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Repulsion between shared electron pairs shapes the molecule in such a way that enables the pairs to be as far apart as possible. Furthermore, the repulsion between shared and unshared pairs forces bonded pairs closer to one another. These principles of the VSEPR theory and orbital hybridization, in addition to Van der Waals forces, explain the overall shape of a molecule.

In my textbook, however, there was little mention to how polar bonds and the partial charges they cause affect the structure of a molecule. I would expect that poles of the same charge would repel each other and have a profound affect on molecular shape. However, many compounds make me second guess this assumption, including $\ce{H2O}$.

Do polar bonds at all affect molecular shape? Why or why not? If so, to what extent?

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    $\begingroup$ If I understand correctly, you are saying that by this logic water should be linear like CO2 for example? If this is the case, think about other forces at play here. Drawing Lewis dot structures of water, and CO2 for my example of a linear molecule, might illustrate why H2O is not linear and CO2 is. Let me know if I'm completely off base regarding your question and I'll just delete this ;) $\endgroup$ – airhuff Feb 3 '17 at 22:38
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Yes, of course they affect the shape. But only slightly so. More precisely, bond angles can change, however this does not change the general shape predicted by VSEPR.

For example, if two atoms (one with a positive pole, and one with a negative pole) are connected by another atom, then the two atoms would make an angle that is smaller than the angle predicted by VSEPR.

And also, polar bonds are capable of intramolecular bonding. Such as in protein-folding. So they definitely contribute to molecular shape.

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