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As According to this answer

Now, note that the hydride ion is "hard", having high negative charge density. While the carbonate ion is "soft", having a lower negative charge density.

Applying the "hard-soft acid base" (HSAB) principle, we would expect the "hard" ions to form more stable compounds with each other and the "soft" ions to form more stable compounds with each other.

Since the heavier alkali metal ions are softer, we would expect them to form more stable compounds with carbonate. As for the harder lighter metal ions, we would expect them to form more stable compounds with hydrides.

As Bonding of Potassium and Biarbonate ions is stronger than bonding between Sodium and Bicarbonate( Aka Lattice energy) but the hydration of Sodium is higher than Potassium so according to my intuition when we dissolve Sodium Bicarbonate in water as its hydration energy is more it can overcome it small Lattice energy and it will get solvated by water further and get dissolved and the opposite should be the case for Potassium Bicarbonate, but Experimental answer is completely opposite? Can it be explained or it is a anamlous behaviour cannot be explained by current science

  • $\begingroup$ :/ I guess you didn't consider such details like molar masses... $\endgroup$
    – Mithoron
    Aug 15 '21 at 0:59
  • 4
    $\begingroup$ Predicting correct solubility from the first principles is far from reality. Nobody can look at structure or a compound and tell you its solubility in water. You need experimental information. All those internet stories are just hand-waving stories. Take them with a grain of salt. No pun intended :-) $\endgroup$
    – M. Farooq
    Aug 15 '21 at 4:09
  • 4
    $\begingroup$ Note that both salts do not differ only in the cation size, but also in the ionic lattice structure. Even though both sodium bicarbonate and potassium bicarbonate shows hydrogen bonding in the crystal structure, sodium bicarbonate forms a polymer and potassium bicarbonate forms a dimer. $\endgroup$
    – Poutnik
    Aug 15 '21 at 5:35
  • $\begingroup$ Related: chemistry.stackexchange.com/questions/85312/… $\endgroup$ Aug 15 '21 at 7:01
  • $\begingroup$ Basically, the solubility is an interplay of both lattice energy and hydration enthalpy (if you are talking about the enthalpy of dissolution). In both NaHCO3 and KHCO3, the lattice energy (represented by the product of charges upon internuclear distance) the charges are same (+1 and -1) so the differentiating factor is internuclear distance. But due to the large size of the anion and a relatively small difference in the radii of K and Na, this factor doesn't influence the dissolution energy a lot. So the major difference falls in the hydration enthalpy, which is where NaHCO3 wins $\endgroup$ Aug 15 '21 at 7:12

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