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Why is $\ce{CF2Cl2}$ polar? The Lewis dot structure shows the molecule as being symmetric. Wouldn't that make it nonpolar?

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closed as off-topic by Jannis Andreska, Pritt Balagopal, bon, M.A.R., Jan Oct 19 '17 at 12:12

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    $\begingroup$ The Lewis dot structure doesn't show a thing. It is drawn on a planar sheet of paper, and the real molecule is 3D. Ever heard of a tetrahedral geometry? $\endgroup$ – Ivan Neretin Oct 11 '17 at 4:28
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    $\begingroup$ I encourage you to make that as an answer @IvanNeretin. It seems to be a precise and complete answer to me. $\endgroup$ – Pritt Balagopal Oct 19 '17 at 5:53
  • $\begingroup$ @PrittBalagopal Maybe, but it is no improvement over the existing answer, hence no. $\endgroup$ – Ivan Neretin Oct 19 '17 at 6:23
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The problem is that the first Lewis diagrams you learn to draw don't take 3D structure into account. A more realistic structure for dichlorodifluoromethane looks like

where the dashed wedged bond is meant to represent a single bond pointing away from you ("into" the plane of the screen), and the filled bold wedged bond is meant to represent a single bond pointing toward you (coming out of the screen). This is more than a traditional Lewis structure though, which is more about notating single/double/triple bonds and any nonbonding electrons.

You could also deduce the 3D structure from the fact that carbon centers with 4 separate bonds are always tetrahedral, and there is no way in a perfect tetrahedron to place either the 2 chlorines or the 2 fluorines exactly opposite each other. This means there is no way for the individual bond dipoles to cancel.

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