[iScience]

Let's say we're dealing with sodium bicarbonate.

Theoretically, where does the solubility of a solute (in a given solvent) derive from?

Please confirm this for me:  When i put sodium bicarbonate in a volume of water, the sodium dissociates from the bicarbonate and the bicarbonate will dissolve since it will interact with the water solvent right? And what determines the saturation concentration is (in this case), the equilibrium for Na⁺ to attach/dissociate right? (please confirm)


Continuing with the assumption that I'm correct, does this mean that if somehow stripped the solution of all the Na⁺ counterions, then there will be no saturation concentration? Well, i suppose if i take out all the counter ions, i'll have a colloidal effect wouldn't i?
Ignoring this effect, would there technically be no theoretical saturation point? (also ignoring the volume of the solute to that of the solution's..)

[mjc123]

Theoretically, where does the solubility of a solute (in a given solvent) derive from?

In a nutshell, Gibbs free energy. There is an equilibrium constant (for ionic salts, the solubility product) given by ΔG° = -RTlnK. This determines how much salt will dissolve in solution in contact with excess solid. To see where the different contributions to ΔG come from, draw a Hess's law cycle. It involves quantities such as the lattice energy, the energy (and entropy) of ion solvation, and the entropy of mixing. These quantities are often large and of opposite sign, so subtle factors may determine whether solubility is low or high in a particular case.

When i put sodium bicarbonate in a volume of water, the sodium dissociates from the bicarbonate and the bicarbonate will dissolve since it will interact with the water solvent right?

The sodium bicarbonate will dissolve - not just the bicarbonate. You have to have charge balance - you can't have a solution of anions alone.

And what determines the saturation concentration is (in this case), the equilibrium for Na+ to attach/dissociate right?

Not quite, it is the equilibrium for Na⁺ and HCO₃^- ions to come together in bulk in a solid lattice, versus the lattice splitting up into separate solvated ions. It is not like a dissociation on a molecular scale, like HOAc  H⁺ + AcO^-.

does this mean that if somehow stripped the solution of all the Na+ counterions, then there will be no saturation concentration?

1. You can't physically do this; you can't have a solution of anions with no cations. (You could remove the Na⁺ ions by replacing them with other cations, e.g. on an ion exchange column, but then you would just have a different salt with a different solubility.)
2. Saturation concentration OF WHAT? Answer: of sodium bicarbonate (or whatever other salt). NOT "of bicarbonate", that is meaningless because, as explained, you can't have a solution simply of bicarbonate. There is "no theoretical saturation point" for such a solution, not because you can add anions without limit, but because YOU CAN'T ADD ANY AT ALL without cations.