[mdk2121]

Hi,

I recently came across a passage in a textbook that is confusing it. It says, "In general, solutes are considered soluble if they have a molar solubility above 0.1 M in solution. Others have only slightly negative changes in free energy, so the equilibrium position lies closer to the undissociated (reactants) side of the reaction. Those solutes that dissolve minimally in the solvent are called sparingly soluble salts."

It was my understanding that if a reaction has a negative Gibbs free energy, it means that it is spontaneous in the forward direction. Correspondingly, it means that the equilibrium constant is positive and that the equilibrium position therefore was closer to the products.

Is my reasoning here wrong? Is the textbook wrong? Or is it just trying to say that in sparingly soluble salts, products are favorited but to a lesser degree than in other soluble salts?

[mjc123]

it means that the equilibrium constant is positive

All equilibrium constants are positive. You mean K > 1 (or ln K > 0).

Or is it just trying to say that in sparingly soluble salts, products are favorited but to a lesser degree than in other soluble salts?

Yes. It's not an either/or, all reactants or all products. If ΔG° is slightly positive or slightly negative, there will be a significant amount of both reactants and products at equilibrium.

Note that there is an error in the quotation from the textbook. If the free energy of dissolution is zero, then K = 1 and the solubility is 1M. So if ΔG° is "slightly negative", solubility > 1M, i.e. it is a soluble salt. For a 1:1 salt with a solubility of 0.1M, K_sp = 0.01, implying ΔG° = +11.5 kJ/mol. So you need a significantly positive ΔG° for it to be "sparingly soluble".