odd here. Assuming $\ce{P2O5}$ to not be a dimer, then a valid Lewis structure drawn for it of the type $\ce{O2P}-O-\ce{PO2}$, which would be polar similarly to $\ce{Cl2O7}$. I expect the question is about the bond dipole moment of the X-O bond, in which case, the answer given in the solutions manual makes sense. @Tyrell, could you clarify whether the question asked for bond dipole moments or molecular dipole moments. $\endgroup$ – Ben Norris Aug 10 '12 at 1:29

I wondered about that for a long time, but Ben Norris and BigGenius have the truth of it: the question is about the bond dipole moment, otherwise the answer would make sense. Here's why.

Molecular dipoles

  • Sulfur trioxide has planar trigonal geometry, and its molecular dipole is zero. Similarly, molecular $\ce{SiO2}$ has the same geometry as $\ce{CO2}$, i.e. linear, and thus its molecular dipole is zero.

  • $\ce{Cl2O7}$ is bent at its central oxygen atom:


    and thus has a non-zero dipole moment. While $\ce{P2O5}$ generally occurs as its “dimer” $\ce{P4O10}$, it is possible to write a molecular structure for $\ce{P2O5}$ which then strongly resembles that of $\ce{Cl2O7}$, being bent at the central O atom. It thus also has non-zero molecular dipole.

  • Finally, ordering between $\ce{Cl2O7}$ and $\ce{P2O5}$: because the P is less electronegative than Cl, the partial charges borne by the P atoms will be larger than those of the Cl atoms, and thus I expect the molecular dipole of $\ce{P2O5}$ to be greater than that of $\ce{Cl2O7}$.

Bond dipoles

If the question is talking about the $\ce{X–O}$ (or $\ce{X=O}$) bond dipole in the series where X = S, Si, P or Cl, it becomes easier to understand. For one thing, because you're asked about $\ce{X–O}$ bonds where X spans a row of the periodic table, which is typical of this sort of problems. The answer would then be that bond dipole moment get bigger when the electronegativity difference of X and O becomes larger (the sharing of electrons gets more asymmetric). This leads to the answer you gave.


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