Why is Hg2Cl2 less soluble in water than HgCl2?

Question

I learned that $\ce{Hg2Cl2}$ is almost insoluble to hot water and $\ce{NH3}$ water in my textbook, but $\ce{HgCl2}$ dissolves well in water.

I wanted to know the reason and searched for it. It's written on some website that thanks to low shielding effect of f-orbitals and the fact that Hg has full f-orbitals, Hg has a high effective nuclear charge. And this makes it possible for the $\ce{Hg+}$ ion to form covalent metal–metal bonds. I searched further and found some information that bond between $\ce{[Hg-Hg]^2+}$ and $\ce{Cl-}$ is more ionic than the bond between $\ce{Hg^2+}$ and $\ce{Cl-}$, according to Fajans' rules.

I thought if the $\ce{[Hg-Hg]^2+}$ and $\ce{Cl-}$ bond is more ionic than the bond between $\ce{Hg^2+}$ and $\ce{Cl-}$, then $\ce{Hg2Cl2}$ would be more likely to dissolve in water because water is a polar solvent. Why is this not the case? Is the $\ce{HgCl2}$ covalent bond so weak that it will dissociate in water? Why?

I came across the theory of negative and positive hydration [1]. Does anyone think this theory could explain the difference of solubility of these two salts?

Reference

1. Rodnikova, M. N. Negative Hydration of Ions. Russian Journal of Electrochemistry 2003, 39 (2), 192–197. DOI: 10.1023/A:1022317227140.

Answer 1

According to Fajan's rules, due to more polarising power of small cation $\ce{Hg^2+}$, the $\ce{HgCl2}$ bonds will be more covalent than those in $\ce{Hg2Cl2}$. But more ionic character of a bond does not imply more solubility.
When you put a compound, say $\ce{NaCl}$ in water (a polar solvent), the ions get solvated. Now as $\ce{Na+}$ ion is small, it has high charge density and more of water molecules surround it, increasing the stability (due to decrease in potential energy), hence higher will be the hydration enthalpy, which breaks the ionic lattice of $\ce{NaCl}$. As ionic compounds arrange themselves into lattice structure, also taking into account the coulombic force of attraction, generally ionic bond is stronger than covalent bond, especially when looking into pair of atoms forming different kinds of compounds.
Similarly, $\ce{Hg^2+}$ has higher charge density so it will realtively be more hydrated with respect to $\ce{Hg2^2+}$, also due to more covalent character of $\ce{Hg-Cl}$ bond in $\ce{HgCl2}$, its lattice would somewhat be weaker, lacking the polarity ionic compounds have. So it can qualitatively be assumed for $\ce{HgCl2}$ to be more soluble in water than $\ce{Hg2Cl2}$.
Also, as pointed by Maurice, calculating, or even comparing solubility remains much of a mystery, but many of times, in common compounds, general trends and classical theories can help explain the solubilities.