Based on the solubility of copper (II) oxalate it should be supersaturated at any oxalate concentration in the uM range, however when I add 1 mM copper nitrate and more I don't see a precipitate forming in a 300 uM oxalate solution in water. Are the particles too small to see, too sparse to notice, or not forming at all?
I. Copper nitrate is the salt of a weak base (cupric hydroxide) and a strong acid. Hence; this salt is acidic. The dissolved copper(II) oxalate is partially ionized. Some of the oxalate anions may be protonized by protons contributed by the acidic salt; hence ceasing to contribute to the solubility product of copper(II) oxalate; which relates to the equilibrium: CuC2O4(s)↓ ⇌ Cu2+(aq) + C2O42-(aq); Ksp = [Cu2+]·[C2O42-]. That opposes to the additional cupric cation contribution, by copper(II) nitrate, for the same solubility product. Besides the the CuC2O4 neutral complex; also the the cationic complex CuHC2O4+ should form: Cu2+ + HC2O4- ⇌ CuHC2O4+(aq). That can be expected to increase the solubility of copper(II) oxalate ― given in terms of maximum solution formality at equilibrium (equals molarity; yet not accounting for speciation). Solubility then depends on both the stability contant of the CuHC2O4+cationic complex and on the solubility product of copper(II) oxalate. Speciation of copper(II) also includes Cu(OH)+; contributed by Cu2+ hydrolysis: Cu2+ + 2H2O ⇌ Cu(OH)+ + H3O+.
II. Significant disagreement is found on some solubility data reported for copper(II) oxalate (in water). The solubility of copper(II) oxalate in water was said to be 1.6·10-4 M, at 18 ºC (*). The solubility product of copper(II) oxalate was reported as Ksp = 4.43·10−10 M2 (**); this corresponding to a solubility of 2.1·10-5 M, calculated as √Ksp.
(*) https://pubs.rsc.org/-/content/articlelanding/1960/tf/tf9605601165/unauth#!divAbstract
(**) Article Selective Copper Recovery with HCl Leaching from Copper Oxal...
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