Why the "Diverse Ion Effect" increases Ksp?
To illustrate the "diverse ion effect", we will study how the solubility of $\ce{AgCl}$ changes in absence and presence of $\ce{NaNO3}$, given that the molar solubility product of $\ce{AgCl}$ is $1.76 \times 10^{-10}$.

The exact definition of the solubility constant is: $$K_{sp}= a(\ce{Ag+}).a(\ce{Cl-})=[\ce{Ag+}].\gamma(\ce{Ag+}).[\ce{Cl-}].\gamma(\ce{Cl-})$$
Where $a(\ce{Ag+})$ and  $a(\ce{Cl-})$ are the activities of these ions, and $\gamma(\ce{Ag+})$ and $\gamma(\ce{Cl-})$


In dilute solutions, i.e. absence of $\ce{NaNO3}$:  $$\gamma(\ce{Ag+})=\gamma(\ce{Cl-})=1$$
So, the solubility of $\ce{AgCl}$ is equal to $$[\ce{Ag+}]= \sqrt{K_{sp}}= 1.33 \times 10^{-5}$$
In presence of $\ce{NaNO3}$ of a concentration of $0.1 \mathrm{mol/kg}$, then
$$[\ce{Ag+}]= \sqrt{K_{sp}}/ \sqrt{(\gamma(\ce{Na+}).\gamma(\ce{NO3-})}
= 1.33 \times 10^{-5}/ (0.769)= 1.73 \times 10^{-5} $$
i.e. in the presence of a soluble salt, the molar solubility of a slightly soluble salt increases.