# Will the solubility product constant differ in this case? [closed]

I am considering the solubility product of sodium citrate. I need this value, I do not have it. However, I do know that the solubility product of calcium citrate is (7±2)×10^−14.

Is it safe to assume that sodium citrate will have a very similar Ksp?

I am concerned about whether or not the addition of sodium citrate will force borate ions out of the water as sodium tetraborate.

Furthermore, could I hypothesise on the solubility based on the likely strength of the hydrogen bonds? I am able to view the charges of each individual atom within a molecule, and the citrate's oxygen anions are far more charged than those in the borate ions. The sodium citrate's hydrogen bonds should therefore be stronger, and result in a higher solubility?

Thank you.

## closed as off-topic by Todd Minehardt, Mithoron, andselisk, airhuff, Geoff HutchisonJan 25 '18 at 15:47

This question appears to be off-topic. The users who voted to close gave this specific reason:

If this question can be reworded to fit the rules in the help center, please edit the question.

• No, because sodium is singly charged and calcium is doubly charged, so there will be a different number of terms in your solubility product formula, and you should expect quite a large difference in the $K_\text{sp}$ between the two salts. – a-cyclohexane-molecule Jan 25 '18 at 1:53

No. The reason you can't find the $K_{\rm sp}$ of sodium citrate is most likely because it is highly soluble. Except for a few exotic cases, sodium salts are very soluble in water.
Adding sodium citrate to the solution does reduce the solubility of borax, however this is not going to be a problem unless you are using solution with really high concentration. As a reference, the solubility of sodium citrate and borax are $\pu{770 g/L}$ at $\pu{25^\circ\!C}$ and $\pu{51 g/L}$ at $\pu{20^\circ\!C}$, respectively.