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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.
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  • $\begingroup$ Ionic nature might not always help in checking solubility. A compound is soluble if its hydration enthalpy overcomes the lattice enthalpy. The linear structure of Hg2Cl2 and tetragonal packing might have effects on its lattice enthalpy. Did you check the enthalpies? $\endgroup$
    – Harshil
    Sep 24, 2022 at 11:14
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    $\begingroup$ Nobody is able to calculate the solubility of a chemical substance in water. It is one of great mysteries of the present science. Why is potassium perchlorate the only practically insoluble potassium compound (apart from silicates) ? Why is calcium fluoride insoluble in water, when calcium chloride, bromide and iodide are soluble in less than their volume of water ? Why ? $\endgroup$
    – Maurice
    Sep 24, 2022 at 12:53
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    $\begingroup$ @Maurice Learn about lattice and hydration energy or whatever, but stop posting this annoying comment all the time. Software can nowadays even design new alloys having desired properties and you keep talking about "mysteries". $\endgroup$
    – Mithoron
    Sep 24, 2022 at 14:20
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    $\begingroup$ @Maurice This is a q & a site. If you have questions then ask them, but first check if they aren't here already. $\endgroup$
    – Mithoron
    Sep 24, 2022 at 16:17
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1 Answer 1

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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.

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  • $\begingroup$ Thank you for explaining why ionic character doesn't mean more solubility. I don't think covalent character always explains less lattice enthalpy because , for example diamond which is covalent crystal, has high lattice enthalpy (713kJ/mol). How could we tell that Hg-Cl covalent bond is weaker than ionic bond between Hg2+ and Cl-? Or am I getting something wrong? $\endgroup$ Sep 29, 2022 at 1:50
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    $\begingroup$ You are correct, my saying that covalent bond is weaker is not correct. It should be, for compounds formed by same atoms, generally covalent bond is weaker than ionic bond, we compare taking the atoms same. $\endgroup$
    – Harshil
    Sep 29, 2022 at 2:04
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    $\begingroup$ chemistry.stackexchange.com/q/11048/124959 $\endgroup$
    – Harshil
    Sep 29, 2022 at 2:04
  • $\begingroup$ I read through your answer again and came to understand what you meant. Thank you so much for taking time to explain your thought. $\endgroup$ Oct 2, 2022 at 11:48

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