Does the distance increase as the polarity increases?
I'm not sure if the guideline actually works out in empirical observations. There are at least two factors affecting bond length - atomic radii - and polarity of the bond. A more polar bond will have a less even distribution of electrons. The electron density will generally be closer to the more electronegative atom. This isolation of electrons may have the effect of "stretching" the bond. 

Consider the following two molecules, both of which are linear and polar:





From polarity considerations alone, we might expect the $\ce{H-F}$ bond to be longer than the $\ce{H-Cl}$ bond because fluorine is more electronegative than chlorine. However, chlorine is a bigger atom than fluorine by nature of having additional sublevels of electrons. 

Also, consider the below data: bond length actually increases with decreasing polarity (quantified through dipole moments). 

Fluoromethane Bond length 136pm, Bond strength 108 kcal/mol, Dipole moment (D) 1.85

Bromomethane Bond length 178pm, Bond strength 84 kcal/mol, Dipole moment (D) 1.87

Chloromethane Bond length 193pm, Bond strength 70 kcal/mol, Dipole moment (D) 1.81

Iodomethane Bond length 214pm, Bond strength 56 kcal/mol, Dipole moment (D) 1.62

The unfortunate thing about your guideline is that I thought that it would work at least across periods - in which atomic size stays relatively constant. However, the facts do not fit the case. Consider $\ce{H-N}$ and $\ce{H-O}$. Oxygen is more electronegative than nitrogen; we'd expect the $\ce{H-O}$ bond to be more polar, and thus longer than the $\ce{H-N}$, at least according to your guideline. The opposite is actually true; the $\ce{H-O}$ bond is actually shorter than the $\ce{H-N}$ bond.