Lead(II) iodide is sparingly soluble in water at room temperature $(\pu{0.76 g/L}$ at $\pu{20 °C})$ and a bit more soluble in hot water $(\pu{4.1 g/L}$ at $\pu{100 °C}).$ In solution, all of the lead iodide is dissociated into ions. At high concentrations of iodide, the lead ions form two complexes, $\ce{PbI3-}$ and $\ce{PbI4^2-}$ [1]:
The increase of solubility of lead iodide caused by the presence of iodide ion in concentration greater than 0.1 molal may be explained by the formation of the complex ions $\ce{PbI3-}$ and $\ce{PbI4^=}.$
In any case, all of the dissolved lead iodide is present as ions (there is no neutral $\ce{PbI2(aq)}$ species), so it is a strong electrolyte. This is different from weak electrolytes such as, say, acetic acid, where there is a neutral species $\ce{CH3COOH(aq)}$ in solution which dissociates into the charged $\ce{CH3COO-(aq)}$ and $\ce{H+(aq)}$.
As an aside, the difference in color (colorless in solution, yellow as a solid) and the temperature-dependence of the solubility leads to a beautiful effect that has been called toxic golden rain, see e.g. video YouTube — Golden Rain - Growing crystals of lead iodide. Chemical reaction.
Reference
- Lanford, O. E.; Kiehl, S. J. The Solubility of Lead Iodide in Solutions of Potassium Iodide-Complex Lead Iodide Ions. J. Am. Chem. Soc. 1941, 63 (3), 667–669. DOI: 10.1021/ja01848a010.