# Why are entries missing on a solubility data chart for ionic compounds? [closed]

There is a solubility chart in my college chemistry class text for "Solubility of Ionic compounds in water". I have copied and pasted image of chart below. I was surprised to find a dozen ionic compounds marked "--" meaning "no data".

Among the entries marked "no data" are

iron(III) iodide
iron(III) sulfide
iron(III) carbonate
iron(II) chromate
copper(II) iodide
magnesium chromate
silver hydroxide
and carbonate, chromate, and nitrate for Sn2+

It's hard to believe there is no solubility data on said ionic compounds. They should be either soluble, insoluble, slightly soluble, or an impossible compound. I asked my professor, and he said most of said compounds are too unstable to get reliable solubility data. In other words, the said cations and anions would not form a stable ionic compound long enough to test solubility. He also said it might not be the case for all the "--" no data compounds. Which of the said "no data" compounds are impossible compounds as distinguished from possible stable ionic compounds for which there is really no solubility data?

• Perhaps nobody has measured them yet. Or the authors did not find a readily available source. – Jon Custer May 26 at 16:20
• All of iron(III) iodide, iron(III) sulfide, iron(III) carbonate, iron(II) chromate, copper(II) iodide either don't exist or are extremely unstable (guess for IronII chromate, but Fe2+ and )CrO4)2- will almost certainly undergo mutual redox). In fact of those listed it's only Magnesium Chromate that mildly surprises me, all the rest I suspect have at best very limited stability – Ian Bush May 26 at 18:08
• I would ratrher say the data are incomplete. you may want to consult wikipedia - Solubility_table – Poutnik May 27 at 9:11

In the case of magnesium chromate, good solubility in water has been known since 1931. From Hill et al. [1], citing [2]:

The only figure on the solubility of the salt [as of 1940] appears to be that of Kohlraush[2], who reported that at 18°[C] 100 ml of saturated solution contained 60 g of magnesium chromate.

Most of the other compounds listed, as described in comments to the question, do not exist or are highly unstable because the ions are incompatible in the proposed solid. Iodide and sulfide ions are oxidized by iron(III), iodide is also oxidized by copper(II), chromate ions are reduced by iron(II) and tin(II), and tin(II) also reduces nitrate. Carbonate ion is not redox-active, but fails to form a stable salt with some weakly-basic or amphotetic oxide forming metal ions like tin(II) or iron(III); a similar failure occurs with proposed silver hydroxide.

References

1. Arthur E. Hill, Glenn C. Soth, and John E. Ricci, "The Systems Magnesium Chromate—Water and Ammonium Chromate—Water from 0 to 75°", J. Am. Chem. Soc. 1940, 62, 8, 2131–2134. https://doi.org/10.1021/ja01865a059

2. Joseph Mellor, Treatise on Inorganic and Theoretical Chemistry, Vol. II (London: Longmanns, 1931), p. 275.

Besides missing the reference on magnesium chromate, the chart has some errors associated with oxides and sulfides.

Magnesium oxide is listed as insoluble in water, which it is if it is calcined at a high temperature, but otherwise it becomes (kinetically) reactive forming the hydroxide; see this answer. Barium oxide is not "soluble", but reacts in the same way as calcium, strontium, and low-temperature calcined magnesium oxides.

Among sulfides, the $$\ce{S^{2-}}$$ ion is a relatively strong base, so the alkali and alkaline earth metal sulfides that "dissolve" in water are in fact all prone to hydrolysis. Calcium and magnesium sulfides do undergo more complete hydrolysis, as the corresponding hydroxides have limited solubility and their precipitation limits the basicity of the solution. Sulfide ion is also too strongly basic to coexist with ammonium ion, so ammonium sulfide as $$\ce{(NH4)_2S}$$ properly does not exist.

Most of the $$10$$ compounds from your list simply do not exist. This is why they are not reported in your table. For example iron(III) iodide does not exist. When trying to produce it by mixing $$\ce{Fe^{3+}}$$ and $$\ce{I-}$$ ions, a redox chemical reaction occurs spontaneously according to : $$\ce{2 Fe^{3+} + 2 I^- -> 2 Fe^{2+} + I2}$$ For the other compounds, similar reactions happen spontaneously which prevent the compound to exist : $$\ce{2 Fe^{3+} + 3 S^{2-} -> 2 FeS + S}$$ $$\ce{2 Fe^{3+} + 3 CO3^{2-} + 3 H2O -> 2 Fe(OH)3 + 3 CO2}$$ $$\ce{3 Fe^{2+} + CrO4^{2-} + 4 H2O -> Cr(OH)3 + 5 OH- + 3 Fe^{3+}}$$ $$\ce{2 Cu^{2+} + 4 I^- -> 2 CuI + I2}$$ Magnesium chromate does exist, as mentioned by Oscar Lanzi. The other compounds of the list do not exist. For example $$\ce{AgOH}$$ and tin compounds, according to : $$\ce{2AgOH -> Ag2O + H2O}$$ Sn compounds, like tin(II) chromate and tin(II) carbonate are spontaneously decomposed exactly like similar iron compounds. See above. Tin nitrate is spontaneously oxidized according to : $$\ce{Sn^{2+} + 2 NO3^- -> SnO2 + 2 NO2}$$

• According to Housecroft and Sharpe FeI3 has been isolated in "inert conditions" via a photolytic reaction: 2Fe(CO)4I2 + I2 -> FeI3 + 8CO. They do say, though, that it "readily decomposes" – Ian Bush May 27 at 10:35
• "Magnesium chromate does exist". And is used industrially, and has been known to be soluble since 1931 (see my answer). -1 for the people who made that chart for missing that. – Oscar Lanzi May 27 at 11:09
• @Ian Bush. You are right. But I presume this table was done for students trying to mix solutions of different cations and anions. And not in inert conditions like you state. – Maurice May 27 at 11:37