[mdk2121]

I am trying to understand the following question: The molar solubility of Fe(OH)₃ in an aqueous solution was determined to be 4 x 10^-10 M. What is the value of Ksp for Fe(OH)₃ at STP?

I understand that Ksp = [Fe3+] [OH-]^3. I know that we are meant to look at the mole ratios of products to reactants, substituting x to get Ksp = (x)(3x)^3. And I see that we are meant to then reason that x = 4 x 10^-10 M.

However, this is where I get lost: I thought that the whole point of the Ksp was to convey the relative amount of product vs reactant we have at equilibrium/saturation. If we have a very high Ksp, for example, I thought that it meant that the equilibrium favored the products, signifying high solubility (and low Ksp would favor the reactants, signifying low solubility.) If that is the case, then how can we assume that the concentration of Fe3+ is equal to the concentration of Fe(OH)3 at saturation?

If the concentration of Fe3+ ions, for example, is always the same as the molar solubility of Fe(OH)3, then I don't see how the Ksp means anything at all. I would have thought that in a situation of very high Ksp, the [Fe3+] would be much higher than the molar solubility of Fe(OH)3 and that in a situation of a very low Ksp, the [Fe3+] would be much lower. But this question seems to want us to use simple stoichiometry.

Am I missing something fundamental regarding the meaning of Ksp and its relationship with concentrations? Thank you so much!

[Borek]

I thought that the whole point of the Ksp was to convey the relative amount of product vs reactant we have at equilibrium/saturation.

Where is the reactant in the Ksp?

how can we assume that the concentration of Fe3+ is equal to the concentration of Fe(OH)3 at saturation?

Do we?

If the concentration of Fe3+ ions, for example, is always the same as the molar solubility of Fe(OH)3

Is it?

(side note: the question about Fe(OH)₃ is tricky and your solution looks incorrect to me, it is in a way related to some of the mistakes in your thinking)